Chemistry 1A: General Chemistry - Las Positas...Chemistry 1A

Chemistry 1A: General Chemistry

Laboratory Manual

Prepared by

Las Positas College

Chemistry Faculty and Staff

Past and Present

Fall 2012 Edition

Page 2

Las Positas College, Chemistry 1A Lab Manual Fall 2012 Page 3

Table of Contents Safety

Laboratory Safety: Laboratory Rules, Self Protection, and Handling of Chemicals and

Glassware………………………………………………………………………………….6

Laboratory Assignments

Math Review 11

Writing Formulas and Nomenclature Worksheets 13

Experiment 1 Significant Figures in Data Collection and Calculation 17

Experiment 2 Library Assignment 21

Experiment 3 Composition and Formula of a Hydrate 25

Experiment 4 Mixing Alcohol and Water—A Thumbsucking Exercise 33

Experiment 5 Identification of Reaction Products 35

Worksheet Stoichiometry Problem Set 43

Experiment 6 Ions in Solution: Electrolyte Strength and Electrical Conductivity 45

Experiment 7 Net Ionic Equations and Reactions in Aqueous Solution 53

Experiment 8 Determination of Copper in a Coin 61

Experiment 9 Oxidation –Reduction Reactions: Predictions and Equations 73

Experiment 10 Determination of the Gas Constant, R 81

Experiment 11 Determination of Sodium Bicarbonate in Alka-Seltzer 89

Experiment 13 Determination of Heat of Reaction 97

Experiment 14 Determination of Crystal Violet by Spectrometry 109

Experiment AA Measurement of Iron by Atomic Absorption (AA) Spectrometry 115

Experiment 15 Group Relationships and Periodic Properties 123

Experiment 16 Model Making and Geometry 145

Experiment 17 Metallic and Ionic Crystal Lattices 157

Experiment 18 Evaporation and Intermolecular Attractions 171

Experiment 19 Using Freezing-Point Depression to Find Molecular Weight 177

Experiment 20 Acid Rain 181

Experiment 21 Chemical Equilibrium: Finding a Constant, Kc 189

Experiment 24 Measuring Sulfur Dioxide in Wine 195

Reference Material

Lab Report Format

Periodic Table

Errors, Precision and Accuracy

Treatment of Experimental Data

Statistics and Uncertainty in the Laboratory

Names, Formulas and Oxidation Numbers of Some Common Ions

Net Ionic Equations

Solubility Tables (4 versions)

Colors of Ions in Aqueous Solutions

Common Oxidation States of Six Elements Important in Redox Chemistry

Activity Series of Metals and Nonmetals

Acids and Bases

Properties of Water: Density and Vapor Pressure

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Safety

Las Positas College

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Laboratory Safety

Read this section on Laboratory Safety. The material contained in this section will be

discussed and augmented by your instructor. All students must also complete “Your

Safety in the Laboratory” on the following pages. Each student must turn in the

completed answer key and the signed “Safety Rules Agreement” before beginning

laboratory work. You should have no trouble if you have done your reading and

participated in the discussion.

A. Laboratory Rules

1. Never work alone! You will work individually, but you should never work in the lab if

the instructor is not present.

2. Report any injury to your instructor at once, no matter how slight it may appear to be.

3. Pregnant students must be especially careful around hazardous chemicals. Pregnant

students must see their instructor for additional safety considerations.

4. Maintain an orderly and clean laboratory desk. Immediately clean up anything you

spill or break. Use a dustpan and brush for broken glass and dispose of broken glass in

the specially labeled container. Keep drawers closed while you are working and keep

stools and backpacks from obstructing the aisles. You should find a safe place to store

your bags and backpacks while you are working.

5. Do not perform any experiments other than those authorized for use that day, unless

you first secure permission from your instructor. Horseplay or gags of any kind are

strictly prohibited.

6. No smoking, drinking, eating or chewing is permitted in the laboratory at any time.

Smoking is only allowed on campus in parking lots or in designated smoking areas, NOT

OUTSIDE LABS! The fume hoods draw smoke directly back into the labs.

7. Do not put equipment in your drawer except for the equipment you were originally

issued. You are encouraged to store your goggles, gloves, and hair bands in your drawer.

8. Do not leave a heat source unattended.

9. Never take a strong whiff of any chemical. If you are instructed to smell a substance,

always waft a small sample of the vapor toward your nose and smell it cautiously.

10. Know the location of and how to use the emergency equipment in your area. Be

familiar with emergency procedures.

11. Avoid distracting or startling any other person. Practical jokes or horseplay cannot be

tolerated.


12. Most experiments are to be performed individually. Experiments in which partners

are allowed or recommended will be specifically identified by your instructor. In any

case, lab reports must be made and submitted individually.

13. Do not leave the laboratory before the period is over, unless you have completed all

calculations and write-up for the experiment. You may find that you have questions or

need to repeat part of the experiment.

14. When you finish your experiment, clean and return equipment to its proper storage

area. Clean the lab benches with a damp towel, clean the sink of all debris, and discard

all chemically contaminated paper and chemical waste in the designated containers. Ask

your instructor if you are unsure about where something should go.

B. Self Protection

1. Safety goggles, of the type approved for chemistry, must be worn in the laboratory.

These goggles seal to the face and have only indirect air vents to prevent chemicals from

dripping or splashing into your eyes.

2. Do not wear contact lenses in the laboratory.

3. If any substance should get in your eye, flush the eye (or eyes) thoroughly (at least 15

minutes) in the eyewash fountain.

4. It is often advisable to wear a lab coat or plastic apron in the lab to protect your

clothing.

5. Wear only fully enclosed shoes; sandals are not permitted.

6. Tie back long hair so that it doesn’t contact flame or chemical solutions.

7. If chemicals are spilled over a large part of the body, use the safety shower at once.

Remove the contaminated clothing. Flood the chemically burned area with water for 15

minutes. Notify your instructor immediately.

8. Do not taste chemicals or anything in the lab. Wash your hands before leaving the lab

so that you do not accidentally ingest chemicals.

9. If you feel faint or dizzy, sit down on the floor. Don’t walk for help, let your neighbor

do that.

10. Don’t panic, whatever the emergency.

C. Handling Chemicals and Glassware

1. Always wash your hands before leaving the laboratory.

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2. Do not take reagent bottles to your desk or into the balance room. Instead, use them at

the side shelves and bring your own container to take what you need.

3. Do not lay the stopper from a bottle down on the table; learn to hold it between your

fingers. Be careful not to exchange stoppers from one bottle to another.

4. If the outside of a bottle becomes wet with liquid, wipe the bottle with a damp towel.

If the bottle has left a ring of liquid on the countertop, wipe up that contamination, too.

5. Whenever pouring liquids from a bottle, “guide” the liquid down a glass stirring rod

held against the lip of the bottle. Done properly, this will practically eliminate the

dribbling problem.

6. Check the label on the bottle both before and after removing the reagent. Check not

only name and formula, but the concentration.

7. Never pipet by mouth or put anything from the lab in your mouth!

8. Never return any excess chemical to a reagent bottle. Dispose of it as directed in the

experiment or ask your instructor.

9. Never use more of a chemical than is called for in the experiment.

10. Always make dilutions by pouring concentrated solutions slowly into water, not the

reverse. Much heat may be evolved, especially in diluting acids or bases. Remember,

add acid to water, not the other way around.

11. Do not grasp recently heated glassware or iron rings, etc. These objects may still be

hot enough to burn you. If you should burn yourself, notify your instructor immediately

and run cool tap water on the burn.

12. Handle glass thermometers carefully. Thermometers with a metallic liquid contain

mercury. Mercury spills require immediate attention. A technician will clean up the spill

immediately.

13. Make sure glassware and equipment are clean before you start your experiment.

14. Disposal of hazardous chemical wastes will be spelled out in the lab manual and/or

by your instructor. Always ask if you aren’t sure where to dispose of something.

15. Before use, inspect all glassware for damaged edges and cracks. Start heating test

tubes and other glassware slowly. When heating a test tube, make sure you are not

pointing it at anyone.

16. Beware of table edges when setting glassware down. Do not place items near the

edge where they can roll or be knocked off.



Laboratory Assignments

Las Positas College

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Name_______________________________

Las Positas College

Chemistry 1A

Math Review

Answer the following problems in the space provided. Show units and significant

figures.

1. 1.61 x 106 + 1.9 x 10

5 =

2. 1.61 x 106

- 1.9 x 105 =

3. 1.91 x 10-5

+ 1.6 x 10-6

=

4. 1.91 x 10-5

– 1.6 x 10-6

=

5. (2.6 x 10-8

)(6.02 x 1023

) =

6. (2.6 x 10-8

)/(0.52 x 10-9

) =

7. (2.6 x 10-8

)(0.25 x 1017

)/(4.6 x 10-9

)

=

________________

8. (2.46)(1.98)/(0.82)(273) =

________________

9. (3.2)1/2

=

10. (4.0 x 10-6

)1/2

=

11. (4.0 x 107)1/2

=

12. (4.0 x 10-7

)1/2

=

13. (4.0 x 10-7

)2 =

14. (3.0 x 10-4

)3 =

15. (3.0 x 103)4 =

16. (12.35)-2

=

17. (1.5)-3

/(2.5)-2

=

18. (2.5)3(1.5)

-2 =

19. (5.0 cm)3/(2.0 cm) =

20. (4.0 cm)2 / (3.0 cm)

3 (2.0 cm)

4

21. (4.0 cm2)2 =

22. (7.5 cm)/(1.5 cm-2

)-3

=

23. Express 4521.3 in scientific notation

=

24. Express 0.0000456700 in scientific

notation =

25. Express 6.72123 x 104 without

exponents =

26. Express 78.7654 x 10-3

without

exponents =

27. Find A where 2A + 2 = 25

A =

28. Find X where XY = 16 and Y2 = 225

X=

29. Find Y where Y2/(0.1) = 4.0 x 10

-9,

Y =

30. (6.4 x 10-14

)1/3

=

31. Find Z where Z2

+ 3Z – 10 = 0, Z =

32. Find Y where Y2 + 2Y = -0.3 Y =

33. (0.070)(0.6023 x 1023

)/(22)=

34. 77.777 – 44 =

35. How many significant figures

in 6.040 x 108?

36. (-2.2 x 104)2 =

37. (-1.3 x 103)3=

38. (-2.0 x 10-2

)4=

39. (4.5)(-2.0) =

40. –(-2.5)(-3.0) =

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Writing Formulas Worksheet Name: _________________________ Fill in the blanks with the correct formulas.

Anions

Chloride

(Cl-)

Sulfate Carbonate Nitrate

Cations ↓

Ammonium

(NH4+)

(NH4Cl)

Silver

Copper (II)

Calcium

Potassium

Mercury (II)

Anions

Hydroxide Oxide Phosphate Sulfide

Cations ↓

Sodium

Magnesium

Iron (II)

Copper (II)

Zinc

Aluminum

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Nomenclature Worksheet Name __________________________________

I. Name the following binary compounds, all of which are composed of non-metals:

NO2___________________________ NO___________________________

CO2 ___________________________ CO___________________________

CS2___________________________ CBr4___________________________

PCl3___________________________ PCl5___________________________

N2O3___________________________ H2S___________________________

II. Name first by the –ous/-ic system and secondly by the IUPAC System

Compound -ous/-ic name IUPAC System name

CuCl

Cuprous chloride

Copper(I) chloride

Hg2O

SnF4

Hg(NO3)2

Fe2O3

III. Name each of the following compounds.

KBr

Be(NO3)2

(NH4)2S

Ag3PO4

Ca(HCO3)2

Li2CO3

BaCrO4

KH2PO4

CdSO4

NaH

MgSO3

Mg(HSO4)2


IV. Write formulas for the following:

Zinc hypochlorite Sodium

permanganate

Silver oxide

Sulfurous acid

Carbonic acid

Zinc phosphate

Ammonium iodide

Stannous fluoride

Hydroiodic acid

Copper(I) oxide

Mercuric oxide

Calcium carbonate

Iron (III) sulfate

Zinc oxide

Nickel (II)

carbonate

Ammonium sulfate

V. Adjacent to the formula, write the name of the salt, then give the formula and the

name of the acid from which each salt may be derived.

FeSO4 Iron (II) sulfate H2SO4 Sulfuric acid

Fe2(SO4)3

Mg(NO3)2

Na2HPO4

KHSO4

NiCl2

SnS2

CaF2

VI. What is the formula (not the symbol) for each of the following elementary substances

when it occurs free in nature?

Hydrogen H2 Nitrogen Oxygen Fluorine

Chlorine Bromine Iodine Helium(!)

VII. Name the following:

H2O ___________________ NH3 ____________________ CH4 _________________

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Experiment 1

Significant Figures in Data Collection and Calculation

PURPOSE

The objective is to show how the number of significant figures is limited by the precision of

the measuring instruments. Here the measurement will be volume, and the measuring

instruments will be beakers, graduated cylinders, pipets and burets.

Method

Various measurements are to be made and reported to the precision allowed by the

measuring instrument. In data collection, always include all certain digits and the first

uncertain (estimated) digit beyond these. This set of digits constitutes the significant figures

appropriate to that measurement.

How the number of significant figures depends upon the measuring instrument is illustrated

below. The same linear dimension is measured with rulers of different degree of precision.

Enter your estimate of the dimension in both situations; you should have two significant

digits for Figure A and three for Figure B.

(A) ____________________ cm*

(B) ____________________ cm*

*Note: enlarged view, not the true size of

a centimeter.

If the instrument (such as an electronic balance) has a digital readout, the last digit may

fluctuate. Take what seems to be a middle value.

PROCEDURE AND DATA SECTION: Write up the procedure and data section below

in your notebook before coming to lab.

Throughout the semester, record all measurements in INK, including proper units,

immediately in your laboratory notebook! The spaces given in this manual are an

indication that you should write something in your notebook, not a proper place to

record data!

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Please read other the reference material for information on errors, uncertainty, statistics and

calculations in the laboratory.

1. BEAKERS AND ERLENMEYER FLASKS

Take a 600 mL beaker and a 500 mL Erlenmeyer flask from your locker. Note the

graduations stenciled on the side of each. Note also that the precision of calibration claimed

is only ±5%. If you were to fill the beaker to the 500 mL mark, how should you express this

volume to show the absolute uncertainty? Fill in the blank below.

Volume = 500 mL ± ______ mL

Now, actually fill the beaker as carefully as you can to the 500 mL mark with tap water.

Then, carefully transfer the water from the beaker to the Erlenmeyer flask.

Questions (answer in your notebook):

1. How well does the new volume reading agree with the old?

2. Although both pieces of glassware have the same estimate of accuracy, why might you

find the Erlenmeyer flask more reliable?

2. GRADUATED CYLINDERS

Visit Station 2 in the lab, where two graduated cylinders partially filled with water are on

display. As precisely as you can, read the volume level in each. Enter your readings in your

notebook in the format shown below. Note that a ± uncertainty is asked for also.

10 mL graduated cylinder ___________ ± ______

100 mL graduated cylinder ___________ ± ______

3. BURETS

Burets are carefully calibrated to be very precise, usually on the order of ±0.01 or ±0.02 mL.

Visit Station 3 in the lab, where two burets are set up as if involved in a titration. The buret

on the left represents the initial reading, and the one on the right represents the final reading

(after you have drained some liquid out of the buret). Record each buret reading to the

proper number of significant figures and uncertainty.

Final buret reading* ___________ ± ______

Initial buret reading* ___________ ± ______


Now calculate the volume of liquid delivered. Report your answer to the correct number of

significant figures and include the total uncertainty. (The total uncertainty will be the sum

of the absolute uncertainties.)

___________ ± ______

Question 3: Why do we say "reading" here, and not "volume"?

4. PIPETS

NEVER use your mouth for suction. An essential part of the exercise is to learn to use a

rubber bulb or other device for filling the pipet.

You will be asked to weigh the water (to ±0.001 g) after transfer to a small flask, and to

repeat this operation two times. The reproducibility (precision) of your work will be a good

indicator of the quality of your technique. One measure of precision, as used here, is to

calculate the average deviation of the data. First, find the average mass of water, then the

deviation of each trial from the average, and finally the average of the absolute values of

these deviations. (Absolute values must be used. Otherwise the sum of the deviations will

equal zero.) Even more significant than average deviation is relative average deviation. To

calculate this, divide the average deviation by the average value of the three trials, and

multiply by 100 to get relative average deviation in parts per hundred (pph or %).

Description of pipet:

Size: ______ Accuracy (from label or manufacturer's catalog): ______

(Also note representative values in other handouts.)

Volume (expressed to proper number of significant figures): _________

Description of water used (This must be deionized water kept at room temperature for some

time.)

Temperature of the water: _________

Density of the water at this temperature : ____________

(from the CRC Handbook of Chemistry and Physics)

Trial #1 #2 #3

mass of water + flask __________ __________ __________

mass of flask empty __________ __________ __________

mass of water __________ __________ __________

Average mass of water __________

Page 20

Deviation from average __________ __________ __________

Average deviation __________(Show calc.)

Relative average deviation __________

( = average deviation/average mass)

Calculated volume of average mass of water __________

(Show calculation in notebook.)

Percent error for the volume of water: (Show calculations in notebook.)

experimental value (label) - true value (calculated)

x 100% =true value (calculated)

5. BAROMETER

Anticipating its relevance to future experiments, we will ask you to practice reading the

barometer and then applying corrections to that reading.

The major correction is that for temperature. As the temperature rises, mercury expands

inside the glass tube and stands higher even though the atmospheric pressure has not

changed. The other correction is for the fact that the gravitational pull of the earth varies

with latitude.

Record the "uncorrected" reading of the barometer: _______________

Record the temperature of the barometer: _______________

Record the temperature correction: _______________

Record the gravitational (latitude) correction: _______________

Calculate the corrected barometric pressure: _______________

Next: Find the atmospheric pressure from the Official LPC Weather site at

http://www.aws.com/AWS/wx.asp?id=LPCAL

Pressure: ________________ inches of Hg.

Convert to mmHg: _________________mmHg

http://www.aws.com/AWS/wx.asp?id=LPCAL


Experiment 2 Name ______________________

Library Assignment **You can do this assignment here in the manual, instead of your lab notebook.

The objective is to give you experience in using our Learning Resource Center (the Library)

to look up scientific and especially chemical information. The three main areas will be (1)

browsing through the stacks of Non-fiction, (2) using the Reference section, with its various

chemical handbooks; and (3) scanning the various scientific magazines to find some

chemical articles that interest you.

Part I Browse through the open stacks of the nonfiction section of the Library. Find one

book in each of the following Library of Congress classifications that both interests you and

would seem useful in your study of chemistry. Note that two letters of the classification

have been given (except for the one which is just Q). Give the rest of the classification

(numbers and letters) and complete the rest of the table.

First two

call letters

Rest of the

Classification

Author(s) Title

Q

QC

QD

QH

QP

RS

TA

TD

As much as you can generalize, what subject matter is treated by books whose call letters

start with:

Q _________________ QC ___________________

QD _________________ QH ___________________

QP _________________ RS ___________________

TA __________________ TD ___________________

Page 22

Part II Reference Section

Before visiting the Reference Section of the Library, use the library catalog to find the call

letters and numbers for the following works and enter them below.

CRC Handbook of Chemistry and Physics (___________________)

Merck Index. (___________________)

McGraw-Hill Encyclopedia of Science & Technology (___________________)

Now choose one of the chemical elements to look up in the Encyclopedia; choose an

element that starts with the same letter as your last name. [If your name begins with J,

choose iodine ("Jod" in German); if Q, choose mercury ("quicksilver"); if W, choose

tungsten ("Wolfram" in German)].

Element _______________ volume & page from Encyclopedia ____________

1) How is this element obtained in the free state?

2) List some commercial applications of this element.

3) Next refer to the Handbook of Chemistry & Physics. Use the index to find the Table of

Isotopes. Find the element you have chosen. How many different isotopes are listed in the

table for this element? Fill in the three blanks below.

Handbook: (______) Edition;

page reference (______);

Number of isotopes (____)


4) Before this next assignment in the Reference Section, find the chemical name of an

ingredient listed on some prescription or over-the-counter drug found in the medicine

cabinet at home. Then look it up in the Merck Index. (Merck is one of the principal

pharmaceutical companies in the world.) Look up the side effects associated with this drug.

Fill in the four blanks below:

Name(s) of drug: _______________________

Merck Index: (_____) Edition; page (_____);

side effects:

4b. Now, search for the same medicine on the Web.

What are the uses of this medicine?

URL (web address): _________________________________

Part III Browse through the Periodicals Section of the Library. List below the names of

five magazines (or "journals", the more professional-sounding term) that regularly seem to

contain articles on chemistry.

1) 2) ___________________________________

3) 4) ___________________________________

5)

Now, from one of these magazines, choose an article on chemistry that interests you. In the

space below write a paragraph telling a little about the content of this article and why it

interests you. Preface the paragraph with title of the article, name of author, name of the

magazine, date of issue and page.

_______________________________________________

_______________________________________________

_______________________________________________

_______________________________________________

_______________________________________________

Page 24


Experiment 3

Composition and Formula of a Hydrate

Prelab – Complete the Prelab on page 31 before coming to lab.

PURPOSE

To determine the percent, by mass, of water in a hydrate compound and to determine the

number of moles of water per mole of hydrate compound – the empirical formula of the

hydrate compound.

Introduction

What is a hydrate? A hydrate is a solid ionic compound that contains water molecules within

its crystalline lattice structure. Frequently the water is chemically bonded less tightly than

other bonds within the compound. Subjecting a sample of hydrate to the heat of a laboratory

Bunsen burner for 5-10 minutes will often drive off all of the water and leave an anhydrous

salt compound residue.

CuSO45H2O(s) + heat CuSO4(s) + 5 H2O(g)

(hydrate) (anhydride)

From the masses of the hydrate, anhydride, eliminated water, and the formula of the

anhydrous salt the composition and the empirical formula of the hydrate can be

determined. Proper technique and careful heating and cooling of the sample should give

precise and accurate results.

PROCEDURE

A. Verify the formula of a known hydrate and perfect your technique with BaCl22H2O.

1. Thoroughly clean a crucible and crucible lid. Place them in a clay triangle suspended

on a ring on a ringstand. Heat them for about 3-5 minutes. Using crucible tongs,

remove the crucible and the lid and let them cool to approximately room temperature.

Carefully weigh the crucible and lid on the analytical balance. Don’t touch the

crucible, even when it’s cool, because fingerprints can affect your results! Use your

time efficiently by cleaning and heating the second sample while the first one is

cooling.

2. When the crucible mass is obtained, add 1-2 grams of solid barium chloride hydrate

to crucible and weigh it – all weights should measured to the nearest milligram. Then

heat the crucible and contents strongly for 4-8 minutes. The crucible should be

cherry red during the last three minutes of heating. The lid of the crucible should be

in place during the heating. Cool and weigh.

Page 26

3. Repeat with a second sample.

4. Calculate the percent water in the hydrate compound.

5. If your two runs do not agree within 2% relative range, consult with your instructor

before going to part B. If time allows, you probably need to do a third trial.

B. Unknown Hydrate

1. Clean your crucibles and lids after Part A is completed and heat them carefully as

before. Add 0.5-0.6 g of unknown hydrate to your crucible and heat as in Part A. Be

careful to avoid splattering. Cool and weigh.

2. Repeat with a second sample – a third if time permits.

3. Calculate the percent by mass of water in the unknown hydrate.

4. When you have completed the calculations in part 3 above, see your instructor for

the formula of the anhydrous compound and then compute the formula for the hydrate

(i.e. how many moles of water are there per mole of the anhydrous compound).

POST-LAB QUESTIONS: Include the answers to the following questions with your

laboratory report.

1. How can an analysis have good precision and poor accuracy?

2. List possible errors that will cause your results to be too low or too high. Indicate if

each error will have a large effect on the results or a minor effect.

3. In a similar analysis, a student determined that the percent of water in the hydrate was

25.3%. The instructor informed the student that the formula of the anhydrous

compound was CuSO4. Calculate the formula of the hydrated compound.

DATA

Enter the sample data report form below in your notebook before coming to lab. During

the experiment, record all data (with proper units and number of significant digits)

directly into your notebook in ink. Show sample calculations where an asterisk, *,

appears at the beginning of the entry description.

Part A.

Run 1 Run 2

1. mass clean empty crucible and lid ______ ______

2. mass crucible, lid, and BaCl2 .

2H2O ______ ______


3. mass BaCl2 .

2H2O (hydrate only) ______ ______

4. mass crucible, lid & anhydride after 1st heating ______ ______

5. mass crucible, lid & anhydride after 2nd

heating ______ ______

6. mass crucible, lid & anhydride after 3rd

heating ______ ______

(if necessary)

7. mass crucible, lid & anhydride after 4th

heating ______ ______

(if necessary)

8. mass BaCl2 (anhydrous salt only) ______ ______

9. mass water, H2O, driven off ______ ______

10. * % water by mass in hydrate, BaCl2 .

2H2O ______ ______

11. * average % water by mass in hydrate, BaCl2 . 2H2O ______

12. * % precision or % relative range of % water by mass

in hydrate, BaCl2 .

2H2O (see error references)

[(high-low)/average] . 100% ______

13. * theoretical % water by mass in hydrate, BaCl2 .

2H2O ______

14. * percent accuracy or relative error of % water by mass

in hydrate, BaCl2 .

2H2O (see error references)

[(expt'l ave. - theo.)/theo.] . 100% ______

Page 28

Part B. Unknown Number ______ (required)

Run 1 Run 2

1. mass clean empty crucible and lid ______ ______

2. mass crucible, lid, and unknown ______ ______

3. mass unknown (hydrate only) ______ ______

4. mass crucible, lid & anhydride after 1st heating ______ ______

5. mass crucible, lid & anhydride after 2nd

heating ______ ______

6. mass crucible, lid & anhydride after 3rd

heating ______ ______

(if necessary)

7. mass crucible, lid & anhydride after 4th

heating ______ ______

(if necessary)

8. mass unknown anhydride only ______ ______

9. mass water, H2O, driven off ______ ______

10. * % water by mass in unknown hydrate ______ ______

11. * average % water by mass in unknown hydrate ______

12. * % precision or % relative range of % water by mass

in unknown hydrate (see error references)

[(high-low)/average] . 100% ______


To calculate formula of unknown hydrate:

13. formula of your unknown anhydride (given by instructor) ___________________

14. * number of moles of water driven off

from unknown hydrate ______ ______

15. * number of moles of anhydride after heating ______ ______

16. * ratio of moles water to moles anhydride ______ ______

17. * average ratio of moles water to moles anhydride _____

18. * formula of unknown hydrate ______________________________

Page 30


PRE-LAB Name_________________

(To be completed before coming to lab)

Experiment 3 Composition and Formula of a Hydrate

1. If you start with 0.572 grams of unknown hydrate and end with 0.498 grams of

anhydrous compound, what is the percentage of water by mass in the sample?

2. How could you determine if you have heated the sample long enough to drive off all of

the water?

3. If you begin with 1.534 grams of BaCl2.2H2O and heat it for 4 to 8 minutes, what will

be the mass of BaCl2(s) that remains?

Page 32


Experiment 4

Mixing Alcohol and Water--a Thumbsucking Exercise

Despite its frivolous-sounding title, this exercise will provide you with an opportunity to

speculate about what must be happening on the molecular level when these two simple

substances are mixed.

PURPOSE

Very little introduction will be given to this exercise. You are to perform two rather simple

operations, make close observations of the phenomena observed, and then speculate about

an explanation.

PROCEDURE

SAFETY FIRST

The hazard level is low in this exercise. However, there is a risk of splashing alcohol in the

eyes, and, as is usual, goggles are required. Also, the use of your thumb to close off the top

of a test tube, as called for here, is not allowed in any future work.

1. Remove two small (10-cm) test tubes from your locker along with a small beaker to

hold them upright. Locate two plastic wash bottles: one, containing deionized water,

will be unlabeled; and the other, containing ethyl alcohol, will be labeled with that name,

along with the warning "Flammable!” There won’t be a flame in today's assignment, but

you should be aware that alcohol, like many non-aqueous solvents, will burn.

2. Fill one test tube about half full with water from the unlabeled wash bottle.

3. Now fill it to the top with alcohol from the labeled wash bottle. Do this carefully. Try

not to mix the two liquids unnecessarily. (They will dissolve in each other, given the

chance, but it is possible to layer the alcohol above the water.)

4. When the tube is "brimful", slide your thumb over the top of the tube carefully so that no

bubbles of air are trapped. (Again, we will not use our thumbs this way in any further

chemical work. It is not good for the thumb or the chemical study. Here, there is no

hazard, and direct contact is necessary to the study.)

5. Be sure your thumb is squarely and firmly positioned on top. Now invert the tube once

or twice and return to the original position, keeping your thumb on top. Observe any

changes in appearance inside the tube and any sensations in your thumb. (Some people

find that they can suspend the tube from their thumb without supporting it! But keep

your free hand underneath, just in case!)

Page 34

6. Discard the solution into the proper waste container.

7. Try the experiment again, but now this time add the alcohol first, and then fill with

water. Do you get the same effect? If not, why not?

POSTLAB QUESTIONS: Answer the following questions in your notebook.

1. In your lab write-up, describe what you saw and experienced in the two different

exercises. Please use complete sentences.

2. What is your explanation for what happened in A, and for the difference in behavior

between A and B? You are not expected to know the whole story. Any hypothesis,

however "wild" is invited.


Experiment 5

Identification of Reaction Products

Prelab – Complete the Prelab on page 41 before lab.

PURPOSE

The purpose of this experiment is to recognize evidence of chemical change, and to write

proper equations, both complete and net ionic, for reactions observed to occur.

Discussion (Read the discussion below but you do not need to enter this section into your

notebook.)

For reactions in this experiment, we will limit our study to the category variously known

as "metathesis," "double displacement," or "partner exchange." Partner exchange can be

illustrated by the generalized equation

AB + CD AD + CB

It can be assumed that the reaction is conducted in water solution, and that, in this

experiment, reactants AB and CD are not only soluble but are dissociated into their

ions.

One can write on paper many chemical equations of this form, with proper

formulas and properly balanced equations, but whether or not they really represent

true reactions depends on experimental investigation.

1) One of the products (AD or CB) is not soluble; that is, it precipitates. (See solubility

rules in the reference section.)

2) One of the products (AD or CB) is a weak acid or base formed from a strong acid

or base (AB or CD). Examples are the formation of the weak acid acetic acid,

HC2H3O2, from the strong acid HCl, or the formation of the weak base NH3 from

the strong base NaOH. (See listing of common weak and strong acids in the

reference section.)

3) One of the products (AD or CB) is a water molecule formed from the

reaction of an acid and base (AB and CD). The reaction, commonly called

"neutralization," has the form "acid + base --> salt + water."

You will be able to detect these chemical changes by careful observation.

1) Precipitates usually form first as cloudy suspensions which then settle out

as visible particles. Some are white; others are colored.

Question: Why are the products not AC and BD? And why are they not DA and BC?

Page 36

2) Strong acid weak acid reactions often show very little sign of chemical change.

In the reactions performed here, the weak acid is unstable and decomposes into a

gaseous component which escapes. The two most common examples are the

formation of CO2 (from H2CO3) and SO2 (from H2SO3). Strong base --> weak base

reactions are similarly hard to detect. However, the most common weak base, NH3,

has a characteristic odor.

3) For neutralization reactions heat evolution may be the only obvious evidence of

reaction. This criterion for reaction is tricky. For instance when concentrated

sulfuric acid is added to water, much heat is evolved but the chemical change

involves just the ionization of the acid, actually:

H2SO4(l) + H2O(l) H3O+

(aq) + HSO4-(aq)

PROCEDURE

In Table I, some ten or so pairings of possible reactants are listed for you to investigate.

Work independently! Aqueous solutions of these compounds in the recommended

concentrations have been prepared for you. Volumes suggested (often just 2 mL) can be

estimated rather than using the graduated cylinder each time. (You can use your

graduated cylinder to measure 2 mL of water into a test tube and keep that tube for

comparison.)

Use your large size test tubes; they will allow you more freedom in mixing. Observe

and record original colors. Add the second reactant to the first gradually, swirling the

tube between additions.

Do not be too quick to report "no reaction". Some precipitates are slow in forming.

Rubbing the inside wall of the test tube with a glass rod sometimes initiates

precipitation in the solution that wets it.

Some precipitates are so finely divided they stay suspended and look like an emulsion--

they look "milky". With time, however, they usually coagulate and settle out. To speed

this process, we often use a centrifuge.

To use a centrifuge:

always insert pairs of tubes (either two or four of your small test tubes, use only

small test tubes); be sure each tube contains approximately the same volume of liquid; fill one with

plain water if necessary;

insert each pair of tubes in the tube holders at 180° (i.e. opposite) to each other (to

keep the centrifuge balanced);

SAFETY FIRST

Wear safety goggles throughout the experiment. Exercise care in handling acids and

bases. Note the caution below on handling AgNO3. Compounds of heavy metals (Ni,

Cu, Ag, etc.) are poisonous if ingested. Wash hands frequently.


keep the lid down until the centrifuge has completely stopped; usually one minute of

centrifuging at full speed is sufficient.

One reason for using the centrifuge is that a clean separation of precipitate is often

necessary before the true color of the precipitate can be distinguished from that of the

supernatant solution. If the color of the precipitate is still not evident, pour off the

solution, add about 1 mL of deionized water, stir the mixture well, and centrifuge it

again.

After the reaction has been studied, dispose of the contents of the test tube into the

proper waste container. Rinse tubes but it is not necessary to dry them before going to

the next test pair.

T ab l e I : Partner Exchange: Mixing Directions

Set Reactants Mixing Directions

1 FeC13 and NaOH 0.1 M FeC13 (2 mL)

6 M NaOH, added dropwise; 10 drops

2 FeC13 and NH3 0.1 M FeC13 (2 mL)

6 M NH3, added dropwise; 10 drops

3 FeC13 and H2SO4 0.1 M FeC13 (2 mL)

3 M H2SO4, added dropwise; 10 drops

4 NiCl2 and Na2CO3 0.1 M NiC12 (2 mL)

0.3 M Na2CO3 (2 mL); centrifuge

5 NiC12 and AgNO3 0.1 M NiC12 (2 mL)

0.1 M AgNO3 (2 mL)

AgNO3 stains the skin; wash hands after contact.

6 Na2CO3 and HC1 1.0 M Na2CO3 (2 mL)

3.0 M HC1, added dropwise; 10 drops

7 NH4NO3 and NaOH 0.2 M NH4NO3 (2 mL)

3 M NaOH (1 mL)

Heat mixture by placing test tube in small beaker of

hot water; hold moist red litmus in vapors; waft

vapors to nose.

8 CuSO4 and Na3PO4 0.1 M CuSO4 (2 mL)

0.1 M Na3PO4 (2 mL); centrifuge

Page 38

9 ZnSO4 and(NH4)2S 0.1 M ZnSO4 (2 mL)

0.1 M (NH4)2S (2 mL)

Mix in HOOD!

10 NiCl2 and(NH4)2S 0.1 M NiCl2 (2 mL)

0.1 M (NH4)2S (2 mL)

Mix in HOOD!

11 H2SO4 and NaOH 3 M H2SO4 (2 mL)

6 M NaOH (2 mL)

Mix all at once!

DATA

Enter a data table for all 11 sets of reagents in the style of the following Sample Data

Table in your notebook before class, but leave enough space for observations. Remember

that the reader might not have performed this experiment, so write sufficient detail to be

informative. Use the format below to record data. For each set of reagents, you should

have five lines blank; A, B, C, D, & E, intended to receive the following information.

Report your observations in ink directly into your lab notebook as you go along.

#1 FeC13 and NaOH

A:

B:

C:

D:

E: ____________________________________________________________

A: Give the complete or "molecular" equation for the predicted partner

exchange reaction.

B: Describe original colors. Describe changes (if any) upon mixing the proposed

reactants. Report any color change; separation of a solid (and its color);

evolution of a gas (and its odor--but waft only a very little to your nose!); and

evolution of heat.

C: If there was a reaction, what is the likely identity of the new substances

formed?

D: Write the “Total Ionic Equation,” that is write each aqueous compound from

part A as separate ions (labeled with the correct charges). Don’t remove

spectator ions yet!


E: Write the net ionic equation for the reaction. If there was no reaction, say so,

"no rxn". (Hint: at least one is a "No Reaction" pair.)

Example of Entry into Report

Set: BaCl2 and CuSO4

A: BaCl2(aq) + CuSO4(aq) BaSO4(aq) + CuCl2(aq)

B: CuSO4 sol'n blue, the other colorless. When mixed, white

precipitate gradually settled out from a blue solution.

C: There was a precipitation reaction. The copper ion remained in

solution. The precipitate was probably BaSO4.

D: Ba2+

(aq) + 2Cl-(aq) + Cu

2+(aq) + SO4

2-(aq) BaSO4(s) + 2Cl

-(aq) + Cu

2+(aq)

E: Ba2+

(aq) + SO42-

(aq) BaSO4(s)

Page 40


PRE-LAB Name_________________


Experiment 5 Identification of Reaction Products

1. Find all the pairs of ions in the following list which cannot exist together (in

appreciable concentration) in the same aqueous solution without precipitating. That is,

write any combination of the following ions which would form a precipitate, a weak acid,

a weak base, or water. [Note: PbCl2 is slightly soluble.]

Fe3+

H+ OH

- Na

+ Cl

- Pb

2+ NO3

- Mg

2+

Answer: __________

2. Convert the following (a) complete or "molecular" equation to (b) the total ionic, and then to (c) the net ionic equation.

(a) Na2 CrO4 ( a q ) + 2AgNO 3 ( a q ) 2 NaNO3(aq) Ag2CrO4(s)

(b)___________________________________________________

(c)____________________________________________________

3. Write all three balanced equations (a, b, and c, as above), given this word equation for a chemical reaction:

barium hydroxide + sulfuric acid barium sulfate + water

solution solution solid

(a) __________________________________________________________

(b) __________________________________________________________

(c) ___________________________________________________________

Page 42


Stoichiometry Problem Set Chem 1A (LPC)

1. a.) How many moles of H3PO4 are represented by 294 grams of H3PO4?

b.) How many atoms of hydrogen are in 196 grams of H3PO4?

c.) What is the % composition of each element in phosphoric acid?

d.) How many grams of phosphorus are in 294.1 grams of H3PO4?

2. a.) Lactic acid is a metabolite formed in the body during muscular activity. It is

composed of 40.00% carbon, 6.71% hydrogen and 53.9% oxygen by weight. What is the

empirical formula of lactic acid?

b.) The molecular weight of lactic acid is 90.08 g/mole. What is the molecular formula?

3. Some commercial baking powders contain a mixture of sodium bicarbonate (baking

soda) and calcium dihydrogen phosphate. When the powder is moistened, carbon dioxide

gas is liberated and makes the batter or dough rise. Balance the equation for the

following reaction.

NaHCO3 + Ca(H2PO4)2 Na2HPO4 + CaHPO4 + CO2 + H2O

4. Chlorine for use in water purification systems may be obtained from the electrolytic

decomposition of seawater. The unbalanced chemical equation for this reaction is:

NaCl(aq) + H2O(l) NaOH(aq) + H2(g) + Cl2(g)

a.) Balance the equation.

b.) What weight of sodium chloride would be consumed in the production of 25 metric

tons of chlorine (1 metric ton = 1000 kg).

5. Iron oxides found in iron ores can be reduced to metallic iron when reacted with

carbon monoxide. The equation for this reaction is:

Fe2O3 + 3CO 2Fe + 3CO2

a.) How many kilograms of elemental iron can be formed if 16.0 kg Fe2O3 is reacted

with 10.0 kg CO?

b.) How many kg of CO2 will be produced in the reaction described in 5a?

c.) 16.0 kg Fe2O3 + 10.0 kg CO = 26.0 kg of reactants

11.2 kg Fe + 13.2 kg CO2 = 24.4 kg products

Explain this apparent inconsistency in the conservation of mass.

6. Xylocaine, a local anaesthetic which has largely replaced novocaine in dentistry, is a

compound of carbon, hydrogen, nitrogen and oxygen. Combustion of a 0.4817 g sample

of xylocaine yielded 1.2665 g of CO2 and 0.4073 g of H2O. A separate nitrogen assay,

using another 0.4817 g sample of xylocaine formed 0.07006 g NH3. What is the

empirical formula of xylocaine?

Page 44


Experiment 6

Ions in Solution: Electrolyte Strength and Electrical Conductivity

Prelab – Please complete the pre-lab on page 51 before lab.

PURPOSE

to relate electrical conductivity to the presence of ions in aqueous solution

to classify substances as strong, weak, and non-electrolytes

to relate electrical conductivity and chemical reactivity to the presence of

hydronium ions in solutions of acids and hydroxide ions in solutions of bases

to classify acids and bases as weak or strong


notebook.)

This experiment addresses the question “Do ions form when a given substance is placed

in solution?” The test we apply is the electrical conductivity of the solution. Whereas

electrical conductivity in metals is based on the movement of electrons, that in solution is

based on the movement of + and – ions. On the basis of the observed conductivity of an

aqueous solution of a substance, we will classify that substance as one of the following:

1. strong electrolytes: substances whose aqueous solutions conduct electricity very

well because many ions are formed

2. weak electrolytes: substances whose aqueous solutions conduct only slightly

because only a few ions are formed

3. non-electrolytes: substances whose conductivity in solution is less than or equal to

that of water because they form hardly any ions or no ions at all

In this experiment you will test the electrical conductivity of aqueous solutions of a

variety of substances that differ in their bond types. These will vary from ionic

substances (salts with a difference in electronegativity, ∆EN, greater than about 1.8 to

2.0, to polar covalent molecules with ∆EN between about 0.4 and 1.8. You will not test

any non-polar covalent molecules (those with ∆EN of 0.4 or less) because they are

generally not soluble in water! Your observations will correlate well with the bond type

of the substances. A few generalizations will help.

1. Solid salts do not conduct electricity when not in aqueous solution. Even ionic

compounds such as NaCl, which exist in the solid state primarily as distinct Na+

and Cl– ions, do not conduct electricity. Because of the solid lattice, ions are

prevented from moving and thus cannot pass charge through the solid.

2. Molten salts do conduct electricity quite well. Melting the solid frees the ions to

move around in the liquid and thus to move charge through the liquid.

3. Salts of high solubility show high conductivity in aqueous solution and thus are

strong electrolytes. Dissolving a salt in water separates the ions, primarily due to

Page 46

the ion/dipole attractions between the ions and the water molecules. The hydrated

ions are free to move through the solution and so conduct electricity quite well.

An example is sodium chloride:

NaCl(s) Na+(aq) + Cl

–

4. Salts of low solubility do not dissolve sufficiently to free many of their ions into

solution. Because so few ions go into solution, no increase in conductivity is seen.

An example is calcium carbonate:

CaCO3(s) Ca2+

(aq) + CO32–

(aq)

A substance must dissolve in water before it can be classified as a strong, weak,

or non-electrolyte. When asked to classify such a substance that does not dissolve

in water, reply “not soluble”.

5. Many polar molecules show very low electrical conductivity in aqueous solution

and are non-electrolytes. (The major exceptions are acids and bases, discussed

below.) Water itself ionizes slightly (a process called autoionization).

H2O(l) + H2O(l) H3O+ + OH

–

However, the concentration of ions is very low, about 10

–7 M in pure water at

25°C, and so the conductivity is very low. Most other polar molecules show even

less tendency to form ions. Even some metal/nonmetal compounds with low ∆EN

values show very little ion formation. Mercury(II) chloride, HgCl2, ∆EN = 1.2, is

a water soluble solid, but it is a non-electrolyte. HgCl2 dissociates to form fewer

ions than does water!

HgCl2(s) HgCl2(aq) HgCl+(aq) + Cl

–(aq)

6. The conductivity of acids of general formula HX is due to their reaction with

water to form H3O+ and X

– ions. Strong acids, such as HCl, react extensively.

Weak acids, such as acetic acid, HC2H3O2, react only slightly.

HCl(aq) + H2O(l) H3O+(aq) + Cl

–(aq)

HC2H3O2(aq) + H2O(l) H3O+(aq) + C2H3O2

–(aq)

Strong acids are strong electrolytes and form almost exclusively H3O+ and X

– ions

in solution. Weak acids are weak electrolytes. They exist primarily as

undissociated HX(aq) molecules in solution, with only minor amounts of H3O+

and X– ions.

The common strong acids are HCl, HBr, HI, HNO3, H2SO4, HClO4, and HClO3.

All other common acids are weak. Only the first hydrogen of sulfuric acid is

strong, essentially 100% ionized. The second is weak, only about 20% ionized at


typical concentrations. Note the positions of the equilibrium arrows in the

ionizations of the two H’s of sulfuric acid:

H2SO4(aq) + H2O(l) H3O+(aq) + HSO 4

–(aq)

HSO4–(aq) + H2O(l) H3O

+(aq) + SO 4

2–(aq)

7. The conductivity of bases is due to the presence of OH– ions. The strong bases are

the soluble hydroxides such as NaOH. For these ionic substances, the ions are

already present and are freed into solution when the substance dissolves.

NaOH(s) Na+ + OH

–

Weak bases such as ammonia, NH3, react only slightly with water. (Most neutral

molecules that are weak bases are amines, organic derivatives of ammonia. Their

general formulas can be written NR3, where R represents either a hydrogen and an

organic carbon group.)

NH3(aq) + H2O(l) NH4+(aq) + OH

–(aq)

Weak acids and bases are weak electrolytes and form only a few ions in aqueous

solution. Strong acids and bases are strong electrolytes and form lots of ions in

aqueous solution. This ion formation will be detected in this experiment by

observing the conductivity of various solutions.

Chemicals

NaCl(s)

1 M NaCl(aq)

CaCO3(s)

CH3CH2OH(l) and (aq, 3 M)

C12H22O11 (aq, 3 M)

1.0 M HCl(aq)

1.0 M NaOH(aq)

1.0 M NH3(aq)

15 M NH3(aq)

HC2H3O2(aq, 1 M)

HC2H3O2(l) [glacial, 17 M]

saturated HCl in toluene

Equipment

conductivity apparatus 50-mL beakers for each substance tested instructor demo:

crucible clay triangle KC2H3O2(s)

Page 48

PROCEDURE

Wear safety goggles. Substance 14 should be worked with under a fume hood. Several

workstations will be set up around the lab. At each station you will measure the

conductivity of one or two substances or solutions. Classify each observation using the

level of conductivity (see instructions on the next page on how to use the conductivity

meter). The instructor will begin by demonstrating some of the conductivities. The

following conductivity tests should be performed:

1. deionized water 10. CH3CH2OH(aq)

2. tap water 11. C12H22O11(aq)

3. NaCl(s) 12. HC2H3O2(l) (17 M)

4. NaCl(aq) 13. HC2H3O2(aq)

5. CaCO3(s) + H2O(l) 14. HCl in toluene

6. KC2H3O2(s) (demonstrated) 15. HCl(aq)

7. KC2H3O2(l) (demonstrated) 16. NaOH(aq)

8. KC2H3O2 (aq) (demonstrated) 17. 1.0 M NH3(aq)

9. CH3CH2OH(l) 18. 15 M NH3(aq)

DATA

Enter a complete data table for all the substances you will test (see sample below) in your

notebook before coming to lab. During the experiment, record all data directly into

your notebook in ink.

Substance Conductivity

Type

(high, medium,

low…)

Electrolyte

Type

(strong, weak,

non)

Major Species

(sol’ns only)

Minor Species

(sol’ns only)

1. H2O

deionized


How to Use the Double-Diode Conductivity Meter

The conductivity meter will enable you to learn whether charged particles — electrons or

ions — are free to move in the substance you are examining. This meter detects five or

six levels of conductivity.

High conductivity means that electrons are free to move easily in the substance (as in a

metal), or that ions are present in large numbers and are free to move.

Low conductivity means that electrons in the substance are not free to move (as in a

nonmetal), or that there are few free ions in the substance or solution.

Before You Use the Meter 1) Turn the switch ON.

2) Pinch the electrodes together. If both red and green diodes glow brightly, the meter is

working and the battery is probably strong enough. The electrodes are best viewed

from the side rather than straight on. (If they don’t light up, or don’t glow brightly,

tell the technician.)

When You Use the Meter

1) Turn the switch to ON.

2) Test the meter frequently, as described above, to be sure it is working.

3) Touch both electrodes simultaneously to the solution or surface you are

testing. On a solid surface, scratch them back and forth several times to

improve contact.

GREEN diode RED diode CONDUCTIVITY

Bright Bright High

Medium Bright Medium

Dim Medium Low

Off Dim Very low

Off Very dim Extremely low

Off Off None

4) When testing a solution, it is sometimes helpful to pinch the electrodes

together and release them a few times in succession while they are

immersed in the solution in order to distinguish between bright and

medium conductivity levels.

5) Clean the electrodes frequently with a tissue to remove traces of the substance you

tested.

6) Turn the switch to OFF when you finish using the meter.

Las Positas College

Livermore, California

JHA110-8-96

Page 50


PRE-LAB: NAME______________________________


Experiment 6 Ions in Solution: Electrolyte Strength and Electrical

Conductivity

1. For each of these substances identify the bond types present (covalent, ionic, or

both). Hint: Check EN (difference in electronegativities).

A. BaC12 B. SO2

C. HBr D. KNO3

E. PbI2 F. H2C2O4

2. Classify each of these ionic compounds as soluble, not soluble, or slightly soluble.

A. Mg (NO3)2 B. PbCl2

C. Fe(OH)3 D. K2SO4

3. Classify each of these as a strong acid, weak acid, strong base, or weak base.

A. H2SO4 B. Ba(OH)2

C. CH3CH2CO2H D. HClO3

E. NH3 F. HF

4. a. Write balanced molecular (not ionic) equations for each of the following.

b. Identify each as PRECIPITATION (forming a solid which is not soluble),

ACID-BASE (formation of a weak acid or base), NEUTRALIZATION (acid +

base forms a salt and water), or DECOMPOSITION (one substance breaks up

into two). If no reaction occurs, write NO OBSERVED REACTION. (Some

reactions may involve more than one type.)

A. HNO3 + KC2H3O2

B. Mg(ClO3)2 + Fe2(SO4)3

C. H3PO4 + NH3

D. Ba(OH)2 + H2SO4

Page 52


Experiment 7

Net Ionic Equations and Reactions in Aqueous Solution

PURPOSE

to predict products for reactions among salts, acids, and bases

to write balanced net ionic equations for those reactions


notebook.)

This experiment presents a brief introduction to reactions involving ions in aqueous

solution and how to write correct balanced net ionic equations for those reactions. The

reactions that will be demonstrated fall into three major types:

1. Precipitation reactions: aqueous solutions of two substances are mixed and form a solid precipitate.

Equations for these reactions may be written by:

a. predicting products by double replacement (partner exchange)—remember that the products

of ammonia reactions can be predicted by thinking of aqueous ammonia as if it existed in

solution as ammonium ions and hydroxide ions

b. applying solubility rules to products to determine whether either of the products (or one of the

reactants!) is insoluble—if no precipitate will form, look for acid-base reactions

c. balancing the equation in molecular form

d. writing the total ionic equation by separating strong electrolytes into ions but keeping weak

and non-electrolytes as molecules—remember to write “NH4OH” as NH3(aq) + H2O(l)

e. canceling unreacted spectator ions, and simplifying coefficients to obtain the net ionic

equation

2. Acid-base reactions: the transfer of H+ between the reactant acid and base. Equations for these

reactions may be written by:

a. predicting products by double replacement (partner exchange)—remember that the products

of ammonia reactions can be predicted by thinking of aqueous ammonia as if it existed in

solution as ammonium ions and hydroxide ions

b. looking for reaction as strong acid or base weak acid or base or as neutralization: acid +

base salt + water—without the formation of a precipitate, a weak acid or base, or water, no

reaction will occur, so write no observed reaction/no expected reaction.

c. balancing the equation in molecular form

d. writing the total ionic equation as above

e. writing the net ionic equation as above

3. Complex formation reactions: a Lewis acid and a Lewis base combine to form a complex. One

type of complex formation reaction will be performed. (See note in the procedure.) Students are not

responsible for writing net ionic equations for complexation reactions at this time.

PROCEDURE

Safety

Many of the chemicals used are toxic, corrosive, or both. Handle them with care. Wear

your goggles at all times. Wear gloves. Wear a lab coat, or else wear old clothes to the

lab. Dispose of the chemicals in the waste jars provided.

Page 54

To perform each reaction, mix about 3 mL (estimated, do not waste time carefully

measuring exactly 3 mL) of each reactant. Record observations of colors, odors, and so

forth before mixing.

Mix the reactants well (the best way to do this is to cover the test tube with parafilm and

shake thoroughly), and then observe the result. Record changes in color, state (formation

of a precipitate or a gas), and temperature (note that temperature changes are often small

and not detectable). Note any odors produced. Remember the technique for detecting

odors: Hold the test tube about 6 inches from your nose. With your free hand, waft the

vapors toward your face. If no odor is detected, gradually move the test tube toward you

nose until an odor is detected (or until the test tube reaches your nose!) Another way to

achieve mixing is to pour the contents of the test tube back and forth between 2 test tubes.

Note also that the test tubes need not be dry if you are going to place aqueous solutions in

them. Just clean, rinse, give a final rinse with deionized water, and shake out the water.

For two of the reactions, you will be asked to use a precipitate from the previous reaction.

To do this, centrifuge the mixture, pour off most of the solution, add about 10 mL of

deionized water, stir the mixture well, centrifuge it, and again pour off most of the

solution. Now proceed with the reaction of the solid. The following reactions should be

performed:

1. Pb(C2H3O2)2(aq) + Na2CO3(aq)

2. The precipitate from reaction 1 plus excess HNO3(aq)

3. Pb(C2H3O2)2(aq) + H2SO4(aq)

4. The precipitate from reaction 3 plus excess HNO3(aq)

5. Pb(C2H3O2)2(aq) + NaOH(aq), add NaOH dropwise until precipitate forms

6. Continue adding NaOH(aq) until the precipitate in reaction 5 dissolves (*)

7. Pb(C2H3O2)2(aq) + 1M NH3(aq)

8. Co(NO3)2(aq) + Na2CO3(aq)

9. Co(NO3)2(aq) + H2SO4(aq)

10. Co(NO3)2(aq) + NaOH(aq)

11. Co(NO3)2(aq) + 1 M NH3(aq)

12. Co(NO3)2(aq) + 6 M NH3(aq) (**)

13. Co(NO3)2(aq) + 15 M NH3(aq) (**)

14. NH4C2H3O2(aq) + Na2CO3(aq) (***)

15. NH4C2H3O2(aq) + H2SO4(aq) (***)

16. NH4C2H3O2(aq) + NaOH(aq) (***)

17. H2SO4(aq) + Na2CO3(aq)

18. H2SO4(aq) + NaOH(aq) (****)

19. H2SO4(aq) + 1M NH3(aq) (****)

(*) complex forms:

Pb(OH)2(s) + OH–(aq) Pb(OH)3

–(aq)


(**) complex forms:

Co(OH)2(s) + 6NH3(aq) Co(NH3)62+

+ 2OH–(aq)

(***) The only observation possible for these three is an odor.

(****) The only observation possible for these two is a temperature change.

DATA

Enter the sample observations table below in your notebook before coming to lab. Make

sure to leave enough space. During the experiment, record all data/observations

directly into your notebook in ink.

Reaction Observations

1. Pb(C2H3O2)2(aq) + Na2CO3(aq)

2. ppt #1 + HNO3

3. Pb(C2H3O2)2(aq) + H2SO4(aq)

4. ppt #3 + HNO3

5. Pb(C2H3O2)2(aq) + NaOH(aq)

6. ppt #5 + excess NaOH

7. Pb(C2H3O2)2(aq) + NH3(aq) (1 M)

8. Co(NO3)2(aq) + Na2CO3(aq)

9. Co(NO3)2(aq) + H2SO4(aq)

10. Co(NO3)2(aq) + NaOH(aq)

11. Co(NO3)2(aq) + 1 M NH3(aq)

12. Co(NO3)2(aq) + 6 M NH3(aq)

13. Co(NO3)2(aq) + 15 M NH3(aq)

14. NH4C2H3O2(aq) + Na2CO3(aq)

15. NH4C2H3O2(aq) + H2SO4(aq)

16. NH4C2H3O2(aq) + NaOH(aq)

17. H2SO4(aq) + Na2CO3(aq)

18. H2SO4(aq) + NaOH(aq)

19. H2SO4(aq) + NH3(aq)

Page 56

Balanced Formula, Complete Ionic, and Net Ionic Equations

For each of the mixtures above for which there was a reaction, write the balanced

molecular, total ionic, and net ionic equations. If no reaction occurred, write the phrase

“No observed reaction”. The numbers correspond to the reaction list. Equations may be

written in pencil. See sample data table below.

1. molecular

total ionic

net ionic

2. molecular

total ionic

net ionic

3. molecular

total ionic

net ionic

4. molecular

total ionic

net ionic

5. molecular

total ionic

net ionic

6. molecular

total ionic

net ionic

Pb(OH)2(s) + NaOH(aq) Na[Pb(OH)3](aq)

Pb(OH)2(s) + Na+ + OH

– Na

+ + Pb(OH)3

–

Pb(OH)2(s) + OH– Pb(OH)3

–

7. molecular

total ionic

net ionic

8. molecular

total ionic

net ionic

9. molecular

total ionic

net ionic

10. molecular

total ionic

net ionic

11. molecular

total ionic

net ionic

12. molecular

total ionic

net ionic

Co(NO3)2(aq) + 6NH3(aq) [Co(NH3)6](NO3)2(aq)

Co2+

+ 2NO3– + 6NH3(aq) Co(NH3)6

2+ + 2NO3

–

Co2+

+ 6NH3(aq) Co(NH3)62+

13. molecular

total ionic

net ionic

Co(NO3)2(aq) + 6NH3(aq) [Co(NH3)6](NO3)2(aq)

Co2+

+ 2NO3– + 6NH3(aq) Co(NH3)6

2+ + 2NO3

–

Co2+

+ 6NH3(aq) Co(NH3)62+

14. molecular

total ionic

net ionic

15. molecular

total ionic

net ionic


16. molecular

total ionic

net ionic

17. molecular

total ionic

net ionic

18. molecular

total ionic

net ionic

19. molecular

total ionic

net ionic

Page 58


POST-LAB Name_________________

(Problems and discussion to be turned in with the lab report)

Experiment 7 Net Ionic Equations and Reactions in Aqueous Solution

Write correct balanced molecular, total ionic, and net ionic equations for the reactions

that occur when the following substances are mixed. All are in aqueous solution except as

noted. Answers may be written in pencil.

1. magnesium chloride and sodium carbonate

2. aqueous ammonia and acetic acid

3. nitric acid and magnesium acetate

4. ammonium chloride and sodium hydroxide

5. barium chloride and calcium nitrate

Page 60

6. sulfuric acid and excess potassium hydrogen carbonate

7. nitric acid and solid silver chloride

8. solid aluminum hydroxide and nitric acid

9. excess aqueous ammonia and sulfuric acid

10. magnesium nitrate and aqueous ammonia


Experiment 8

Determination of Copper in a Coin

Prelab: Complete the Prelab questions on page 71 before lab.

PURPOSE

The purpose of this experiment is to determine the mass percent copper in a nickel coin.

Method

A nickel coin is an alloy of two metals, copper and nickel. The alloy dissolves readily in

nitric acid, yielding a solution of Cu2+

ions and Ni2+

ions. Excess nitric acid, a volatile

acid, is removed by adding the less volatile sulfuric acid, and heating.

The acidity of the solution is reduced somewhat by adding a controlled amount of

ammonia water. Under these conditions, the Cu2+

ion can be reduced to Cu1+

by the

hydrogen sulfite ion, HSO3-, and then precipitated as CuSCN by the thiocyanate ion,

SCN-. The precipitate, copper(I) thiocyanate, is collected on a filter paper, washed, dried

and weighed. From the two weights--that of the original alloy sample and of the final

precipitated compound--the mass percent copper in the original alloy can be calculated.

PROCEDURE

Safety First

As is usual, goggles must be worn throughout this experiment. Hot solutions of nitric

and sulfuric acids are involved. Wash with water at once after any accidental contact of

skin with acid. At the beginning of the experiment, evolution of noxious gases (oxides of

nitrogen) will require use of the fume hoods.

Some more safety notes: NH4SCN: toxic by ingestion; CuSCN: moderately toxic;

DMG: no particular hazards; Ni(NO3)2: strong oxidizing agent, tolerance (as dust) = 1

mg/m3, dangerous fire risk for solid; NaHSO3: contact with solid causes burns to

skin/eyes, strong irritant to skin and tissue, tolerance (as dust) = 5 mg/m3; SO3:

poisonous if inhaled, highly toxic, strong irritant to tissue, oxidizing agent, fire risk with

organics, forms H2SO4 with water producing heat

This procedure assumes a sample size of about 1.25 g, or one-fourth of a nickel. A whole

nickel can be used if the amounts or reagents are scaled up (four times), but the problem

with using the larger size sample is that longer times are required for filtering, washing

and drying the precipitate at the end.

The nickel will have been cut for you. You will be provided with two pieces (for two

runs) along with the year and mint of issue.

Page 62

Plan on making two trials or runs. Read through the directions that follow and see how a

second trial can be carried along at the same time as the first. Be sure to label weighing

dishes, flasks and filter papers to distinguish one sample from the other. Do not get the

two runs mixed up!

Work independently. Steps 1 through 15 must be completed during the first laboratory

period. Read labels carefully; be certain to use the specified concentration. Several

reagents are used with different concentrations at different times during this experiment.

Part I Separation of copper and precipitation as copper(I) thiocyanate.

1. Rub one of the pieces of coin with a paper towel and avoid touching it with your

fingers thereafter. Weigh it to ±0.001 g in a preweighed plastic weighing dish.

2. Transfer the sample to a 250 mL Erlenmeyer flask. Add 15 mL of 6 M HNO3 (dilute

nitric acid, from the reagent shelf). In the fume hood, heat gently (hot plate or low

flame) to get (react) the nickel into solution. Note color changes and fumes emitted.

Do not breathe the fumes. CAUTION! DO NOT LET THE FLASK DRY OUT

OR THE SOLID WILL BEGIN TO "BUMP!" Have some extra 6 M HNO3(aq)

readily available in order to add more 6 M HNO3 as needed to always keep some

liquid in the flask. If someone else cannot use your extra nitric acid, dispose of any

excess nitric acid in the appropriate waste container.

3. When the coin is completely dissolved (10 to 20 minutes), add 10 mL of 3 M H2SO4

(dilute sulfuric acid from the reagent shelf). Heat the solution to boiling. If a solid

forms, add 2 mL more of the 3 M H2SO4. White fumes should appear, at least faintly,

above the mouth of the flask when moist breath is blown across it. This would be

gaseous SO3 (from decomposition of H2SO4) reacting with water vapor to re-form a

fog of H2SO4. This usually takes 10 to 20 minutes also.

Make sure you start the second trial while you are waiting for the first to react.

4. Allow the flask to cool (Never put a very hot flask in cold water, but you may speed

the cooling by setting the flask in cool water) and move it to your desk.

5. Add about 25 mL of deionized water.

6. Measure out into your graduated cylinder 15 mL of 6 M NH3 (dilute aqueous

ammonia, sometimes labeled "ammonium hydroxide", from the reagent shelf). Start

adding the ammonia in small increments to the solution in the flask. You will soon

see a milky blue precipitate, Cu(OH)2, form where the ammonia hits the acidic

solution, and then go away as you swirl the flask. The object is to continue adding

ammonia only until that light blue precipitate just persists after swirling. (It may be

necessary to use more than 15 mL of the ammonia (possibly as much as 25 to 30 mL

more, depending on the amounts nitric and sulfuric acids used earlier) to get the light

blue precipitate. If a dark blue solution containing NO precipitate forms, you have


added too much NH3. See the instructor.) This step takes a lot of ammonia. After

you have slowly added the 15 mL of ammonia, add more using small increments. It

is critical not to overshoot. Be patient! Discard the unused ammonia solution in the

appropriate waste container.

7. Add 15 mL of 1.0 M H2SO4. This solution will be prepared especially for this

experiment, and is not the reagent shelf sulfuric acid used earlier. The precipitate

should dissolve leaving a clear blue or blue-green solution. The solution is now at the

proper acidity, ready for the reducing agent.

8. Add 15 mL of 5% NaHSO3 (sodium hydrogen sulfite, or sodium bisulfite solution).

Note, very cautiously, the characteristic choking odor of SO2 that emanates. Very

little additional change will be observed at this point.

9. Promptly add 15 mL of 10% NH4SCN, ammonium thiocyanate solution, the

precipitating agent. Swirl the flask as the precipitate forms. Continue to swirl the

flask from time to time as the precipitate coagulates; that is, changes from a very fine

grain to a larger grain size. The larger grain size will be more filterable. Warming

will help this process, but do not boil the mixture. If you do not coagulate the

precipitate, there is a good chance that much will pass through the filter paper and be

lost. Contact instructor in this case before disposing of the filtrate.

10. While the precipitate is coagulating, prepare the funnel and filter paper for filtration.

Obtain a piece of 15 cm filter paper (Whatman #40 or equivalent). Fold the paper as

directed by your instructor. Briefly, the directions are to fold the paper in half, then

fold again but stop short of a 90o fold. Weigh the folded portion (after the corner is

torn off). Do not wet the paper while folding; it must be weighed dry.

11. Open up the larger quadrant of the weighed filter paper and fit it snugly into the

funnel. With your wash bottle, moisten the filter paper and press its top edge against

the glass to make a tight seal.

12. Start filtering the reaction mixture. Make a quantitative (without loss) transfer of

precipitate to the filter cone, using squirts of water from the wash bottle to assist you.

The filtrate should be collected in a beaker or flask and saved for proper disposal

later.

13. After all the precipitate has been collected on the filter paper, it must be washed until

free of Ni2+

ions. A suitable test for traces of nickel ion in the wash water is provided

by the reaction of Ni2+

with DMG in the presence of ammonia. DMG is

dimethylglyoxime, an organic compound that forms a bright red precipitate with the

nickel ion. Of course, what we want here is the absence of a red color, proving the

absence of nickel ion in the wash water.

a. Squirt about 1 mL of deionized water onto the solid and let it drain through. Repeat

this twice more.

Page 64

b. Collect the last mL of wash water directly from the funnel stem in a test tube. Be

very careful to avoid ANY contamination from previous washings as you collect

this mL sample of wash.

c. Then add about 1 mL of 6 M NH3 and about 1 mL of 1% DMG solution to the test

tube containing the last mL of wash.

d. Repeat the washing process until the test sample does not show the presence of Ni2+

ion. Collect a fresh sample of the latest washing, that is NOT contaminated with

any of the earlier wash solutions, for each repeated test to determine whether the

precipitate is now free of Ni2+

ion.

14. When washed free of Ni2+

ions, the precipitate is allowed to drain in the funnel. With

a spatula, open up a channel between the paper and the glass to allow the stem to

drain. Later, carefully lift the paper plus precipitate and rest it on a watch glass. Still

later, as it dries, open up the filter paper with your spatula, being careful not to lose

precipitate. In the meantime, dispose of the filtrate in the large jar labeled for this

purpose.

15. Place the watch glass in a safe place in your drawer to dry for two days or more. If

time remains in this lab period, you may start drying your sample in an oven set at

95oC for one hour and you may start the test tube experiments described later in the

section entitled "Some Reactions of Nickel and Copper Ions." Otherwise, do this part

next lab period. It is safe to stop the experiment at this step.

16. On the next lab day, complete the drying of the paper and precipitate by placing the

watch glass either (a) in an oven set at 95oC for one hour, or, (b) under a heat lamp

for one hour. While the sample is drying, perform Part II.

17. Allow the glass to cool, and then weigh just the paper and precipitate. Protect the

balance pan with a piece of paper which has been previously weighed, so that you can

quantitatively determine the weight of only your precipitate and filter paper.

18. If time permits, continue the drying operation for a half hour more, or so. Cool and

reweigh. If the second weighing is within 0.01 g of the first, the sample may be

considered dry. If there is a loss in weight greater than 0.01 g, you would ideally heat

it again until there was no significant change. Such a routine is known as "drying to

constant weight". Dispose of the precipitate in the disposal jar provided.

19. Calculate the mass percent of copper in the nickel coin, assuming that all of the

copper in your alloy sample has ended up as CuSCN.


DATA



directly into your notebook in ink. Show sample calculations for all calculated data

for at least one trial if multiple runs were made.

Part I Description of coin: Date: _______ Mint: _______

Trial #1 Trial #2

Mass of weighing paper ____________ ____________

Mass of weighing paper and nickel ____________ ____________

Mass of sample of nickel ____________ ____________

Mass of dry filter paper ____________ ____________

Mass of protective paper (optional) ____________ ____________

Mass of papers + precipitate (1st) ____________ ____________

Mass of papers + precipitate (2nd) ____________ ____________

Mass of papers + precipitate (3rd) ____________ ____________

(if necessary)

Mass of precipitate ____________ ____________

Calculation of copper content Preserve proper precision, significant digits, in all

calculations!

Mass of copper present ............. ____________ ____________

(Show a clear sample calculation.)

Mass percent copper in coin ........ ____________ ____________

(Show a clear sample calculation.)

Average.................................. ____________

Precision

Deviation from average.............. ____________ ____________

Average deviation ............................ ____________

Relative average deviation, % ............... ____________

(Average deviation/average mass percent)*100%

Accuracy (to be filled in by instructor) ____________% Error

Page 66

Part II - Reactions of nickel and copper ions

PURPOSE

In this second part of the experiment we will investigate some reactions of nickel(II) and

copper(II) ions which will allow us to tell them apart.

Method

A solution containing Ni2+

ions and a separate solution containing Cu2+

ions will be

subjected to the following test reactions (which are run independent of one another):

1) add aqueous ammonia solution

2) add aqueous sodium hydroxide solution

3) add ammonia followed by DMG (dimethylglyoxime) solution

PROCEDURE

For this part of the experiment you will need seven clean test tubes in all. You may work

with a partner, provided that each of you participates in all tests and that each of you

records your own observations. NO more than two students should be working together

on these test reactions!

1. Place 1 mL (20 drops) of 0.1 M Ni(NO3)2 in each of three test tubes. Now in three

other clean test tubes place 1 mL (20 drops) of 0.1 M Cu(NO3)2.

2. Ammonia test.

a. To one of the test tubes containing Ni(NO3)2 add 6 M NH3 drop by drop

until 10 drops have been added. Record your observations after the first

drop has been added and after all 10 drops have been added.

b. To one of the test tubes containing Cu(NO3)2 add 6 M NH3 dropwise as

before until 10 drops have been added. Record your observations.

3. Sodium hydroxide test.

a. To another portion of Ni(NO3)2 add 6 M NaOH dropwise until 10 drops have

been added. Record your observations.

b. To a sample of Cu(NO3)2 add 6 M NaOH dropwise as above.

4. Dimethylglyoxime test.

a. Place 1 drop of Ni(NO3)2 solution in the seventh clean test tube.


b. To the tube with 1 drop of Ni2+

solution add 20 drops of deionized water.

Observe the effect of dilution on the color of the Ni2+

by comparing the color

in this test tube to the color in the one remaining tube containing the 20 drops

of Ni(NO3)2 solution that you prepared above.

c. To both of these solutions in b. immediately above (the test tube with 1 drop

of Ni2+

and 20 drops of water AND the remaining tube with the 20 drops of

Ni2+

) add 1 drop of 6 M NH3 followed by 10 drops of DMG solution. Do you

get a positive test for Ni2+

in both?

d. Perform this test on the remaining Cu(NO3)2 solution from above. Record

your observations.

DATA

Report in ink all observations on mixing reactants directly into your notebook. See

sample observations table below.

Solutions Test tube with 1 mL 0.1 M

Ni(NO3)2

Test tube with 1 mL 0.1 M

Cu(NO3)2

6 M NH3

(1st drop)

6 M NH3

(10 drops)

6 M NaOH

(10 drops)

Test tube with 1 drop

Ni(NO3)2 + 20 drops

H2O (blank)

Test tube with 1

mL Ni(NO3)2

Test tube with 1 mL

0.1 M Cu(NO3)2

6 M NH3 +1%

DMG

Answer the post-lab questions on the following page.

Page 68


POST-LAB Name_________________


Experiment 8 Part II: Reactions of nickel and copper ions

1. How do Ni2+

and Cu2+

differ in their reactions with NH3? with NaOH?

2. Calculate the concentration after 1 drop of 0.1 M Ni2+

solution was diluted with 20 drops

of water. (HINT: M1V1 = M2V2 where M1 and M2 are the molarities before and after

dilution and V1 and V2 are the volumes before and after dilution.)

3. Could you still see the green color of Ni2+

even after dilution with water?

4. Could you still detect the characteristic color of the DMG test for Ni2+

in the diluted

solution?

Equations - Complete and balance these ionic equations:

Ni2+

(aq) + NH3(aq) + H2O(l) Ni(OH)2(s) + NH4+

(aq)

Cu2+

(aq) + NH3(aq) + H2O(l) Cu(OH)2(s) + NH4+

(aq)

Ni(OH)2(s) + NH3(aq) Ni(NH3)62+

(aq) + OH-(aq)

Cu(OH)2(s) + NH3(aq) Cu(NH3)62+

(aq) + OH-(aq)

Ni2+

(aq) + OH-(aq)

Cu2+

(aq) + OH-(aq)

Page 70


PRE-LAB NAME___________________


Experiment 8 Determination of Copper in a Coin

1. Why does nitric acid dissolve (react with) so many more metals than does either

sulfuric or hydrochloric acid? They are all strong acids. (Hint: Dissolving (reacting) a

metal in an acid is an oxidation-reduction reaction.)

2. What is a volatile acid? (See a dictionary).

Which of the following are volatile? HCl HNO3 H2SO4 H3PO4

(Circle the volatile ones. Hint: look up the boiling points and characteristics of each acid,

possibly in the Merck Index).

3. Balance the following net ionic equations:

a. For dissolving metals into solution by reaction with nitric acid

Cu(s) + H+

(aq) + NO3-(aq) Cu

2+(aq) + NO2(g) + H2O(l)

Ni(s) + H+

(aq) + NO3-(aq) Ni

2+(aq) + NO2(g) + H2O(l)

b. For reducing Cu(II) to Cu(I):

Cu2+

(aq) + HSO3-(aq) + H2O(l) Cu

+(aq) + HSO4

-(aq) + H

+(aq)

c. To precipitate Cu(I) as copper(I) thiocyanate:

Cu+

(aq) + SCN-(aq)

4. What is the mass percent copper in copper(I) thiocyanate?

Show clear calculations.

Page 72


Experiment 9

Oxidation –Reduction Reactions: Predictions and Equations


PURPOSE

To use patterns of reactions involved in oxidation-reduction in predicting redox

behavior.

To confirm these predictions from observations of some redox reactions.

To represent these redox reactions using balanced chemical equations.

Introduction

"Redox" is a convenient term for oxidation-reduction reactions. These reactions often

involve a transfer of electrons from one reactant to another. They can be recognized by

noting a change in oxidation number (charge) in one or more elements as they move

from being part of a reactant to being part of a product.

Other references are extremely helpful for writing correct chemical equations which

describe the reactions that you observe. See other sources for discussion and drill on:

1) ionic equations

2) solubility

3) assignment of oxidation numbers

4) common compounds to illustrate various oxidations states of N, Mn, O, Cl, S, etc,

5) methods of balancing redox equations: the half-reaction method

6) acids and bases

Method

Redox reactions cannot be predicted as easily as simple partner exchange reactions, but it

can be done.

Transfer of electrons depends on two things:

1) the ease with which the reducing agent parts with electrons, and

2) the strength with which the oxidizing agent attracts the electrons.

You will not be required to make predictions of reaction products at this time. The factors

involved in such predictions will be discussed later. However, for this lab, we will

introduce the patterns of reactions involved in oxidation-reduction. For this purpose we

can make use of additional information that can be found at the back of this lab manual in

the Reference section:

The Activity Series of Metals

Nonmetals Common Oxidation States of Various Elements

An example of how the second of these tables can be used follows.

Page 74

Example: Predict the products of reaction between K2Cr2O7 and H2O2 in acidic

medium. (It makes a big difference in many redox reactions whether the solution is

acidic or basic.)

Answer to above example:

System:

KCr2O7 and H2O2, acidic solution

Prediction:

From information on oxidation states, we see that Cr in Cr2O72-

is in its highest

oxidation state (+6). Therefore, Cr2O72-

cannot be a reducing agent, but just might be a

good oxidizing agent. We also see that oxygen in H2O2 is in the -1 state, and can go

either up to zero (as in O2) or down to -2 .(as in H2O) . However, if there is to be a

reaction with KCr2O7 H2O2 must act as a reducing agent, and yield O2. The fate of

chromium in Cr2O72-

may be revealed by a change in color.

Experiment:

A dilute solution of K2Cr2O7, was acidified with dilute hydrochloric acid, and H2O2

was added dropwise. The characteristic orange color of the Cr2O72-

ion changed to a

greenish color, and bubbles of a gas were evolved.

Because Cr3+

is greenish in color in this case, and because a gas (O2) is

expected from the reaction of H2O2, we feel justified in proposing the skeleton

equation:

(Acidic) Cr2O72-

(aq) + H2O2 (aq) Cr3+

(aq) + O2(g)

To balance this equation, we use the technique of balancing redox equations discussed

in lecture and in your textbook. In the course of this process we will arrive at the two

half reactions:

6 e- + Cr2O7

2-(aq) + 14 H

+(aq) 2 Cr

3+(aq) + 7 H2O(l)

and H2O2(aq) O2(aq) + 2 H+

(aq) + 2 e-

Continuing the process, these two half reactions, properly multiplied and added, yield

the net ionic redox reaction:

Cr2O72-

(aq) + 3 H2O2(aq) + 8 H+

(aq) 2 Cr3+

(aq) + 3 O2(g) + 7 H2O2(aq)


PROCEDURE

Four possible redox systems will be investigated. Work individually! Specific directions

are given.

Redox Reaction Mixing Directions

System 1: Cu and HNO3(aq) Perform in HOOD!

Place 2 or 3 copper shot (about the size of BB's) in a small test tube in the hood. Add

about 5 mL dilute HNO3 (6 M). Look for a reaction over a period of 5 minutes or so.

Note the color of the fumes above the test tube. Covering mouth of test tube with

parafilm, waiting, and holding tube in front of a white background may help in seeing

the color of gas produced.

System 2: Cu and HCl(aq)

Repeat the procedure of System 1, but substitute 5 mL 6 M HCl for HNO3.

System 3: NaHSO3 and KMnO4

Do three variations: [3A], [3B] & [3C]

[3A] Acidic solution

In large test tube, 0.1 M NaHSO3 (2 mL) + 3 drops of 0.5 M H2SO4; then add 10

drops of 0.1 M KMnO4, drop by drop. The sign of reaction will be the loss of the

permanganate color.

[3B] Slightly basic solution

0.1 M NaHSO3 (2 mL) + 2 drops 1 M NaOH; then add 0.1 M KMnO4 dropwise

(about 10 drops).

[3C] Strongly basic solution

0.1 M NaHSO3 (2 mL) + 1 mL 6 M NaOH; then add 0.1 M KMnO4

dropwise (about 10 drops). Note two reactions: the first, a color change, is

unique to system [3C]; the second, a precipitate, is like that of [3B].

System 4: KI and FeCl3

In a large test tube place 0.1 M KI (2 mL) + 0.1 M FeC13 (2 mL)

Note the color and form of the precipitate of I2(s) formed. This precipitate

may be very slow to form.

SAFETY FIRST: Wear goggles throughout the experiment. Exercise care in handling

acids and bases. Brown fumes from reaction of nitric acid are toxic. Perform such

reactions in the hood. Do not breathe or smell the fumes. To quench (stop) a reaction,

add several volumes of water.

Page 76

DATA

Enter the sample report form below in your notebook before coming to lab. Record

your observations, as you go along, directly into tables in your lab notebook in ink,

not into the sample tables provided here. Remember that the other references on

oxidation and reduction are extremely helpful in writing the correct equations to describe

these redox reactions. Answer as many of these questions as possible before coming to

lab.

SYSTEM #1: Cu and HNO3(aq) #2: Cu and HCl(aq)

Prediction: Do you expect

H2(g) to be evolved? Why

or why not?

What gas(es) might be

emitted if the nitrate ion

reacts?

Observations:

Net ionic equation for reaction, if any.

System #1: Cu and HNO3(aq)

_____________________________________________________________________

System #2: Cu and HCl(aq)

_____________________________________________________________________

System 3: NaHSO3 and KMnO4

Prediction:

Is the HSO3- ion a possible oxidizing agent? _________________________

If 'yes', to what species could it change? _________________________

Is it a possible reducing agent? _________________________

If 'yes', to what species could it change? _________________________

Is the MnO4- ion a possible oxidizing agent? _________________________


If 'yes', list the species to which it could change, the color, and conditions that favor each

change:

species color conditions

Why is the MnO4- ion not a possible reducing agent? ________________________

Observations:

[3A] acidic: ____________________________________________________

[3B] slightly basic: ____________________________________________________

[3C] basic: ___________________________________________________

Net Ionic Equations: Write balanced redox equations for the three reactions in System 3

using the half reaction method. (Hint: one of the products for [3A] is HSO4-(aq))

[3A] __________________________________________________________

[3B] __________________________________________________________

[3C] __________________________________________________________

System 4: KI and FeCl3

Prediction:

What possible redox role(s) can the iodide ion assume? ____________

To what species could it be converted? ____________

What possible redox role(s) can the Fe3+

ion assume? ____________

To what species could it be converted? ____________

Observations: ___________________________________________

Net Ionic Equation (Use half reaction method):

_______________________________________________________________________

Page 78


PRE-LAB NAME___________________


Experiment 9 Oxidation –Reduction Reactions: Predictions and Equations

See Reference: Common Oxidation States of Various Elements with Representative

Compounds

1. Determine the oxidation number of each element in the following ions or compounds:

(a) OH- (b) H2O2 (c) O2

2- (d) S2O7

2 - (e) Cr(OH)4

-

2. For each of the following reactions, tell whether or not it is a redox reaction. If it is, tell which reactant is the oxidizing agent (OA) and which is the reducing agent (RA).

Yes/No OA RA

(a) Cu(s) + 2 AgNO3(aq) Cu(NO3)2(aq) + 2 Ag(s)

(b) NaHCO3(s) + HC1(aq) NaCl(aq) + H2O(l) + CO2(g)

(c) Fe(s) + Cl2(g) FeCl3(s)

3. In the following problem, "o.s." means "oxidation state". Write the formula of any salt

you can think of which contains:

S in +6 o.s. S in +4 o. s. S in -2 o. s.

N in +5 o.s. N in +3 o.s. N in -3 o.s.

Cl in +7 o.s. C1 in +5 o.s. Cl in -1 o.s.

4. Balance the following half reaction, and tell whether it is a reduction or an oxidation half

reaction.

NH4+ (aq) NO3

-(aq) (Acidic)

Which is it?

Page 80


Experiment 10

Determination of the Gas Constant, R


PURPOSE

In this experiment we will evaluate the magnitude of the gas constant R in the ideal gas law

equation, PV = nRT. We will assume that we don’t know the value of R and we will

measure P, V, n, and T.

Method

To solve the ideal gas law, PV = nRT, for R, we need to have values for all the others: P, V,

T and n, describing an “ideal gas.” A direct approach would be to weigh a known volume

of a known gas and convert it to moles using molar mass. However, because of their low

densities, gas weights are difficult to measure directly. In this experiment the gas (H2) will

not be weighed. Instead, the number of moles present will be calculated from the

stoichiometry of the reaction which yielded it, namely

Mg(s) + 2 HCl(aq) MgCl2(aq) + H2(g)

The mass of one of the reactants, magnesium, will be measured carefully and then the molar

amount can be calculated. The other reactant, hydrochloric acid, will be taken in excess.

And, of course, the volume, temperature and pressure of the gas will be recorded.

The volume and temperature of the gas present no problems because they are direct

observations. The volume will be measured with a gas buret. (See the figure on the next

page.) The temperature will be measured with a thermometer. The number of moles of gas

is derived from the moles of magnesium reacted and the stoichiometry of the reaction

equation.

Measuring the pressure of the gas is complicated by two factors:

1. The gas pressure inside the buret is not the same as the barometric pressure.

2. The gas inside the buret is really two gases: H2 and H2O(g) (water vapor). Each gas

occupies the same volume but each at its own partial pressure.

First, we will find the total pressure of the gases, Pgas. It is the corrected barometric pressure

of the air, Pair, minus the pressure difference between the outside air and the trapped gases,

H2(g) and H2O(g). This pressure difference, ΔP, is due to the height of the column of water, h.

(Again, refer to the figure on the next page.) How to make this correction as well as the

reasoning behind it is given elsewhere. The relationship is:

Pgas = Pair - ΔP

Page 82

Second, we will find the vapor pressure of water at the temperature of the water, PH2O, (from

Appendix) and subtract it from the total pressure of the gases, Pgas, to yield the partial

pressure of the hydrogen, PH2.

PH2 = Pgas - PH2O

It is this partial pressure of hydrogen that we will use in the ideal gas law

PH2V = nH2RT .

PROCEDURE

SAFETY

Wear your safety goggles throughout this experiment. It involves an acid solution from

which a gas is evolved. The gas, hydrogen, is explosive. No flames or hot plates are

allowed in the vicinity.

1. Assuming you are using a gas buret of 100 mL capacity, your magnesium sample must

not exceed 0.080 g (about 8 cm). However, the oxide coating usually present must first be

removed before the sample is weighed. Place the ribbon on a piece of cardboard, and polish

each side with a piece of emery paper until the metal is

shiny. Wipe the ribbon free of dust with a paper towel.

Until it is weighed, avoid touching the strip with your

fingers, which could leave oils and moisture on the

surface.

2. Carefully weigh the magnesium to ±0.001 g.

Compress the weighed sample into a small bundle.

(Fingers are okay now.) You can coil it loosely around a

pencil and then squash it. Next take about 20 cm of fine

(#24) copper wire and wrap it all around the bundle like a

cocoon. Leave only uniformly small holes. The idea is to

make a cage about the magnesium to allow the acid to

move through the holes to react with the magnesium and

to also hold the magnesium in place as it disintegrates

during reaction. Check to see that the wire cage will slip

inside a gas buret tube and then set it aside.

3. Along with the 100 mL gas buret, obtain a ring stand

with buret clamp attached. Rinse the buret with tap

water. If beads of water remain on the inside wall when

it drains, the buret is not clean. Scrub it with a buret

brush and detergent solution, and rinse it thoroughly.


4. Into the gas buret pour about 10 mL of 6 M HCl (dilute hydrochloric acid). Use a funnel

in this operation to avoid getting acid on your hands. If you do, wash your hands promptly.

5. Remove the funnel and start adding deionized water to the buret from a small beaker.

Hold the buret at an angle from the vertical, and pour the water down the side slowly. The

idea is to layer the water over the acid with the least possible mixing. Fill it to the very top,

brimful. Clamp the buret in the vertical position while you prepare the rest of the assembly.

6. Have ready a 400 mL beaker, half full of water. (Tap water is okay here.) Also have

ready a one- or two-hole rubber stopper that will fit the buret.

7. Insert the magnesium sample about 4 cm into the buret. Lock it in place by inserting the

rubber stopper, pinching the wire stem. Be sure your stopper has a hole (or two) in it!

There should be no air bubbles below the stopper. See figure (a).

8. Quickly cover the stopper hole with your finger, and invert the buret into the beaker of

water. Remove your finger when the stopper is immersed. Clamp the buret in this position.

See figure (b).

9. The more dense acid, now at the top, will sink to the bottom and start to react with the

magnesium. As hydrogen is generated, it rises to the top, displacing water, which leaves

through the hole in the stopper. Toward the end of the reaction, small bits of magnesium

may break free and float to the top. (How can magnesium metal float on water?) Do not let

these pieces get stranded on the walls. Cautiously tilt the buret, if necessary, to release

them, but be sure to keep the rubber stopper submerged at all times.

10. When all the magnesium has disappeared and no more bubbles of gas arise, the

reaction is over. However, do not be too quick to disassemble the apparatus. You first must

make three measurements.

a. Read the volume of gas collected.

b. Measure, with a millimeter ruler, the height, h, in figure b, the height the water

stands in the buret above the surface of water in the beaker.

c. Measure the temperature of the water, which we will take to be the temperature of

the gas overhead.

11. Read and record the room temperature and uncorrected barometric pressure.

12. Empty the buret and beaker into the waste container. Rinse the buret, and repeat the

determination at least two more times.

DATA



Page 84



Room temp._________

Barometer temp. latitude corrected

reading...__________corr'n________corr'n.._______bar. reading_____________

Trial 1 Trial 2 Trial 3

1. Mass of magnesium __________ __________ __________

2. Volume of gas collected __________ __________ __________

3. Temperature of water (& gas) __________ __________ __________

4. Height, h, elevation of water

level in buret (in mm) __________ __________ __________

5. Vapor pressure of water, PH2O __________ __________ __________

CALCULATIONS

Show sample calculations where an asterisk, *, appears at the beginning of the entry

description.

1. * Partial pressure of H2, PH2. __________ __________ __________

(corrected for vapor pressure and height difference of water)

2. * Moles of hydrogen, nH2 __________ __________ __________

3. * Calculated value of R __________ __________ __________

(from above exp'tl data)

4. Average __________

5. Deviation from average __________ __________ __________

6. Average deviation __________

7. Relative average deviation __________

8. * % error (line 4 value compared to

accepted value, 0.08206 L.atm/mole.K) ______________


POST-LAB Name_________________


Experiment 10 Determination of the Gas Constant, R

Another way of expressing the results of this experiment is to calculate the molar volume,

V/n, of the gas at 0oC and 1.000 atm pressure, conditions known as STP, standard

temperature and pressure. Calculate from your data the value of molar volume, V/n, in units

of liters per mole at STP. Note that you are to use your experimental values for pressure,

volume and temperature in this calculation and NOT use either the experimental or

theoretical value of the gas constant, R, anywhere during this calculation.

Ans.

Would you expect to get the same % error as you got for your determination of R? Try it

and see. The accepted value for molar volume is 22.41 L/mol at 273.15 K and 1.000 atm

pressure. Show your supporting calculations.

% error __________

Is it the same yes or no ?

Page 86


PRE-LAB Name_________________


Experiment 10 Determination of the Gas Constant R

1. Write two balanced equations:

a. the reaction of magnesium with hydrochloric acid:

b. the reaction of magnesium oxide with hydrochloric acid:

2. In this experiment, care is taken to remove any "oxide coating" from the magnesium

metal before it is weighed. If a student failed to do this, how would this error affect the

calculation of R?

Too high______?

Show your reasoning:

Too low_______?

No effect_____?

Hint:

Solve the ideal gas law equation PV = Nrt for R. Then see how the error would affect

a. the numerator

b. the denominator

3. The magnesium sample is wrapped in fine copper wire before insertion into the acid

solution. Why isn't fine iron wire used instead? It's cheaper!

4. The experiment calls for about 0.08 g of Mg and 10 mL of 6 M HCl. What volume of

6.0 M HCl is theoretically required to react with 0.080 g of Mg? Show calculations:

Page 88


Experiment 11

Determination of Sodium Bicarbonate in Alka-Seltzer


PURPOSE

The purpose of this experiment is to determine the mass percent of sodium hydrogen

carbonate (sodium bicarbonate) in an Alka-Seltzer tablet.

Method

The carbon dioxide evolved when Alka-Seltzer is added to water will be collected. Its

volume will be determined by measurement of the volume of water it displaces.

Application of the ideal gas law will allow calculation of moles of CO2. The number of

moles of CO2 is related to moles of NaHCO3 by the net ionic equation:

HCO3 -(aq) + H

+(aq) H2O(l) + CO2(g)

Ordinarily, the acidity necessary for the reaction comes in the tablet itself in the form of a

solid organic acid, citric acid. However, in this experiment an inorganic acid, HCl, will be

added to the water as well to speed the reaction to completion.

One problem that must be faced is the appreciable solubility of CO2 in water. To reduce this

source of error, the water in the collection system will be pre-saturated with the gas. Then,

in calculating the final result, we will correct for the solubility of CO2 in the reaction

solution.

Another problem is that the Alka-Seltzer tablet must be kept dry until the controlled reaction

starts. Direct contact with the fingers and other sources of moisture must be avoided.

PROCEDURE

SAFETY

Wear your safety goggles throughout this experiment. It involves handling an acid solution

from which a gas is evolved. The gas, CO2, is harmless, but could expel acid if released

suddenly. Note: instructor's approval of apparatus is required before first run.

1. Handling the Tablet

a. Work individually. Each student will receive a pair of tablets sealed in foil, and

several small plastic weighing dishes. Use a pencil to make identifying marks on the

weighing dishes.

Page 90

b. Accurately pre-weigh (± .001 g) one of the dishes and record the mass. Open the foil

packet and, using the pre-weighed dish, accurately weigh one of the tablets to ±0.001 g.

Record the mass. Avoid wasteful opening of any more foil packets than absolutely

necessary.

c. Pick up the tablet using a piece of the foil or a piece of Parafilm to cover your

fingers. Place a ruler as a guide across a diameter of the tablet and score (scratch) the

tablet with a knife point or large pin. Still using the foil or Parafilm snap the tablet into

two halves. Shake off any crumbs or powder.

d. Using a pre-weighed dish, accurately weigh one of the halves (±.001 g). Then affix

one end of a fine 10 inch thread to the half tablet. To do this, use a very small piece of

cellophane tape (about one-fourth inch). Save the other half-tablet for saturating the

water in the collection system with CO2.

2. Saturating the Water with CO2

a. Fill a 500 mL Florence flask about half full with tap water. Pour in about 10 mL of 6

M HCl. Swirl to mix the solution, and add more tap water until the level is about three-

quarters up the neck of the flask.

B.Into the flask, drop the unweighed half-tablet, and set the flask aside as the solution

effervesces.

3. Establishing the Siphon

a. The siphon system, shown in Figure 1 to the

right, has been preassembled for you. Insert it

into the Florence flask after effervescence

(fizzing) has stopped. Have available a 600

mL beaker and a pinch clamp.

B.Using your pipet bulb, squeeze air into

stopper D to start a flow of water into the large

beaker, C, in Figure 2. When you see no more

air bubbles passing through the siphon, apply

the pinch clamp E to the rubber tubing to stop

the flow.

c. If the level of water in B is now appreciably below the neck of the flask, lift beaker C

above the level of water in flask B, loosen the clamp E and allow water to return from

beaker C to flask B, while you always keep the end of the glass tube below the surface of

the water in beaker, C, but retain some water in the beaker C for pressure equalization

(discussed later). Note that you must keep the end of the glass tube submersed under the

water level in beaker C as you run the water from beaker C back to flask B. Be sure the

stopper makes a very tight seal in flask B.


4. Preparing the reaction flask, A

a. Insert a funnel into a 250 mL Erlenmeyer flask. Pour in 25 mL of 6 M HCl through

the funnel so as to not wet the sides of the flask. Remove the funnel and clamp the flask

to a ring stand.

b. Pick up the thread with the half tablet attached, and lower it inside the flask (very

carefully!) until it is about 2 to 3 cm above the acid solution. Be sure the tablet stays dry

on the way down. Immediately insert stopper D to secure the thread and to make an

airtight seal in flask A.

Stoppers must be tightly pushed into flasks B & A. Check!!!

c. Remove the pinch clamp E. There should be just a very small, brief flow of water

from flask B to beaker C. If water continues to flow, you have a leak and all connections

should be checked.

5. Pressure Equalization

a. The pressure inside flask A will not be the same as that in the room. To equalize the

pressures, keep the end of the glass tube under the surface of the water in beaker C and

raise beaker C until the level of water in beaker C is the same as in flask B. While

holding the water levels in beaker C at the same level as in flask B, put the pinch clamp

back on the tubing. If you need more hands to accomplish this task, ask someone around

you for help.

b. Now carefully hold the glass tube so no water drips out and pour the water from the

beaker C into another container to save as carbonated water for future trials. Drain the

beaker C well but do not dry it. Return the beaker to its normal position C with the glass

tube inside again. (This beaker C will collect the water displaced by the gas evolved

during the reaction.)

c. Remove the pinch clamp. Like before, you may see a very small, brief flow of water

from flask B to beaker C. You are now ready to start the reaction, but only after you get

Page 92

your instructor's approval and initials on your data sheet. Failure to remove the pinch

clamp would cause pressure to build up and the system would blow apart somewhere. So

be sure to get your instructor's approval before the first trial. On subsequent trials, ask a

person near you to provide an extra check to verify that the system is ready.

6. Starting the Reaction

a. Tilt flask A until the tablet sits in the acid solution. Clamp the flask A in this position,

but swirl it from time to time to dissolve the entire tablet. When no more effervescence is

noted, the reaction is over (about 10 minutes).

b. Do not measure the water in the beaker just yet; the gas pressure must first be

equalized as it was originally. Raise or lower the beaker C until the water levels are the

same between beaker C and flask B, just as you did at the start of this run keeping the end

of the glass tube submerged) and replace clamp E on the tubing. You have now captured

a volume of water in the beaker equal to the volume of gas evolved (CO2 and water

vapor) measured at the barometric pressure of the laboratory.

c. Measure the water in beaker C carefully in a graduated cylinder. Do not discard this

carbonated water. Save it for future trials.

d. Open the system and measure the temperature of water in flask B. We will take this to

be the temperature of the gas as well.

e. Empty flask A into the waste container, rinse it, and wipe dry the inside of the neck.

You are now ready for another trial. Using a preweighed dish accurately weigh the

second tablet and record the mass. Score it and snap it into two pieces. Discard small

fragments. Using preweighed dishes, accurately weigh separately each of the two large

pieces and record the masses. Be sure to put identifying marks on the plastic weighing

dishes so that you know which mass belongs to which run.

7. Collecting Other Needed Data

a. Read the barometer, and note the temperature and latitude corrections.

b. Determine the vapor pressure of water at the temperature of the gas collected; consult a

table elsewhere.

c. The solubility of CO2 in the acid solution in reaction flask A is assumed to be 20 mL.

d. Record the manufacturer’s listing of the composition of Alka-Seltzer on the data sheet.


DATA





Name of Alka-Seltzer preparation tested: _____________________

1st tablet 2nd tablet

1. Mass of weighing dish _________ _________

2. Mass of weighing dish and whole tablet _________ _________

3. Mass of whole tablet _________ _________

Trial 1 Trial 2 Trial 3

4. Mass of weighing dish _______ _______ _______

5. Mass of dish and half-tablet _______ _______ _______

6. Mass of half-tablet _______ _______ _______

Approval of apparatus (instructor)

7. Volume of water collected _______ _______ _______

8. Temperature, water (and gas) _______ _______ _______

9. Vapor pressure of water _______ _______ _______

10. Barometric pressure __________ - __________- __________ = __________

uncorrected temperature latitude corrected

reading correction correction reading

CALCULATIONS Show a sample calculation for each one below for at least one trial.

1. partial pressure of CO2 _______ _______ _______

2. solubility of CO2 in flask A 20 mL 20 mL 20 mL

3. corrected volume of CO2 _______ _______ _______

4. moles of CO2 _______ _______ _______

Page 94

5. moles of NaHCO3 reacted _______ _______ _______

6. mass of NaHCO3 reacted _______ _______ _______

7. % by mass NaHCO3 in this

Alka-Seltzer preparation _______ _______ _______

8. average mass percent NaHCO3 _______

9. deviation from average _______ _______ _______

10. average deviation _______

11. relative average deviation, % _______

Calculation of % by mass sodium bicarbonate claimed on label

Name of Alka-Seltzer preparation ____________________

1. from the label: mass of NaHCO3 per tablet __________

2. average mass of a tablet (from your data, line 3) __________

3. calculated mass % NaHCO3, based on label __________

Calculation:

4. taking label for true value, calculate % error __________

Calculation:

POST-LAB QUESTION

Answer the following question in your notebook.

What factors might explain why your determination of mass percent NaHCO3 may not agree

with the information on the label? Be specific, but do not say things like "mistake in reading

the balance" or "mistake in reading the graduated cylinder." We assume you do not make

blunders like that.


PRE-LAB Name_________________


Experiment 11 Determination of Sodium Bicarbonate in Alka-Seltzer

Use the following data obtained in the analysis of an Alka-Seltzer tablet to make the

calculations asked for. They parallel those in the experiment to be performed. See the

sequence of calculations given in the Data and Calculations sheet for helpful hints.

Data

Mass of whole tablet 3.478 g

Mass of half-tablet taken for reaction 1.973 g

Volume of CO2 + water vapor (corrected) 314 mL

Temperature of gas 22.0oC

Vapor pressure of water at 22.0oC 21 torr

Barometric pressure, corrected 757 torr

Calculations (Show your work.)

1. partial pressure of CO2 __________

2. moles of CO2 __________

3. moles of NaHCO3 __________

4. mass of NaHCO3 __________

5. mass % NaHCO3 in Alka-Seltzer __________

6. mg of NaHCO3 per one tablet, calculated __________

On its label, the company (Miles Laboratories) claims 1904 mg NaHCO3 per tablet. Taking

this to be the true value, what is the % error in the above experimental work?

__________

Page 96


Experiment 13

Determination of Heat of Reaction (using Logger Pro®)


PURPOSE

The objects of this experiment are to:

a. Determine the total heat capacity of a simple Styrofoam cup calorimeter

b. Measure the heat (enthalpy) change of two related chemical reactions, and

c. Calculate the heat change of a third reaction by applying Hess' law.

Method

The calorimeter used here will be similar to the one used previously. You will use

LoggerPro® to collect and display data as a graph and table, analyze your experimental

data, and print a graph and data table.

Two chemical reactions will be studied, both of which are exothermic:

(1) Mg(s) + 2 HCl(aq) MgCl2(aq) + H2(g) ΔH1

(2) MgO(s) + 2 HCl(aq) MgCl2(aq) + H2O(l) ΔH2

The amount of heat evolved will be calculated from the temperature rise of the aqueous

solution inside the calorimeter. From the amounts of Mg and MgO reacting, ΔH1 and

ΔH2 will be determined. From these results, ΔH for the following reaction, difficult to

determine experimentally, will be calculated by application of Hess's law:

(3) Mg(s) + H2O(l) MgO(s) + H2(g) ΔH3

In this experiment, the heat absorbed by the calorimeter will be taken into consideration:

heat evolved by reaction = heat absorbed by H2O + heat absorbed by calorimeter

qrxn = -(mH2OcH2OΔT + CcalΔT) (1)

where Ccal is the total heat capacity of the styrofoam cup part of the calorimeter, mH2O is

the mass of water, cH2O is the specific heat of water, and ΔT is the temperature change.

The value of Ccal will be determined by adding a measured mass, mH2O, of hot water at Th

to the same mass, mH2O, of cold water in the styrofoam cup. The cold water and cup are

both at Tc. Heat will transfer from the hot water to the cold water until they reach the

same final temperature, T2.

heat lost by hot water = heat gained by cold water and styrofoam cup

Page 98

-mH2OcH2O(T2 - Th) = +mH2OcH2O(T2 - Tc) + Ccal(T2 - Tc) (2)

The value of Ccal can be calculated from this relationship.

PROCEDURE

SAFETY

Wear your safety goggles as usual. Be careful to keep the cables from the temperature

probes away from the hot plate or the flame of the Bunsen burner!!! They are quickly

destroyed by excessive heat! In Part B and Part C you will be working with acidic

solutions and in Part B a flammable gas, H2, is evolved. Avoid breathing this gas, and

keep it away from flames and heaters.

Carefully measure all volumes, masses and temperatures (read to the estimated limit on

all scales) throughout this experiment!

Part A Determination of the Calorimeter Constant

1. Assemble a coffee cup calorimeter. Use two nested 8 oz. clean dry foam plastic cups,

placed in a 400 mL glass beaker, with a cardboard cover lid which has a hole in the

center. Set up LoggerPro with the Vernier Interface and two stainless steel

temperature probes (using CH1 and CH2). (Double click on the Logger-Pro icon on

the desktop. Change the time interval for data collection by going to EXP DATA

COLLECTION and changing the time to a longer interval, e.g. 800 seconds).

2. Using a graduated cylinder, carefully measure 50.0 mL of water into the calorimeter.

Place the temperature probe connected to CH1 through the hole in a cardboard cover

and into the water in the calorimeter.

3. In a drained wet 150 mL beaker, heat another carefully measured 50.0 mL of water to

about 70oC. Observe its temperature with the temperature probe connected to CH2.

BE CAREFUL TO KEEP THE CABLES CONNECTED TO THE

TEMPERATURE PROBES AND ALL OTHER CABLES AWAY FROM THE

HOT PLATE OR THE FLAME OF THE BUNSEN BURNER!!! (They are

quickly destroyed by high temperatures!)

4. Pour all of the hot water from the 150 mL beaker into a third drained wet Styrofoam

cup and also transfer the temperature probe (connected to CH2) with the hot water in

the Styrofoam cup. Beware, this cup with temperature probe easily tips over, so you

may find it helpful to place this cup in a holder such as a glass beaker to give a little

more stability.

5. Begin data collection by clicking on "Collect". Gently stir both the hot and cold

water in the Styrofoam cups with the temperature probes in their respective cups

(cold still connected to CH1 and hot still connected to CH2) and watch for the


temperatures (both hot and cold) to be holding approximately constant (probably

within about 30 seconds). (The hot water temperature will actually be very gradually

decreasing.)

6. When the two temperatures are holding approximately constant (not the same

temperature, but each separately constant), then quickly pour all of the hot water from

the Styrofoam cup into the cold water calorimeter and gently swirl and stir the

hot/cold mixture with the temperature probe (still connected to CH1) through the hole

in the lid on the calorimeter. Collect data a minute or so after mixing the hot and cold

water with the continuous gentle stirring.

7. When the temperature of the hot/cold mixture becomes constant (actually very slowly

decreasing), you may then stop collecting data by clicking on "Stop".

8. Look at the data you have collected and record the constant temperature of the cold

water just before mixing. Similarly record the best almost constant temperature of the

hot water just before you started pouring. Excepting for a brief spike, which might

appear, record the highest constant (almost) temperature of the hot/cold water mixture

observed soon after pouring together.

9. As an option, you may store your data by clicking on the "Data" Menu (top left of screen) and

then click on "Store Latest Run" in the Data Menu.

10. Drain the calorimeter, and wipe the inside dry.

11. Make at least three repeat runs (at least four runs total)

12. Perform the calculations for the calorimeter constant. Note that if the heat given up

by the hot water is less than the heat gained by the cold water, the calorimeter

constant will appear to be negative! In this experiment, since the mass, mH2O, of the

hot water and the cold water are the same (almost) and we assume the specific heat,

cH2O, of the hot and cold water are the same, a quick check on the experimental

results can be made by comparing the temperature change of the hot water, (T2 - Th),

and the temperature change of the cold water, (T2 - Tc). If the temperature change of

the hot water, (T2 - Th), is less than the temperature change of the cold water, (T2 -

Tc), that is (T2 - Th) < (T2 - Tc), then you will determine a negative calorimeter

constant, Ccal. Calorimeters usually do not have negative calorimeter constants, Ccal.

13. If you get negative values for Ccal, you should rerun the calorimeter constant

determinations. If you are limited by time, it may be possible to make additional

calorimeter constant runs after you finish the other parts of this experiment on

another day.

Very carefully measure all volumes, masses and temperatures (read to the estimated

limit) throughout this experiment!

Page 100

Part B Magnesium Plus Hydrochloric Acid

1. Carefully measure out 100.0 mL of 1.0 M HCl(aq) into the clean dry calorimeter. You

may disconnect the temperature probe connected to CH2. Place the temperature

probe connected to CH1 through the hole in the cardboard cover and into the solution.

The calorimeter should now be covered with the cardboard lid and the temperature

probe should be immersed in the solution. Set this calorimeter aside.

2. Obtain a piece of magnesium, Mg, metal, in ribbon form, about 30 cm long. Because

such ribbon is usually covered with an oxide coating, it should be polished before

weighing. This can be done by stretching the ribbon on top of a piece of cardboard

and rubbing it with a piece of emery cloth. Polish both sides of the ribbon, and wipe

it free of dust. With minimum handling, cut the ribbon into lengths shorter than 2 cm.

Weigh these strips by placing plastic weighing dish on the balance and tare the

balance. Then measure the mass of the magnesium (you should have about 0.25 to

0.30 g of magnesium).

3. Begin data collection by clicking on "Collect". Gently stir the solution in the

calorimeter with the temperature probe and watch for the temperature to be holding

about constant. When the temperature is constant (probably within about

15 seconds), add the Mg strips and replace the cover. Swirl and stir the calorimeter

gently with the temperature probe (still connected to CH1). The temperature will

start to rise. Avoid breathing the gas (H2) evolved.

4. Collect data as the temperature rises to a nearly constant maximum and then begins to

very gradually drop to lower temperatures. You may then stop collecting data by

clicking on "Stop".

5. Look at the data you have collected and record the constant temperature of the HCl(aq)

solution just before adding the magnesium metal. Similarly record the constant (almost)

highest temperature (ignore very brief spikes) of the solution after adding the Mg.

6. You may store your data by clicking on the "Data" Menu (top left of screen) and then click on

"Store Latest Run" in the Data Menu.

7. Discard the contents into the waste container. (There should be solution only and not

any unreacted Mg metal remaining in the calorimeter!). Rinse the calorimeter with

water and wipe it dry.

8. Repeat the complete determination, starting with another carefully measured 100.0

mL of 1.0 M HCl(aq) and weighed cleaned Mg.

Part C Magnesium Oxide and Hydrochloric Acid

1. Carefully measure out 100.0 mL of 1.0 M HCl(aq) into the clean dry calorimeter.

Place the temperature probe connected to CH1 through the hole in the cardboard


cover and into the solution. The calorimeter should now be covered with the

cardboard lid and the temperature probe should be immersed in the solution. Set this

calorimeter aside.

2. Weigh a plastic weighing dish. Add about 0.5 g of MgO to the dish, and carefully

reweigh it to ±0.001g. Reseal the bottle of MgO promptly. (After pouring out the

MgO from the weighing dish, the dish will need to reweighed, because some of the

MgO may stick to it, Step #7 below.)

3. Begin data collection by clicking on "Collect". Gently stir the solution in the

calorimeter with the temperature probe and watch for the temperature to be holding

about constant. When the temperature is constant (probably within about 15 seconds),

dump in the MgO rapidly but carefully so that it does not stick to the sides of the

calorimeter. Set the emptied weighing dish aside to reweigh later. At once, swirl the

cup gently and stir the calorimeter with the temperature probe (still connected to CH1)

to keep the solid from caking on the bottom. The temperature will start to rise.

4. Collect data as the temperature rises to a nearly constant maximum and then begins to

very gradually drop to lower temperatures. You may then stop collecting data by

clicking on "Stop".

5. Look at the data you have collected and record the constant temperature of the HCl(aq)

solution just before adding the magnesium oxide. Similarly record the constant

(almost) highest temperature (ignore very brief spikes) of the solution after adding the

MgO.

6. You may store your data by clicking on the "Data" Menu (top left of screen) and then

click on "Store Latest Run" in the Data Menu.

7. Reweigh the emptied plastic weighing dish.

8. Before discarding the contents of the calorimeter, check to see that no solid is left. (If

there is, you may want to change your technique slightly on the next trial and not use

the results of this run.) All of the solid MgO should be reacted with the HCl(aq).

9. Rinse the calorimeter and dry the inside.

10. Repeat the complete determination, starting with another carefully measured 100.0

mL of 1.0 M HCl(aq) and weighed MgO.

Page 102

DATA





Part A: Calorimeter Constant Run: 1 2 3 4

1. Temperature of 50.0 mL cool water (Tc) ______ ______ ______ _______

2. Temperature of 50.0 mL warm water (Th) ______ ______ ______ _______

3. Maximum temperature on mixing (T2) ______ ______ ______ _______

4. Temperature change of hot water,

(T2 - Th) ______ ______ ______ _______

5. Temperature change of cold water,

(T2 - Tc) ______ ______ ______ _______

6. Heat lost by hot water,

mH2OcH2O(T2 - Th) ______ ______ ______ _______

7. Heat gained by cold water,

+mH2OcH2O(T2 - Tc ______ ______ ______ _______

8. Heat transferred ("lost") to

calorimeter, qcal, ______ ______ ______ _______

9. Total heat capacity of the calorimeter,

qcal/ΔTcal = Ccal, in joules per degree, ______ ______ ______ _______

where ΔTcal = (T2 - Tc). Note: If you get negative values for Ccal, you should rerun

these calorimeter constant determinations after you finish the other parts of this

experiment.

10. Average: __________

(Note: You should repeat runs for the calorimeter constant until you get positive values

for Ccal. If your results continue to be negative, consult your instructor. (In place of

continually negative results, you may be instructed to use Ccal = 10 J/oC.)

Part B: Mg plus HCl(aq) Trial 1 Trial 2

1. Mass of weighing dish __________ __________

2. Mass of weighing dish + magnesium __________ __________


3. Mass of magnesium __________ __________

4. Initial temperature of acid sol'n (T1) __________ __________

5. Final temperature (maximum, T2) __________ __________

6. Heat absorbed by the solution, msolncH2OΔT,

(assume 100.0 g, with c = 4.18 J/goC) __________ __________

7. Heat absorbed by the calorimeter, CcalΔT __________ __________

8. Total heat evolved by reaction, qrxn __________ __________

9. Moles of magnesium __________ __________

10. ΔH (kJ/mol Mg) __________ __________

11. Average __________

Part C MgO plus HCl(aq)

Trial 1 Trial 2

1. Mass of weighing dish __________ __________

2. Mass of weighing dish plus MgO __________ __________

3. Mass of weighing dish, after using MgO __________ __________

4. Net mass of MgO taken for reaction __________ __________

5. Initial temperature of acid solution (T1) __________ __________

6. Final temperature (maximum, T2) __________ __________

7. Heat absorbed by the solution, msolncH2OΔT,

(assume 100.0 g, with c = 4.18 J/goC) __________ __________

8. Heat absorbed by the calorimeter, CcalΔT __________ __________

9. Total heat evolved by the reaction, qrxn __________ __________

10. Moles of MgO __________ __________

11. ΔH (kJ/mol MgO) __________ __________

12. Average ΔH __________

Page 104

Part D Application of Hess's Law

a.Write in the experimentally derived energy terms in the following thermochemical

equations. (These are the results from your data above.)

(1) Mg(s) + 2 HCl(aq) MgCl2(aq) + H2(g) ΔH1 = _________

(2) MgO(s) + 2 HCl(aq) MgCl2(aq) + H2O(l) ΔH2 = _________

b.Use Hess' law to get ΔH3 for reaction (3) below from the experimental ΔH1 and ΔH2

entered immediately above. (This ΔH3 is based on your experimental data in (1) and (2)

above.) Show calculations.

(3) Mg(s) + H2O(l) MgO(s) + H2(g) ΔH3 = _________

c. Now do an independent (separate) calculation of ΔH3 from the standard molar

enthalpies of formation, ΔHfo, found in a reference table elsewhere (e.g. Handbook of

Chemistry and Physics). Note that you do NOT use your experimental data in this

calculation. Show calculations. Draw a box around your answer.

d.How does the theoretical ΔH3 (step c. above) compare with your experimental ΔH3

(step b. above)?


POST-LAB Name_________________


Experiment 13 Determination of Heat of Reaction

1. Various assumptions and simplifications were made in this experiment that reduce

precision. However, considering the greater uncertainty in measuring temperature,

they are probably excusable. But now tell how these ignored errors would tend to

affect the value of ΔH obtained:

(H increases, decreases,

or no effect)

a. Not all of the "100 mL" of acid entered the

calorimeter.

b. The density of the acid solution is not exactly 1.00

g/mL, but slightly higher.

c. The specific heat of the acid solution is less than the

assumed 4.18 J/goC.

d. The "maximum temperature" observed is really lower

than if we had made a time-temperature plot to correct

for heat loss.

e. The thermometer itself has a finite heat capacity; i.e.,

it absorbs some heat.

2. In Part D, Line c, you calculated ΔH3 from the standard heats of formation (standard

molar enthalpies of formation), ΔHfo, taken from a reference table (e.g. Hand Book of

Chemistry and Physics). Now do the same for ΔH1 and ΔH2. (Again note that you do

NOT use your experimental data in this calculation.) Show both calculations below.

ΔH1

ΔH2

How can ΔH1 and ΔH2 be combined to give ΔH3? That is show the simple mathematical

calculation that gives the value of ΔH3 from the values of ΔH1 and ΔH2.

What thermo-chemical law provides the basis for what you are doing here in the simple

mathematical calculation?

Page 106


PRE-LAB Name_________________


Experiment 13 Determination of Heat of Reaction

1. Calculate the final temperature for each of the following mixtures. Assume no heat

loss to the environment. Show your calculations.

a. 50.0 g of water at 25.0oC is mixed with 50.0 g of water at 85.0

oC.

b. 75.0 g of water at 25.0oC is mixed with 25.0 g of water at 85.0

oC

2. a. In Part B of this experiment, 0.20 g of Mg is added to 100 mL of 1.0 M HCl(aq).

Which is the limiting reactant? Show calculations.

b. In Part C, 0.50 g of MgO is added to 100 mL of 1.0 M HCl(aq). Which is the

limiting reactant?

3. When 50.0 mL of 1.00 M H2SO4(aq) at 26.1oC was added to 50.0 mL of 1.00 M NaOH

also at 26.1oC, the temperature rose to 32.6

oC. Assume the resulting solution had a

total volume of 100.0 mL with the same density and specific heat as water. Calculate

ΔH for the reaction described by this equation.

H2SO4(aq) + 2 NaOH(aq) Na2SO4(aq) + 2 H2O(l)

To do this calculation, work through the following questions and steps (a-f):

a. Which reactant limits?

b. How many moles of it are present?

Page 108

c. How much energy is required to heat 100.0 g solution from 26.1oC to 32.6

oC?

d. How much energy is released by the reaction per one mole of the limiting

reactant?

e. How many moles of that reactant are in the given equation?

f. What is ΔH for the reaction as written above in kJ (not kJ/mol)?


Experiment 14

Determination of Crystal Violet by Spectrometry


PURPOSE

In this experiment you will learn to use a spectrophotometer to determine the molar

concentration of Crystal Violet in an unknown solution.

Method

The principles of light absorbance and the theory behind spectrometry are discussed

elsewhere.

chemicals

1.00 x 10-4

M Crystal Violet stock solution

equipment

Vernier Logger Pro Spectrometer and cuvettes

50 mL volumetric flasks

10 mL graduated pipet

Safety

Crystal Violet solutions can be irritating to the eyes, skin and clothing. As usual, use eye

protection while handling these solutions.

PROCEDURE

Part 1. Preparing the Standard Crystal Violet Solutions

a. Five solutions will be prepared below (as directed by instructor). The total volume for

each solution will be 50.00 mL.

Standard Solution Number 1 2 3 4 5

Volume of 1.00 x 10-4

M

Crystal Violet in mL

5.00

4.00

3.00

2.00

1.00

b. For each solution, thoroughly clean a 50 mL volumetric flask, including a final rinse

with deionized water.

c. Thoroughly clean a 10 mL graduated pipet*, again rinsing with deionized water.

Obtain about 10 mL of the 1.00 x 10-4

M Crystal Violet solution in a clean, dry 50 mL

beaker. Rinse the pipet with 2 or 3 small portions of the Crystal Violet solution.

Discard the rinsings and any solution in the beaker.

Page 110

* Graduated pipets, unlike volumetric pipets, require that you start and stop the solution

level at the desired marked line.

d. Obtain a fresh portion of about 30 mL of the 1.00 x 10-4

M Crystal Violet solution in

the 50 mL beaker. Using the pipet just rinsed with the Crystal Violet solution, pipet* the

assigned (1. a. above) volume into the 50 mL volumetric flask. Label the flask with the

appropriate solution number (see 1. a. above)

e. Add deionized water to the flask until the solution level is almost to the mark.

Stopper the flask, and mix the contents by inverting the flask 8 or 10 times. Be sure to

hold the stopper firmly! Allow solution to drain down the neck of the flask and do not

lose any of the solution as you remove the stopper at this point! Using a transfer pipet,

add deionized water to the mark. Again stopper the flask and thoroughly mix by

inverting the flask several times.

f. Prepare solutions 2, 3, 4, and 5. If there are not enough 50-mL volumetric flasks,

transfer solution 1 into a clean container and reuse the same 50-mL volumetric flask.

Do the same for the other solutions.

g. For your report, you will need to know the concentration of each of the 5 diluted

solutions that you have prepared. Calculate the concentration of each diluted solution

and enter into the table (see Data section below). Show your calculations in your

notebook.

Part 2. Determining the Absorbance Curve of the Crystal Violet solution.

a. Prepare the Vernier Spectrometer by plugging in the USB cable and opening the

Logger Pro software. If the software doesn’t immediately recognize the Spectrometer,

choose Connect Interface Spectrometer Scan for Spectrometers from the

Experiment menu. Allow the Spectrometer to warm up for 3 minutes before taking

readings.

b. Prepare a blank by filling a cuvette 3/4 full with distilled water. To correctly use a

cuvette, remember: All cuvettes should be wiped clean and dry on the outside with a tissue.

Handle cuvettes only by the top edge of the ribbed sides.

All solutions should be free of bubbles.

c. Calibrate the spectrometer by choosing Calibrate Spectrometer from the

Experiment menu. Follow the instructions from the dialog box to complete the

calibration using your blank cuvette. You will be asked to insert the blank cuvette into

the cuvette slot. Insert it in such a way that the spectrometer light goes through the

smooth sides and not the ribbed sides of the cuvette. Click “OK.”


d. Rinse a sample cuvette with two or three small portions of the #1 diluted Crystal

Violet solution from its volumetric flask. Fill the cuvette about ¾ full with the #1

diluted Crystal Violet solution. Place the cuvette in the cuvette slot. Click “Collect”. An

absorbance curve should appear on the screen. After viewing the absorbance curve, hit

“Stop”.

e. Click on the “Configure Spectrometer Data Collection” icon, located on the right hand

side of the toolbar to open the display. (The button looks like a rainbow graph.) Click

Abs. vs. Concentration (under Set Collection Mode). The wavelength of the maximum

absorbance will be automatically selected. Double check that 460 nm or somewhere

close (this is the max) is the only wavelength that is selected. Record this max. Click

“OK” to close the display.

Part 3. Measuring the absorbance of each standard solution and your unknown

a. With the first Beer’s law standard solution in the cuvette slot (#1), record the

absorbance value shown on the lower left-hand corner of the screen (the last digit may

fluctuate so do your best to find the average).

b. Transfer a portion of solution #2 into a clean cuvette, pre-rinsed with 2 or 3 small

portions of the solution, and insert into the cuvette slot. Record the absorbance. Repeat

this step for the remaining standard solutions.

c. Transfer a portion of your assigned unknown solution into a clean cuvette, pre-rinsed

with 2 or 3 small portions of the unknown solution. Place the cuvette in the sample

holder. Record the absorbance.

d. Pour out all Crystal Violet solutions into the waste container provided and rinse the

containers with water.

Part 4. Determination of the Concentration of an Unknown Crystal Violet Solution

a. You will use the Excel program to make your Beer’s Law Plot. Display the trend

line, the equation, and the R2 value. Make sure to set the y-intercept at zero. Use a

proper title for each axis and for the plot. Make sure to print your Beer’s Law Plot (or if

your instructor desires, send by e-mail).

b. Use the Beer’s Law Equation generated from your plot to determine the concentration

of your unknown.

Page 112

DATA





1) Table of Concentrations and Absorbance Values λmax __________

Solution Number 1 2 3 4 5

Volume of 1.00 x 10-4

M

Crystal Violet in mL

5.00

4.00

3.00

2.00

1.00

Concentration (M)

Absorbance

Unknown Code Absorbance of the

unknown solution

Calculated

concentration of the

unknown solution

2) Attach your Beer’s Law plot (with the equation displayed) to your lab report (or e-mail

to instructor if told to do so).

CALCULATIONS

1) Show the calculation of the concentration of each standard.

2) Using the equation determined from your Beer’s Law plot, calculate the concentration of

your unknown sample.


PRE-LAB NAME_____________________ (To be completed before coming to lab)

Experiment 14 Determination of Crystal Violet by Spectrometry

1. For this question, assume that you have been assigned solution number 3 to prepare.

Calculate the concentration of Crystal Violet in this solution.

2. A student has found the following %T readings for solutions of the molarities listed

below. All measurements were done at the wavelength of maximum absorbance. For each

%T reading, calculate the absorbance and enter the value in the table below.

(A = 2 – log %T )

%T Molarity Absorbance, A

40.9 1.6 x 10-4

_____________

50.5 1.3 x 10-4

_____________

59.0 9.6 x 10-5

_____________

70.4 6.4 x 10-5

_____________

3. Beer's law: A = εbc, where A = absorbance, ε = molar absorptivity, b = path length, and

c = concentration in M.

Plot the absorbance, A, (y-axis) versus concentration, c, (x-axis) using Excel. Print your

graph and attach to this form (Check with instructor if e-mail is OK).

4. Display the trend line for this linear relationship and the equation for the line.

5. Again following the procedure of the experiment, the same student found a %T of 56.4

for the unknown. Find the concentration of the unknown solution in two ways.

a. First, read the concentration value off the graph. Show how you read the unknown value

on the graph on the next page using dashed lines (----).

b. Second, calculate the concentration value using the slope of your line and Beer's law:

A = εbc, where A = absorbance, ε = molar absorptivity, b = path length, and

c = concentration in M. (Note that εb is the theoretical slope of the line.)

c. What is the theoretical value of the y-intercept (absorbance, A) on a Beer's Law plot?

Page 114


Experiment AA

Measurement of Iron by Atomic Absorption (AA) Spectrometry

Prelab: Complete the prelab questions on page 121 before lab.

PURPOSE

The purpose of this experiment is to learn the technique of atomic absorption

spectrometry. The AA instrument will be used to detect and measure the iron content of

an unknown sample.

Method

Your instructor will discuss the general theory behind atomic absorption spectrometry

and how the instrument works.

Acid will be used to dissolve the iron in the unknown sample. A calibration curve will be

prepared using a series of standard Fe3+

solutions whose absorbances will be measured

using the AA instrument. The absorbance of iron in an unknown sample will be

measured using the AA instrument and the concentration calculated using the calibration

curve.

Reference: D.T. Sawyer, W. R. Heineman, and J. M. Beebe, “Chemistry Experiments for

Instrumental Methods,” John Wiley and Sons, New York. 1984.

PROCEDURE

A. Preparation of the Sample to be Analyzed

If your unknown sample is a solid:

1) Rinse a clean 250-mL flask 3 or more times with deionized (DI) water. Do not insert

a bottle brush into your flask.

2) Measure about 2 grams of the unknown sample and transfer it into the clean 250 mL

flask. Record the exact mass.

3) Add 25 mL of 8 M HCl (aq) to the flask and boil slowly on a hot plate for 5 minutes.

Avoid drying out!

4) Cool the mixture and then add 10 mL of deionized water.

Safety

Be careful when boiling the sample in 8M HCl (aq). Make sure that the fume

hood is over your heating apparatus. Do not leave your flask unattended and do

not let it dry out.

The AA instrument uses an acetylene flame. Avoid getting too close to this part of

the instrument.

Page 116

5) Rinse a clean glass funnel, a 125-mL erlenmeyer flask, and a 50-mL volumetric flask

3 or more times with deionized water.

6) You will carry out 3 levels of filtration. Line your glass funnel with each type of

filtering material and make sure to rinse well between each filtration. The first

filtrations should be done into an Erlenmeyer flask (ask for an extra one if you run

out). The last filtration should be done into the 50 mL volumetric flask.

a) Using a glass funnel, filter first with a double layer of cheesecloth into a clean

125-ml flask to remove the largest pieces.

b) Next, filter the mixture through a “Fast” filter paper into another clean

erlenmeyer flask. (If your filtration is taking a while, proceed to Part B and

come back to this).

c) Next filter the mixture through a Whatman No. 1 filter paper into the 50-mL

volumetric flask.

7) Once all of the liquid has filtered through (this is what you will analyze), make sure

that there are absolutely no solid particles in the solution. If there are solid particles

still left, re-filter the liquid.

8) Dilute to the mark with deionized water. You will use this solution directly for the

absorbance measurement.

B. Preparation of the Standard Iron Solutions:

1) Four solutions of known concentrations of iron will be prepared by dilution from the

1000 mg/L stock solution:

Standard 1 Standard 2 Standard 3 Standard 4

Concentration of

Standard Solution

(mg/L or ppm)

10.00 30.00 50.00 100.0

Calculate in your notebook the volume of the stock solution you will need to dilute for

each standard. Show your calculations to your instructor before proceeding.

2) Thoroughly rinse 4 100-mL volumetric flasks with deionized water. Do not use a

bottle brush! Label each one with the standard number and the concentration.

3) Using a 10-mL graduated pipet, transfer the required volume of the stock solution

into each of the 100-mL volumetric flask.

4) Dilute to the mark with DI water. Cover with parafilm or stopcock.

C. Absorbance Measurements

1) Your instructor will demonstrate how to use the AA instrument.

2) Take your standard solutions and the unknown solution to the AA instrument.

3) The instrument is ready to use when the flame is on. Keep away from this area of

the instrument. The aspirator tube should be sitting in and aspirating DI water

when no samples are being tested.

4) You will first have to do a blank calibration by aspirating the “blank” sample.

Gently transfer the aspirator to the container labeled “blank”. Avoid pinching the


fragile aspirator tube. Make sure that the tip of the aspirator is completely below

the liquid surface.

5) Using the pen stylus, click on the “Analyze Blank” box. As the instrument proceeds

with the multiple measurements, you can watch the progress in the window on the

upper right hand corner of the screen. (See sample screen shot on the next page).

6) When the instrument is done with the blank calibration, your screen should say

“Autozero performed”.

7) Aspirate your first standard by gently transferring the aspirator into Standard 1

flask. Make sure the tip is in the solution.

8) The concentrations of your standards should have been entered by your instructor.

Next to the “Analyze Standard” button, make sure that the correct Standard # and

concentration are showing.

9) Using the pen stylus, click on “Analyze Standard”. Watch the “progress” box on

the upper right hand corner. The measurement is done when the screen says

“Standard 1 applied”:

10) Record the absorbance value into the data table in your notebook.

11) Insert the aspirator back into the DI water beaker to clean the aspirator for a few

seconds.

12) The window next to “Analyze Standard” should say Standard 2 with the correct

concentration.

13) Transfer the aspirator from the DI water beaker into the Standard 2 flask. Click on

“Analyze Standard” as before and wait for the “Standard 2 Applied” message to

come up. Again, record the absorbance.

14) Follow the same steps for Standards 3 and 4, remembering to rinse the aspirator

with DI water before each standard analysis and recording the absorbance

measurement. At the end of the standards analysis, leave the aspirator in the DI

water beaker.

15) After the absorbance measurements are done for the 4 standards, you will now

measure the absorbance and concentration of iron in your unknown solution.

16) Do one last check to make sure that there are absolutely no solids in your sample.

This will clog the very tiny aperture for the nebulizer!

17) Transfer the aspirator into your unknown sample flask. Click on “Analyze Sample”.

Page 118

18) When the sample analysis is done, record the concentration of iron in your sample.

19) If the concentration is higher than that of the most concentrated standard, dilute the

sample ten times and repeat the analysis.

20) Remove the aspirator from your flask and gently transfer back into the DI water

beaker for the next person.

D. Determination of the Absorbance of Iron in Your Unknown.

1) Prepare your calibration curve. Using the Excel program, make a plot of the

absorbance versus standard concentrations.

2) Using the “Add Trendline” function (under the Chart menu), do a linear regression

analysis.

3) In the same “Add Trendline” window, click on option and check the boxes to set the

intercept to zero and display the equation and the R2 value.

4) Use the equation to calculate the absorbance of the iron in your unknown.

5) Print the chart with the correct labels and staple to your report.

DATA





1) Description of the unknown sample (brand, type of cereal, size, texture, etc.)

2) % RDA or DV of iron per serving __________ and serving size in grams _______

3) Mass of the unknown sample


4) Concentration of stock solution

5) Data table below:

Standard 1 Standard 2 Standard 3 Standard 4

Concentration of

Standard Solutions

(mg/L or ppm)

10.00 30.00 50.00 100.0

Volume of stock

needed to make 100

mL of each standard

solution

Absorbance of each

standard

Concentration of iron

in unknown from AA

result

Calculated

absorbance

of unknown

CALCULATIONS

1) Calculations for the volumes needed to prepare the standard solutions

2) Calculation of the absorbance of unknown iron solution

3) Calculate the mass of iron in 1 gram of your food sample using your concentration

result for the unknown.

4) Calculate the mass of iron in 1 g according to the label nutritional information using

% of RDA or DV per serving. Assume an RDA of 15 mg.

5) Do these two values agree with each other? Give some reasons as to why they may

be different.

Page 120


PRE-LAB Name_________________


Experiment AA Measurement of Iron by Atomic Absorption (AA)

Spectrometry

1) Using an acceptable source, look up some information on how an atomic absorption

spectrophotometer works (online or from a text or journal article). Briefly summarize

what you learned below:

2) Explain in your own words how you would prepare Standard 1. Show your

calculation for the volume required of the stock solution with a concentration of 1.00

x 103 mg/L.

3) Using the sample results below, calculate the concentration of iron in a sample that

has an absorbance of 0.72.

Page 122


Experiment 15

Group Relationships and Periodic Properties


PURPOSE

In this experiment we will

• find evidence of similar chemical behavior of elements that have a vertical

arrangement in the periodic table, and

• find progressively different behavior of elements that have a horizontal arrangement

in the periodic table.

Method

Elements that have a vertical arrangement in the periodic table are, of course, members of

the same family or group. Just as in other families we know, all members are not exactly

alike, but they often have much in common that distinguishes them from members of

other families.

Elements that have a horizontal arrangement in the periodic table are members of the

same period. An analogy to the progressive differences in these elements is the ecological

differences you might observe traveling across the state of California due east from the

Pacific coast. Whether you started the trip at Crescent City, Santa Cruz or Santa Barbara,

you would find a sequence of changes that seemed familiar.

For our brief survey of the periodic table we will focus on a dozen or so elements:

IA IIA IIIA IVA VA VIA VIIA

2nd Li C O

3rd Na Mg Al P S Cl

4th K Ca Br

5th I

The chemical behavior of these elements will be investigated in the following categories:

I. Reactivity with water

II. Reactivity with acids

III. Study of oxides: basic and acidic

IV. Redox behavior of the halogens.

Because we are not trying to "reinvent" the periodic table on the basis of empirical

evidence, it will be quite appropriate to relate degree of reactivity to the knowledge of

electronic structure we take from the textbooks. If you know how—and why—ionization

energy varies in a left-to-right direction and in a top-to-bottom direction within the

Page 124

periodic table, you already have a powerful tool for predicting and explaining many

periodic trends.

PROCEDURE Work with one, and only one, partner.

Note: Sprinkled throughout the procedure are numerous questions, labeled Q-1, Q-2,

etc. Answer these questions in the appropriate areas of the Report Sheet. However, be

sure that you have collected all of your data before you begin to answer these questions.

Do not repeat the question, but please do use complete sentences.

I. Reactivity with water

A. Three alkali metals: lithium, sodium and potassium.

Prepare three beakers. Put about 100 mL of deionized water into each. Add 4 or 5

drops of phenolphthalein solution, an acid-base indicator. Then cover each beaker

with a wire screen from your locker.

Notice that these metals are stored under oil to prevent exposure to air and moisture.

Even so, your instructor will probably have to cut off an oxide crust to get a sample

of the soft, shiny metal. Bring your instructor 3 labeled watch glasses. You will be

given a piece the size of a small bean. Return to your desk and, with both partners

watching, drop each piece into a different beaker and immediately replace the wire

screen. (A screen is used instead of a watch glass because with the screen there is

less likelihood of accumulating an explosive mixture of gases.)

Q-1. Compare the relative reactivity of these three metals. What would you

predict for the relative reactivity of rubidium? In other words, rank these

four metals.

Q-2. What is the significance of the color change?

Q-3. Note that the spherical appearance of the sodium suggests that it is now a

liquid. Look up the melting point of sodium. Suggest why the sodium has

melted.

Q-4. Why does the sphere not sink, at least part way, into the water? Give two

reasons.

SAFETY FIRST

Wear your safety goggles throughout the experiment. Access to lithium, sodium and

potassium will be supervised by your instructor, but be aware that these metals react

explosively with water if not properly controlled. Other safety messages will be given

later in this procedure.


B. Sodium and methyl alcohol

There is no water present here, but methyl alcohol can be thought of as derived

from water. Caution: methyl alcohol is flammable; keep flames away.

Obtain a small piece of sodium metal from your instructor. Drop it into a small

beaker containing about 20 mL of methyl alcohol.

Q-5. How does this reaction compare with that of Na and H2O?

Q-6. Why does the sodium not "float" this time? Two reasons.

C. Two alkaline earth metals: magnesium and calcium

Place about 150 mL of water into each of two 250 mL beakers. Into one beaker put

a short (1 inch) strip of magnesium ribbon. Observe the behavior of the magnesium

periodically while you work with the calcium.

Prepare the second beaker for the collection of a gas as in Figure 1. Fill a large test

tube with water. Put your thumb over the end, and invert the tube into the beaker so

that little if any air is trapped at the top of the tube. (Repeat the operation until you

succeed.) Now dry your hands, take a dry watch glass to the bottle marked "Fresh

Calcium metal," and collect from 1 to 5 granules as directed. Follow any other

instructions provided by the laboratory technician on the card in the gray bin.

Figure 1

Return to the beaker. Raise the test tube only enough to trap the pieces of calcium

under the mouth of the tube. Try to collect most of the evolved gas. It may be

necessary to repeat the operation in order to obtain a full test-tube of the gas.

When the tube is full of gas, remove it, keeping it in a vertical position, mouth

down. Then, some distance from the beaker, have your partner bring a lighted

splint to the mouth of the test tube.

Page 126

When the reaction in the beaker has ceased, note the milky appearance of the

liquid. Add a few drops of phenolphthalein indicator solution and note any color

change. Now add some dilute hydrochloric acid dropwise to decolorize the

indicator. What happens to the milkiness at the same time?

Dispose of the contents of the beaker in the proper waste container. (Save the

beaker with the magnesium in it for IIA below.)

Q-7. A "barking" sound when the gas explodes is characteristic of hydrogen.

What would happen to the splint if the gas had been (a) oxygen? (b)

carbon dioxide?

Q-8. What would have happened if you had carried the tube to the flame with the

mouth up?


A. Magnesium

Place about 5 mL of 3 M HCl in a medium sized test tube. Set this test tube upright

in one of the holes in your test tube rack. The idea is to be able to invert a large test

tube over the top of the medium test tube.

Add the magnesium from the previous experiment to the smaller test tube and

immediately invert a large test tube over it. Collect the gas evolved until the

magnesium dissolves. Carefully lift up the top test tube, keeping it vertical. Bring a

match or a burning splint to the mouth of the test tube. Identify the gas.

test tube rack

Mg and acid

Figure 2

Dispose of the contents of the test tube in the proper waste container.


Q-9. Which is easier for a metal to replace: hydrogen from water or hydrogen

from an acid? What do you think would happen if you added sodium to an

acid?

B. Aluminum

Place a small strip of aluminum in a large test tube. Add about 4 mL of 6 M HCl.

Set it aside. If there is no reaction in five minutes, set the tube in a beaker of hot

water. Is any gas evolved?

III. Study of oxides: basic and acidic

Oxides of elements of the third period (plus carbon) will be taken to illustrate

"horizontal" trends in the periodic table. The source of the oxides will be as follows: A.

sodium: sodium peroxide, a commercial product, will be substituted, because the oxide,

difficult to make, converts to the peroxide; B. magnesium: a piece of Mg ribbon will be

burned in air; C. aluminum: the oxide presents problems because of its insolubility, so we

will study the hydroxide which is closely related; D. carbon: a piece of charcoal will be

ignited and then thrust into pure oxygen; E. phosphorus: a small sample of red

phosphorus will be ignited in an iron spoon and thrust into pure oxygen; F. sulfur: same

procedure; G. chlorine: the oxides are dangerously explosive, so the safer reaction

product of oxide with water will be substituted.

A. Sodium

Drop about 0.2 g of sodium peroxide, Na2O2, into a large test tube containing bout

10 mL of deionized water. Set it to one side on your desk while oxygen evolves.

Then test the solution with litmus paper. (In making tests with litmus paper, use a

stirring rod to bring a drop of the solution to the paper. Do not dip the paper into the

solution.)

Dispose of the solution in the proper waste container.

B. Magnesium

Obtain a short piece of Mg ribbon. Put a little deionized water (about 1 mL or less)

in an evaporating dish. Using your tongs, hold the Mg in a gas flame long enough to

ignite it, and then avert or shield your eyes. Let the white "ash" fall from the tongs

into the water. Stir the mixture of solid and water from time to time, and test with

red litmus until you get a color change. (The solid does not dissolve very well, and

vigorous stirring is needed to get even a slight color change.) Do your best with this

one—often there is no discernible color change. If you get no color change, place a

small sample of the oxide of magnesium in an evaporating dish with a small amount

of water and test that as above.

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Before disposing of the mixture, pour about half of it into a small beaker. To one

half add a small volume of concentrated (12 M) hydrochloric acid with stirring. To

the other half add a small volume of 50% (12.5 M) sodium hydroxide with stirring.

Report any evidence of reaction (change!) in either or both, and then dispose in the

proper waste container.

Q-10. If you were to collect the "ash" from the burning of the magnesium

(including the part that went up in smoke), would it weigh less than, more

than, or the same as the original magnesium? Explain.

C. Aluminum

Aluminum oxide is easy to obtain from the hydroxide (by drying to remove water),

but the reverse direction requires extreme conditions we choose not to employ here.

The following exercise will convey the relevant chemistry of aluminum hydroxide.

In a large test tube place about 5 mL of 1.0 M Al(NO3)3. Start adding 6 M NaOH

dropwise. The precipitate may be almost transparent at first. Later it becomes

slushy-like. When it becomes so thick that the contents are a stiff gel, stop the

addition. (Note: Adding too much NaOH will cause the precipitate to dissolve. If

you do not get a precipitate within 20 to 30 drops, discard your mixture in the waste

bottle and try again.) Divide the residue by removing about half to another tube.

(Carry it out on a stirring rod if necessary.) Add a few drops of phenolphthalein

indicator to one tube. Shake the tube regularly as you continue to add NaOH

dropwise. Stop when the solution becomes clear pink (that is, add NaOH until the

precipitate dissolves and the indicator changes from colorless to pink). Set the tube

aside.

To the other tube containing the gel, add 6 M HCl dropwise until the solution

becomes clear and colorless.

Aluminum hydroxide is one of a number of water-insoluble hydroxides that is said

to be amphoteric, that is, it will dissolve not only in a strongly acid solution but also

in a strongly basic one.

D. Carbon

Fill three bottles with water. Then fill these bottles with oxygen as directed by the

instructions posted near the oxygen cylinder. When the bottle is full, turn off the gas

supply (at the needle valve) and slide a glass plate over the mouth. For each of these

bottles quickly remove the glass plate, immediately add about 10 mL of deionized

water, and then recover the bottle with the plate right

Grasp one end of a wooden splint in your tongs, and hold the other end in a gas

flame until it starts to glow. Then lift the glass plate from the first bottle, and thrust


the burning splint into the gas. Use the glass plate to confine gases as well as you

can. When burning slows down, drop the ember into the water, and seal the bottle

with the plate. Shake the bottle for a minute or so to help absorption of any gaseous

products. Test the water with blue litmus paper. (Only a slight color change is

expected.)

E. Phosphorus

Take the two remaining bottles of oxygen—with water in the bottom—to the hood

(one for burning phosphorus, the other for sulfur). Obtain a long-handled iron

spoon, sometimes called a combustion or "deflagrating" spoon.

Use the spatula to pick up a very small amount of red phosphorus—about the size

of a very small pea—and transfer it to the spoon.

CAUTION: Although red phosphorus is not as dangerous

as the other allotrope, white phosphorus, avoid contact with

your skin, and wash your hands after handling it.

Hold the spoon in a gas flame until the phosphorus starts to burn. Then insert it into

the bottle of oxygen, using the glass plate to block some of the opening. When the

combustion dies down, remove the spoon and close the bottle with the glass plate.

Continue to heat the spoon in the flame. You must clean the spoon by burning off

excess phosphorus. If you add too much phosphorus, it will take a long time to burn

off the excess. Shake the bottle to absorb gases in the water.

Test the water solution with both colors of litmus paper. Then discard the solution

in the proper waste container.

F. Sulfur

When the spoon is clean and cool, repeat the procedure of part E above, but

substitute sulfur for phosphorus. The same precautions for size of sample apply.

Note the choking odor of burning sulfur, due to sulfur dioxide. The solution may be

disposed of in the proper waste container.

G. Chlorine

As was mentioned earlier, the various oxides of chlorine are dangerously explosive.

A water solution in which one of these oxides has reacted with the water has been

made available to you. It has been labeled two ways: as an acid "HOClO3" and as a

base "ClO3(OH)."

Page 130

Test the solution with both red and blue litmus paper, and decide which formula is

appropriate. The solution may be discarded in the proper waste container.

IV. Redox behavior of the halogens

A. Oxidation of the bromide ion

Into a small test tube place about 0.1 g of solid KBr. (Weighing is unnecessary. An

amount the size of a pea on the tip of your spatula is about right.)

Hold the tube with your wire test tube holder. (Hold it near the bottom, rather than

at the top as you usually do.)

Drop four drops of concentrated H2SO4 straight on to the solid in the tube. Observe

evidence of reaction. Promote further reaction by rubbing the bottom of the tube on

a warm hot plate. After brown fumes have filled the entire tube, stop heating the

tube.

Dispose of the chemicals in the proper waste container.

B. Oxidation of the iodide ion

Repeat the procedure used in part A, but substitute 0.1 g of KI for the KBr. All

other operations will be the same, except that less heating may be required. Observe

the color of the fumes.

Dispose of the chemicals in the proper waste container.

C. Establishing the "pecking order" among the halogens

All of the halogens are potential oxidizing agents, but not all are equally powerful.

All of the halide ions can be oxidized to the free halogen, but not all with the same

SAFETY FIRST

• All of the free halogens are toxic and irritating. Small amounts of bromine and

iodine will be generated in parts A and B respectively. Do not take more than the

specified amount of starting materials. Work under a hood. Avoid breathing the

vapors.

• Concentrated sulfuric acid is very corrosive to the skin and destructive to clothing.

Wash hands immediately after contact. Remove affected clothing, and rinse with

water.

• Hexane, a solvent something like gasoline, is extremely flammable. Keep it away

from flames and heaters.


ease. From the nine test tube experiments that follow, you should be able to rank the

halogens and relate their ranking—or "pecking order"—to the periodic table.

Label three large test tubes “chlorine water," "bromine water," and "iodine water,"

and obtain about 6 mL of each of these reagents. (Note: Chlorine water is not

Cl2·H2O, that is, it does not have a formula. It is a mixture in the same way as salt

water or seawater.) Label another set of three large test tubes "0.1 M NaCl," "0.1 M

KBr," and "0.1 M KI," and obtain about 5 mL of each of these reagents. In a small

beaker obtain about 10 mL of hexane.

Now, back at your desk, place about 1 mL each of chlorine water, bromine water,

and iodine water into three separate medium test tubes. (Each tube will contain only

one of the above solutions.) Then add 1 mL of hexane to each medium test tube

(These are test tubes 1, 2, and 3 below in the table). Stopper each of the test tubes,

and shake them well. Hexane is volatile; loosen the stopper under a fume hood to

prevent vapor build up inside the tube. Record the color of the upper hexane layer

on the chart in the data section. This is the color of the ‘free’ halogen in a nonpolar

solvent.

The remaining six tests will be mixtures of a halogen (Cl2, Br2, I2) with a halide salt

(NaCl, KBr, KI). Mix about 2 mL of halogen with 2 mL of a salt solution according

to the following mixing chart (test tubes 4 through 9 below). In each case, add 1 mL

of hexane, stopper, and shake well. Loosen the stopper under a fume hood. In each

case record the color of the hexane layer.

Solvent only NaCl(aq) +

solvent

KBr(aq) +

solvent

KI(aq) +

solvent

Cl2 (aq)

Test tube 1:

Hexane and Cl2

only

Test tube 4: Cl2

+ KBr(aq) +

hexane

Test tube 5:

Cl2 + KI(aq) +

hexane

Br2 (aq)

Test tube 2:

Hexane and Br2

only

Test tube 6:

Br2 + NaCl(aq)

+ hexane

Test tube 7:

Br2 + KI(aq) +

hexane

I2 (aq)

Test tube 3:

Hexane and I2

only

Test tube 8:

I2 + NaCl(aq) +

hexane

Test tube 9:

I2 + KBr(aq) +

hexane

Dispose of the contents in the "Halogen and Solvent Waste" container.

ANHYDRIDES

The word anhydride means without water. Anhydrides are binary compounds that do not

contain hydrogen. To obtain the formula of an anhydride, subtract one or more molecules

of water from a formula. Thus, the anhydride of carbonic acid = H2CO3 – H2O = CO2.

If the original compound contains an odd number of hydrogen atoms, it will be necessary

to double its formula first. Also, enough water molecules must be removed so that no

hydrogen remains. Thus the anhydride of phosphoric acid is P2O5.

Page 132


Las Positas College NAME

Chemistry 1A

Dates

Report Sheet: Exp. 15: PERIODIC PROPERTIES

NOTE: All chemical equations must be balanced and must include physical states or

phases.

Name of partner:

I. Reactivity with Water

A. Lithium, Sodium and Potassium

Describe the reaction of each metal with water. Include specific

identification of the products formed.

Write the balanced chemical equation for each reaction:

Q-1 Rank:

Q-2

Q-3

Q-4

Page 134

B. Sodium and Methyl Alcohol

Describe the reaction of sodium with methanol:

Q-5

Q-6

C. Mg and Ca

Observations:

Write the balanced chemical equation for reaction of Ca and water:

Write the balanced chemical equation for reaction of HCl with the white

precipitate:

Q-7

Q-8

Judging by the reactivity of calcium and of magnesium with water, what would you

predict as to the behavior of strontium or barium with water?

________________________________________________________________________

________________________________________________________________________



A. Magnesium

Observations:

Write the balanced equation for the reaction:

Q-9

B. Aluminum

Observations:

Write the balanced equation for the reaction:

Summarize the comparative reactivity of the metals of Group I, of Group II, and of

Group III with water and with dilute acids. What would you predict as to the relative

ease of oxidation of these metals on exposure to air?

________________________________________________________________________

________________________________________________________________________

________________________________________________________________________

________________________________________________________________________

________________________________________________________________________

________________________________________________________________________

Page 136

III. Study of oxides

A. Sodium Peroxide with Water

Observations:

Equation:

Litmus test:

B. Magnesium

Observations:

Equation for burning of Mg:

Litmus test, mixture of ash and water:

Equation for reaction of magnesium oxide and water:

Observations on effect of adding HCl:

Equation for reaction of HCl and magnesium hydroxide:

Q-10.


C. Aluminum

Observations:

Al(NO3)3 and NaOH:

Al(OH)3 and HCl:

Al(OH)3 and NaOH:

Equation for formation of the hydroxide from the nitrate:

Equation for dissolving of the hydroxide in HCl:

Equation for dissolving of the hydroxide in NaOH:

Al(OH)3(s) + NaOH(aq) Na[Al(OH)4](aq)

D. Carbon

Observations:

Equation for the burning of splint (carbon):

Litmus test, water solution of combustion product:

Equation for reaction between carbon dioxide and water:

E. Phosphorus

Observations of combustion:

Page 138

Equation:

P4O10(s)

What is the proper systematic name for P4O10(s)?

Result of litmus test, aqueous solution:

Equation to explain this reaction:

F. Sulfur

Observations of combustion:

Equation:

Result of litmus test, aqueous solution:

Equation to explain this reaction:

G. Chlorine

Result of litmus test, water solution of "HOClO3" or "ClO3(OH)":

Based on the above observation, which formula is correct? _____________

What is the formula of the oxide that is the anhydride of this acid? (See

discussion at the end of the experiment for a description of anhydrides.)

General conclusion concerning the acidity or basicity of the hydroxides (oxides) of the

elements of the third period:

________________________________________________________________________

________________________________________________________________________

________________________________________________________________________


IV. Redox behavior of the halogens

A. Oxidation of the bromide ion

Observations:

Complete the following set of ionic equations:

a) oxidation half-reaction:

b) reduction half-reaction: HSO4–(aq) SO2(aq)

c) overall net ionic equation:

B. Oxidation of the iodide ion

Observations:

Complete the following set of ionic equations:

a) oxidation half-reaction:

b) reduction half-reaction: HSO4–(aq) SO2(aq)

c) overall net ionic equation:

Page 140

C. "Pecking order" of the halogens: Enter the color of the hexane layer in each

box below. In addition, enter "REACTION" or "NO REACTION" in each

case. HINT: Is the color of the hexane layer that of the reactant halogen

(Cl2, Br2 or I2) or is it the color of the product halogen?

solvent only NaCl(aq) +

solvent

KBr(aq) +

solvent

KI(aq) +

solvent

Cl2(aq)

Br2(aq)

I2(aq)

For every combination in which there was a reaction, write the net ionic equation:

____________________________________________________

____________________________________________________

____________________________________________________

On the basis of these results, circle the best

• oxidizing agent: Cl2 Br2 I2 Cl– Br– I–

• reducing agent: Cl2 Br2 I2 Cl– Br– I–

Rank each of the 6 species in this activity series:

oxidizing agents:

best worst

reducing agents:

worst best


POST-LAB Name_________________


Experiment 15 Group Relationships and Periodic Properties

1. Most but not all acids contain oxygen.

Name one that does not:

2. For an acid that does contain oxygen, we can derive the formula of its anhydride.

Write the formula of the anhydride of

H2SO3 HNO2

H3PO3 HOCl

3. Are the above oxides acidic, basic, or amphoteric?

4. Write the formulas of the anhydrides of the following bases:

KOH ________ Ba(OH)2 ________ Al(OH)3 ________ AgOH _______

5. Are the above oxides acidic, basic, or amphoteric?

6. In what region of the periodic table would you expect to find elements that form

amphoteric oxides and hydroxides?

____________________________________

7. Utilizing the general trends you have observed in this experiment, predict whether

the following reactions would occur. If a reaction is predicted to occur, write a

balanced conventional equation for the reaction. If no reaction is predicted to

occur, write the words "no reaction".

a. francium metal is

added to water.

b. strontium metal is

added to water.

c. rubidium metal and

Page 142

chlorine gas are mixed

d. fluorine gas is added

to aqueous calcium

chloride.

e. aqueous cesium bromide

and aqueous iodine are

mixed.


PRE-LAB Name_________________


Experiment 15 Group Relationships and Periodic Properties

1. Define these terms as used in chemistry:

Periodic ___________________________________________________________

Slurry ___________________________________________________________

Allotrope ___________________________________________________________

2. Give the symbol of the element that is:

a) in the second period and in Group IIIA (Group 13): ______________

b) in the third period and in Group IIA (Group 2): ______________

3.

Give the

formula of the

Na Mg Al Si

oxide of this

element

of the

hydroxide

4.

What two acid-base indicators are used in

this experiment?

What is the color of each in...

Acidic solution Basic solution

1)

2)

5. Three members of the halogen family are studied in this experiment; give their

molecular formulas

physical state at room

temperature

color

Page 144

Three halide ions are also studied in the form of salts; give their formulas:

__________ __________ __________

6. The oxide left by the removal of water from an acid or base is said to be the

anhydride of that acid or base. The process is reversible. For instance:

H2SO4 giving up H2O results in SO3 and SO3 + H2O

=

H2SO4

2 KOH giving up H2O results in K2O and K2O + H2O = 2 KOH

a. What is the anhydride of nitric acid, HNO3? ___________________

Remember, there can be no hydrogen in an oxide. (Hint: Start with 2 molecules of

acid.)

b. What is the anhydride of barium hydroxide, Ba(OH)2?_________________


Experiment 16

Model Making & Geometry

PURPOSE

The purpose of this experiment is to gain experience in

a. predicting the shape, drawing stereochemical formulas, and building models of

molecules and ions;

b. predicting hybridization, numbers of sigma and pi bonds, and drawing three-

dimensional orbital overlap sketches (the latter for selected molecules only).

Method

Lewis structures, in conjunction with VSEPR theory and/or valence bond theory, are used

to predict molecular geometries. Molecular models are used as an aid in visualizing and

drawing the shapes/stereochemical formulas for molecules and ions. Hybridization is

determined from the Lewis structures.

PROCEDURE

For each molecule or ion listed below:

1. Calculate AE, defined as available electrons or the total number of valence

electrons in the entire molecule or ion, adjusted for any charge in the case of ions.

2. Draw the complete Lewis structure, including all resonance forms. Satisfy the octet

rule if possible.

3. Count the number of bonded atoms (BA) and lone pairs (LP) for the central atom.

Add them to get the coordination number of the central atom. The coordination

number is used to determine the electron pair geometry (2 = linear, 3 = trigonal

planar, 4 = tetrahedral, 5 = trigonal bipyramidal, 6 = octahedral).

4. Build a model by selecting an atom piece that has the correct number of holes to

correspond with the coordination number. Use sticks with no atoms at the ends to

represent lone pairs of electrons. Note that it is only necessary to build one model

for any species that exhibits resonance. Use an atom piece with an extra hole if

there is a double bond.

5. Draw a stereochemical formula for the model. Include approximate bond angles on

the drawing. Use element symbols in the drawing and do not show lone pairs to

outside atoms.

6. State the electron pair geometry and the molecular shape.

7. Have your model checked by your instructor.

8. Continue in the second part to identify the hybridization of the central atom, the

number of sigma bonds and pi bonds in each species, and the formal charges on

atoms.

9. In the third part, you will use formal charge to select between possible structures

and you will draw three-dimensional orbital overlap sketches for a couple of

molecules.

Page 146

AE

Lewis structure

BA

LP

Coord #

Stereochemical Drawing

Instructor approval

Electron Pair Geometry/

Molecular Geometry

NH3

NH4+

CO2

SO2


AE

Lewis structure

BA

LP

Coord #


Instructor approval


Molecular Geometry

SO3

POF3

PCl5

Page 148

AE

Lewis structure

BA

LP

Coord #


Instructor approval


Molecular Geometry

SF6

ClF3

IF5


AE

Lewis structure

BA

LP

Coord #


Instructor approval


Molecular Geometry

BiCl52–

XeF2

XeF4

Page 150

AE

Lewis structure

BA

LP

Coord #


Instructor approval


Molecular Geometry

TeF4

COCl2

HCN


Part II: Complete the following table. Redraw the Lewis structure for each molecule or

ion, calculate formal charges, and label any non-zero formal charge.

species Lewis structure; label any non-zero

formal charges on each structure

hybridization of the central

atom # #

NH3

NH4+

CO2

SO2

SO3

POF3

PCl5

Page 152




atom # #

SF6

ClF3

IF5

BiCl52–

XeF2

XeF4

TeF4





atom # #

COCl2

HCN

Page 154

Part III

1. Use formal charge to choose the best arrangement for the atoms in a molecule

whose formula is usually written N2O: N—N—O or N—O—N. Include resonance

structures in your evaluation. All molecules should satisfy the octet rule for all

atoms.

2. The following molecules do not have just one central atom. Draw Lewis diagrams,

build the models, and then draw stereochemical formulas.

C2H4Cl2 1,1-isomer: both Cl’s on the same C

C2H4Cl2 1,2-isomer: one Cl on each C

C2H2Cl2 1,1-isomer

C2H2Cl2 1,2 cis isomer: use index of textbook to find description of cis and trans

isomers

3. Draw one or two pictures showing the overlap of orbitals to form sigma and pi

bonds in COCl2.


POST-LAB Name_________________


Experiment 16 Model Making & Geometry

1. In the molecule drawn below, how many sigma bonds are there?

how many pi bonds? what is the hybridization of atom A?

what is the hybridization of atom B?

H O C C C C C C C C C

C

N

Cl

H H H H H H H

H

HH

HAB

2. Draw the Lewis structure for KrCl4:

How many electrons?

What is the shape of the molecule?

What hybrid orbitals are used by Cl?

Is the molecule polar?

3. Assign formal charges to all the atoms in the structure for sulfuric acid. What is the

oxidation number of the sulfur?

H O S O H

O

O

Page 156


Experiment 17

Metallic and Ionic Crystal Lattices


PURPOSE

The purpose of this experiment is to construct models of the common metallic and ionic

crystal lattices and to study their characteristics.

Method

Historically, Styrofoam balls were connected with wooden toothpicks. Now we will use

the Solid State Model Kits to reproduce the following crystal lattices:

IONIC METALLIC

A. sodium chloride (rock salt) D. face centered cubic

(Cubic close packed)

B. cesium chloride E. body centered hexagonal

(Hexagonal close packed)

C. wurtzite (Hexagonal ZnS) F. body-centered cubic

By looking at these models, you will be able to identify many features of these lattices.

Figure 1 shows the unit cells for the metallic crystal lattices and Figure 2 shows the unit

cells for the ionic crystal lattices.

PROCEDURE

You will use the Solid State Model Kit and the Instructions in the kit to answer the

following questions. Refer to the diagrams both here and in the Model Kit for assistance.

1. Find the instruction manual in your Model Kit. Read “Getting Started” along with the

rest of the information on pages 1-5.

2. Do the NaCl example on pages 6-7 using the Model Kit. This example takes you

through the questions step by step, while all other examples are abbreviated. You won’t

have any problems IF you have followed instructions 1 and 2!

3. Answer all of the questions for the NaCl example in part A below, then proceed with

the other parts. Answer all of the questions while you have the model in front of you for

reference.

Page 158

Introduction

The following are the simplest crystal lattice structures found in nature.

a. face-centered cubic b. body-centered hexagonal c. body-centered cubic (cubic close-packed) (hexagonal close-packed) Figure 1: Common metallic crystal lattice unit cells ________________________________________________________________

Black = - ion White = + ion

a. CsCl (simple cubic - ions) b. NaCl (face-centered cubic – ions) cesium chloride sodium chloride (rock salt)

c. ZnS (face-centered cubic – ions) d. ZnS (body-centered hexagonal – ions) cubic zinc sulfide (zinc blende) hexagonal zinc sulfide (wurtzite)

Figure 2. Common ionic crystal lattice unit cells for 1:1 salts


CRYSTAL LATTICES

Construct each model using the Instructions in your Solid State Model Kit, answering

questions as you go. As you progress, have your work checked by the instructor. Be

prepared to answer questions about the information you have gathered.

IONIC LATTICES

A. Sodium Chloride (Rock Salt), Octahedral Sites: Model A

Follow instructions on pages 6-7 of your Model Kit Instruction Manual.

Figure 3. The Na+ ion in the octahedral hole of the NaCl lattice.

1. What is the CN of X? (That is, how many Cl

- ions are around each Na

+ ion?) _____

This is known as the coordination number (CN) or number of nearest neighbors.

Figure 4. Sodium Chloride (Rock Salt). Smaller Na+ spheres in the octahedral holes of a face-centered cubic lattice of Cl- spheres.

Instructor's Initials

2. Within each unit cell, how many net Cl

- ions are:

a. on the corners (net number) _____

b, on the edges (do not count corners again) _____

c. within the faces _____

d. within the center _____

e. net total _____

3. Within each unit cell, how many net Na

+ ions are:

a. on the corners (net number) _____

b. on the edges (do not count corners again) _____


Page 160


e. net total _____

4. The chloride ions alone describe what lattice? _______________________

5. The sodium ions alone describe what lattice? _______________________

6. Calculate the radius of the sphere that will touch all

six surrounding spheres. Hint: tilt the model until

you see the plane of 4 spheres shown to the right.

7. Calculate the radius ratio R/r (or r/R). ______

B. Cesium Chloride. Cubic Sites: Model B

Follow the instructions beginning on page 11 of the instruction manual.

Figure 5. The Cs+ ion in the cubic hole of the CsCl lattice 1. What is the CN of X? (That is, how many Cl

- ions are around each Cs

+ ion?) _____

Now complete the CsCl lattice.

Figure 6. Model B. Cesium Chloride, multiple unit cells

Instructor's Initials _______

2. How many complete unit cells of CsCl are contained in this model? _____

3. Within each unit cell, how many net Cl

- ions are:

a. on the corners _____

Small sphere of

radius r


b. on the edges _____



e. net total _____

4. Within each unit cell, how many net Cs

+ ions are:





e. net total _____

5. The chloride ions alone describe what lattice? _______________________

6. The cesium ions alone describe what lattice? _______________________

C. Hexagonal Zinc Sulfide (Wurtzite), Tetrahedral Sites: Model C

Follow the instructions on page 54 of the Solid State Kit Instruction Manual.

Figure 7. The Zn2+ ion in the tetrahedral hole of the hexagonal ZnS lattice.

1. What is the CN of X? (That is, how many S

2- ions are around each Zn

2+ ion?) ______

Instructor's Initials ______

2. How many complete unit cells of ZnS are contained in this model? ______

3. Within each unit cell, how net many S

2- ions are:



Page 162

b. with in the faces _____

c. within the center _____

d. net total _____

4. Within each unit cell, how many net Zn

2+ ions are:





e. net total _____

METALLIC LATTICES

D. Face-Centered Cubic (FCC) or Cubic Close-Packed: Model D

Follow the instructions on page 27 of the Solid State Model Kit Instruction

Manual.

Figure 8. Model A. Face-Centered Cubic.

Instructor's Initials___________

1. How many spheres are directly touching the sphere marked with an "X" above?_____

2. Within the unit cell, how many (net number) spheres are:

a. on the corners (net contribution) net no.= __________

b. within the faces (do not count the corners again) net no. = __________

c. in the center net no. = __________

d. total all number of spheres in the unit cell ____________

3. FCC is one of the two close-packed lattices. In one of these, every other layer is the

same, called ABAB packing. This means that for every sphere in the 1st layer,-there is


a sphere right above it in the 3rd layer. In the other type of close packing, spheres

repeat every third layer, called ABCABC packing. In this arrangement, each sphere in

the 1st layer does not have a sphere right above it until the 4th layer.

Is the packing in FCC, ABAB or ABCABC ?

(Hint: the four layers you have made are NOT the close packed layers. In order to see the

packing type, use the cube made by the first three

4. What are the relationships among the sides (a, b, c) in the

FCC unit cell (equal or unequal) _____

5. What are the angles () in the FCC unit cell? _____

Use the diagram of the face-centered cubic lattice at the right to

answer the following questions.

6. How many net spheres are in the unit cell? ______

7. What is the volume of a sphere in terms of R, the sphere radius? ______

8. What is the total volume occupied by the spheres in the unit cell in terms of R?

9. What is the face diagonal (corner-to-corner distance on face) in terms of R?

10. What is the value of S in terms of R? Use the Pythagorean Theorem on the right

triangle given. Show work below.

Page 164

11. Give the volume of the cube in terms of S.

12. Give the volume of the cube in terms of R.

13. Calculate the percent of occupied space.

% occupied = (Vspheres/Vcube) x 100%

E. Bodv-Centered Hexagonal or Hexagonal Close-Packed (HCP): Model E

Follow the directions on page 24 of the Solid State Model Kit Instruction Manual.

Figure 9. Model E. Body-Centered Hexagonal.

Instructor's Initials ______

1. What is the coordination number (number of nearest

neighbors) in the hexagonal close packed lattice? ______


a. on the corners (net contribution) net no. = _____

b. within the faces net no. = _____

c. in the center (inside the cell) net no. = _____

d. total net number of spheres in unit cell _____


3. Is the packing in HCP ABAB_________ ABABC_________?

(Hint: the four layers you have made ARE the close packed layers in the

hexagonal close-packed lattice.)

4. What are the relationships among the sides (a, b, c)

in the HCP unit cell (answer: equal or unequal)? _____

5. What are the angles () in the HCP unit cell? _____

F. Body-Centered Cubic (BCC): Model F

Figure 10. Model F. Body-Centered Cubic.

Instructor's Initials _____

1. What is the coordination number (number of nearest

neighbors) in the body-centered cubic lattice? _____


a. on the corners (net contribution) net no = _____

b. within the faces (do not count the corners again) net no = _____

c. within the center net no = _____

d. total net number of spheres in unit cell _____

3. The body-centered spheres in the neighboring unit cells are only 14% farther

away than those touching the body centered sphere in this cell. Such spheres are

called next nearest neighbors. How many next nearest neighbors are in the BCC

lattice?

_____

4. What are the relationships among the sides (a, b, c)

in the BCC unit cell (equal or unequal)? _____

5. What are the angles () in the BCC unit cell? _____

Page 166

Use the two diagrams of the body-centered cubic lattice below to-answer the following

questions.

6. How many net spheres are in the unit cell? _____

7. What is the volume of all the spheres in the unit cell, in terms of R? _____

8. What is the body diagonal (BD) of the cube in terms of R? Note: the central

sphere touches each corner sphere.

_____

9. What is the face diagonal (FD) of the cube in terms of S? Use Pythagorean

theorem on the right triangle which has FD as the hypotenuse. Show your

work below. _____

10. What is the value of S in terms of R? Use the Pythagorean theorem

on the right triangle which has BD as the hypotenuse. Use your

values of BD in terms of R and FD in terms of S. Show your work below.

_____

11. Give the volume of the cube in terms of R. _____

12. Calculate the percent of occupied space. _____

13. Which type of packing is more efficient? ______FCC or ______ BCC


POST-LAB Name_________________


Experiment 17 Metallic and Ionic Crystal Lattices

1. Cesium chloride structure

a. Why is it not proper to say that CsCl has a body-centered cubic structure?

b. Thallium(I) chloride crystallizes in the cesium chloride lattice, as shown in

Fig. 5. The shortest distance between the center of a Tl+ ion and the center

of a Cl- ion is 333 pm.

1) What is the length of the edge of a unit cell of TlCl? Hint: How

many TlCl diameters equal one body diagonal?

2) What is the density in g/cm3 of TlCl?

2. Superconductors. Given the idealized unit cell for the structure of a

superconductor shown below to be orthorhombic, answer the following questions.

Page 168

a. How many copper atoms are there per unit cell? (Show calculations.)

b. What is the number of oxygens per unit cell? and Y? and Ba?

c. What is the formula for this superconductor?

d. The coppers in this unit cell occupy two different kinds of sites. Based on the

number of each of these types and the formula for the compound, what are the

oxidation states for copper in this compound?

The following Questions will likely require some research and reading outside of our

textbook! Go dig out the information that you need!

3. a. What makes a substance a superconductor?

b. Give the properties of superconductors.

4. Explain why ionic solids are brittle and metals are malleable.

5. Explain why metals decrease in conductivity with a rise in temperature while

semiconductors increase in conductivity with temperature.

6. How does the band theory explain luster of metals and heat conductivity? Include

labeled diagrams in your answer.


PRE-LAB Name_________________


Experiment 17 Metallic and Ionic Crystal Lattices

Show your work!

1. Write the equation for the volume, V, of a cube in terms of its side, s:

V =

2. Write the equation for the volume, V, of a sphere in terms of its radius, R:

V =

3. Write the Pythagorean Theorem equation for a right triangle with sides a and b

adjacent to the right angle and hypotenuse c:

4. If a right triangle has two equal sides of length a, express the length of the

hypotenuse c in terms of a:

c =

5. For a cube with sides of length s:

a. Express the length of a diagonal across a face, FD, in terms of s:

FD =

b. Express the length of a diagonal through the body, BD, in terms of s:

BD =

6. Simplify √8:

√8 =

Page 170


Experiment 18

Evaporation and Intermolecular Attractions


In this experiment, Temperature Probes are placed in various liquids. Evaporation occurs when the probe is removed from the liquid’s container. This evaporation is an endothermic process that results in a temperature decrease. The magnitude of a temperature decrease is, like viscosity and boiling temperature, related to the strength of intermolecular forces of attraction. In this experiment, you will study temperature changes caused by the evaporation of several liquids and relate the temperature changes to the strength of intermolecular forces of attraction. You will use the results to predict, and then measure, the temperature change for several other liquids.

You will encounter two types of organic compounds in this experiment—alkanes and alcohols. The two alkanes are n-pentane, C5H12, and n-hexane, C6H14. In addition to carbon and hydrogen atoms, alcohols also contain the -OH functional group. Methanol, CH3OH, and ethanol, C2H5OH, are two of the alcohols that we will use in this experiment. You will examine the molecular structure of alkanes and alcohols for the presence and relative strength of two intermolecular forces—hydrogen bonding and dispersion forces.

Figure 1 Materials

Windows PC methanol (methyl alcohol) Vernier computer interface ethanol (ethyl alcohol) Logger Pro acetone 4 Temperature Probes n-pentane 4 pieces of absorbent paper (2 cm X 4 cm) 1-propanol 4 small rubber bands 1-butanol Colored time tape n-hexane

Page 172

PROCEDURE

1. Obtain and wear goggles! CAUTION: The compounds used in this experiment are flammable and poisonous. Avoid inhaling their vapors. Avoid contacting them with your skin or clothing. Be sure there are no open flames in the lab during this experiment. Notify your teacher immediately if an accident occurs.

2. Prepare the computer for data collection.

Obtain 4 stainless steel temperature probes and plug them into Channels 1-4. Prepare the computer for data collection by opening Logger Pro. You should see 4 temperature readings corresponding to each probe. On the Graph window, the vertical (temperature) axis should be scaled from 0 to 30C and the horizontal (time) axis from 0 to 400 seconds. In the Experiment menu, choose Data Collection, and then set the experiment length to 400 seconds. This is different from changing the display. You must also double click on one of the values from the axis you want to change. Choose axis options and set the range to the correct values.

3. Wrap Probes 1-4 with identical pieces of absorbent paper secured by small rubber

bands as shown in Figure 1. Roll the paper around the probe tip in the shape of a cylinder. Hint: First slip the rubber band up on the probe, wrap the paper around the probe, and then finally slip the rubber band over the wrapped paper. The paper should be even with the probe end.

4. Place 4 large test tubes in you test tube rack and label them Methanol, Ethanol, Acetone, and Pentane. Place enough of each pure liquid in the corresponding test tube to fill it up 2-3 cm. Put Probe 1 in the methanol test tube, Probe 2 in the ethanol test tube, Probe 3 in the acetone test tube, and Probe 4 in the pentane test tube. Make sure the containers do not tip over.

5. Prepare 4 pieces of colored time tape, each about 10-cm long, to be used to tape the probes in position during Step 6.

6. After the probes have been in the liquids for at least 45 seconds, begin data collection by clicking Collect . Monitor the temperature for at least 15 seconds to establish the initial temperature of each liquid. Then simultaneously remove the probes from the liquids and tape them so the probe tips extend 5 cm over the edge of the table top as shown in Figure 1.

7. When each of the temperatures have reached minimums and have begun to increase, click Stop to end data collection. Click the Statistics button, , then click OK to display a box for both probes. Record the maximum (t1) and minimum (t2) values for Temperature 1 (methanol), Temperature 2 (1-propanol), Temperature 3 (acetone), and Temperature 4 (pentane).

8. For each liquid, subtract the minimum temperature from the maximum temperature to determine t, the temperature change during evaporation.

9. Roll the rubber band up the probe shaft and dispose of the absorbent paper in the “Chemically Contaminated Paper” container.

10. Based on the t values you obtained for these 4 substances, plus information in the Pre-Lab exercise, predict the size of the t value for 1-propanol and 1-butanol. Compare their hydrogen-bonding capabilities and molecular weights to those of methanol and ethanol. Record your predicted t, then explain how you arrived at this


answer in the space provided. Do the same for n-hexane. It is not important that you predict the exact t value; simply estimate a logical value that is higher, lower, or between the previous t values.

11. Repeat the experiment using methanol, ethanol, 1-butanol and 1-propanol. The methanol and ethanol samples will act as a control so that you can compare this experiment with your first experiment.

12. Now repeat the experiment using pentane and hexane.

DATA




for at least one trial if multiple runs were made. Trial #1

Substance methanol ethanol acetone n-pentane

t1 (°C)

t2(°C)

t (t1–t2) (°C)

Trial #2

Substance Predicted

t (t1–t2) (°C)

Explanation

t1

(°C)

t2

(°C)

Experimental

t (t1–t2) (°C)

methanol

ethanol

1-propanol

1-butanol

n-pentane

n-hexane

Page 174

DATA ANALYSIS

Answer the following questions in your notebook.

1. Two of the liquids, 1-propanol and acetone, have similar molecular weights, but

significantly different t values. Explain the difference in t values of these substances, based on their intermolecular forces.

2. Which of the alcohols studied has the strongest intermolecular forces of attraction? The weakest intermolecular forces? Explain using the results of this experiment.


PRE-LAB Name_________________


Experiment 18 Evaporation and Intermolecular Attractions

Prior to doing the experiment, complete the Pre-Lab table. The name and formula are

given for each compound. Draw a structural formula for a molecule of each compound.

Then determine the molecular weight of each of the molecules. Dispersion forces exist

between any two molecules, and generally increase as the molecular weight of the

molecule increases. Next, examine each molecule for the presence of hydrogen bonding.

Before hydrogen bonding can occur, a hydrogen atom must be bonded directly to an N,

O, or F atom within the molecule. Tell whether or not each molecule has hydrogen-

bonding capability.

Substance Formula Structural Formulas Molecular

Weight Hydrogen Bond

(Yes or No)

methanol CH3OH

ethanol C2H5OH

1-propanol C3H7OH

1-butanol C4H9OH

acetone C3H6O

n-pentane C5H12

n-hexane C6H14

Page 176


Experiment 19

Using Freezing-Point Depression to Find Molecular Weight

When a solute is dissolved in a solvent, the freezing temperature is lowered in proportion to the number of moles of solute added. This property, known as freezing-point depression, is a colligative property; that is, it depends on the ratio of solute and solvent particles, not on the nature of the substance itself. The equation that shows this relationship is:

T = Kf • m

where T is the freezing point depression, Kf is the freezing point depression constant for a particular solvent (3.9°C-kg/mol for lauric acid in this experiment

1 ), and m is the

molality of the solution (in mol solute/kg solvent).

In this experiment, you will first find the freezing temperature of the pure solvent, lauric acid, CH3(CH2)10COOH. You will then add a known mass of benzoic acid solute, C6H5COOH, to a known mass of lauric acid, and determine the lowering of the freezing temperature of the solution. In an earlier experiment, you observed the effect on the cooling behavior at the freezing point of adding a solute to a pure substance. By measuring the freezing point depression, T, and the mass of benzoic acid, you can use the formula above to find the molecular weight of the benzoic acid solute, in g/mol.

Figure 1: Note-- use a second probe to measure the water bath temperature.

PURPOSE The purpose of this experiment is to: a. study the effect of a solute on the freezing point of a solvent, b. to learn a method of accurate freezing point determination, and c. to use that method to determine the molar mass of an unknown

1 “The Computer-Based Laboratory,” Journal of Chemical Education: Software, 1988, Vol. 1A, No. 2, p.

73.

Page 178

Materials

Power Macintosh or Windows PC Buret clamp, wire gauze Vernier computer interface 18 X 150-mm (medium sized) test tube Logger Pro lauric acid 2 Temperature Probes benzoic acid 400-mL beaker ring stand, ring clamp, 2 utility clamps

Safety: Benzoic acid is moderately toxic by ingestion; irritates eyes, skin and respiratory tract. Use the fume hoods and avoid breathing vapors. There will be open flames in the lab. Keep the cables away from the heat of the Bunsen burner as in figure 1. This is important becaure the cables are quickly damaged by excessive heat. Be sure to dispose of unknowns in marked containers.

PROCEDURE

1. Obtain and wear goggles.

2. Prepare the computer for data collection by opening the Experiment 15 folder from Chemistry with Computers. Then open the experiment file that matches the probe you are using. The vertical axis of the graph has temperature scaled from 0°C to 100°C. The horizontal axis has time scaled from 0 to 10 minutes.

Part I Determine the Freezing Temperature of Pure Lauric Acid

3. Clean and dry a medium test tube. Weigh accurately (to +/- 0.001g) about 7 or 8 grams of lauric acid in a preweighed weighing boat. Transfer all of the lauric acid to the test tube. You might chose to reweigh the boat with any solid dust that remains to correct for that which does not get into the test tube. Label your test tube and place it into one of the hot water baths on the back counter to melt the lauric acid.

4. Add about 270 mL of tap water with a temperature between 20 and 25°C to a 300-mL Berzelius beaker (special from the stockroom). Assemble the apparatus in Figure 1, making sure to add a second temperature probe in the water bath.

5. Insert the Temperature Probe into the hot lauric acid. About 30 seconds are required for the probe to warm up to the temperature of its surroundings and give correct temperature readings. During this time, fasten the utility clamp to the ring stand so the test tube is above the water bath. Then click Collect to begin data collection. CAUTION: Be careful not to spill the hot lauric acid on yourself and do not touch the bottom of the test tube.

6. Lower the test tube into the water bath. Make sure the water level outside the test tube is higher than the lauric acid level inside the test tube. If the lauric acid is not above 50°C, reheat the lauric acid sample and begin again.

7. With a very slight up and down motion of the Temperature Probe, continuously stir the lauric acid during the cooling. Hold the top of the probe and not its wire.


8. Continue with the experiment until data collection has stopped after 10 minutes. Do not attempt to pull the probe out—this might damage it. You can reuse the lauric acid sample for Part II if you have not lost any of the lauric acid.

9. To determine the freezing temperature of pure lauric acid, you need to determine the mean (or average) temperature in the portion of graph with nearly constant temperature. Move the mouse pointer to the beginning of the graph’s flat part. Press the mouse button and hold it down as you drag across the flat part of the curve, selecting only the points in the plateau. Click on the Statistics button, . The mean temperature value for the selected data is listed in the statistics box on the graph. Record this value as the freezing temperature of lauric acid. Click on the upper-right corner of the statistics box to remove it from the graph.

Part II Freezing Temperature of a Solution of Benzoic Acid and Lauric Acid

10. Prepare the computer for data collection. From the Data menu, choose Store Latest Run. This stores the data so it can be used later. To hide the curve of your first data run, click the Temperature vertical-axis label of the graph, and uncheck the Run 1 box. Click OK .

11. Weigh approximately 1g of benzoic acid and add it to the test tube containing lauric acid. Put this test tube back into the hot water bath to melt the solids and mix thoroughly with the temperature probe. Repeat Steps 3-8 to determine the freezing point of this mixture.

12. When you have completed Step 8, click on the Examine button, . To determine the freezing point of the benzoic acid-lauric acid solution, you need to determine the temperature at which the mixture initially started to freeze. Unlike pure lauric acid, cooling a mixture of benzoic acid and lauric acid results in a gradual linear decrease in temperature during the time period when freezing takes place. As you move the mouse cursor across the graph, the temperature (y) and time (x) data points are displayed in the examine box on the graph. Locate the initial freezing temperature of the solution, as shown here. Record the freezing point in the Data and Calculations table.

13. To print a graph of temperature vs. time showing both data runs:

a. Click the Temperature vertical-axis label of the graph. To display both temperature runs, check the Run 1 and Latest boxes. Click OK .

b. Label both curves by choosing Make Annotation from the Analyze menu, and typing “Lauric acid” (or “Benzoic acid-lauric acid mixture”) in the edit box. Then drag each box to a position on or near its respective curve.

c. Print a copy of the Graph window. Enter your name(s) and the number of copies of the graph you want.

Time

Freezing Point

Page 180

DATA





Mass of lauric acid g

Mass of benzoic acid g

Freezing temperature of pure lauric acid °C

Freezing point of the benzoic acid–lauric acid mixture

°C

Freezing temperature depression, t

°C Molality, m

mol/kg

Moles of benzoic acid

mol

Molecular weight of benzoic acid (experimental)

g/mol

Molecular weight of benzoic acid (accepted)

g/mol

Percent error

%

CALCULATIONS

1. Determine the difference in freezing temperatures, t, between the pure lauric acid

(t1) and the mixture of lauric acid and benzoic acid (t2). Use the formula, t = t1 - t2.

2. Calculate molality (m), in mol/kg, using the formula, t = Kf • m (Kf = 3.9°C-kg/mol for lauric acid).

3. Calculate moles of benzoic acid solute, using the answer in Step 2 (in mol/kg) and the mass (in kg) of lauric acid solvent.

4. Calculate the experimental molecular weight of benzoic acid, in g/mol. Use the original mass of benzoic acid from the Data and Calculations table, and the moles of benzoic acid you found in the previous step.

5. Determine the accepted molecular weight for benzoic acid from its formula, C6H5COOH and calculate the percent error.


Experiment 20

Acid Rain


PURPOSE

In this experiment, you will observe the formation of four acids that occur in acid rain:

carbonic acid, H2CO3

nitrous acid, HNO2

nitric acid, HNO3

sulfurous acid, H2SO3

Carbonic acid occurs when carbon dioxide gas dissolves in rain droplets of unpolluted air:

(1) CO2(g) + H2O(l) H2CO3(aq)

Nitrous acid and nitric acid result from a common air pollutant, nitrogen dioxide (NO2). Most nitrogen dioxide in our atmosphere is produced from automobile exhaust. Nitrogen dioxide gas dissolves in rain drops and forms nitrous and nitric acid:

(2) 2 NO2(g) + H2O(l) HNO2(aq) + HNO3(aq)

Sulfurous acid is produced from another air pollutant, sulfur dioxide (SO2). Most sulfur dioxide gas in the atmosphere results from burning coal containing sulfur impurities. Sulfur dioxide dissolves in rain drops and forms sulfurous acid:

(3) SO2(g) + H2O(l) H2SO3(aq)

In the procedure outlined below, you will first produce these three gases. You will then bubble the gases through water, producing the acids found in acid rain. The acidity of the water will be monitored with a pH Sensor.

Materials

Windows PC Solid NaNO2 Vernier computer interface Solid NaHCO3 Logger Pro Solid NaHSO3 Vernier pH Sensor 1.0 M HCl with 2 mL plastic dropper Wash bottle with distilled water 3 2 mL plastic droppers Rubber stopper, size 00 3 special glass sample vials Beakers, 600, 400, 250, 150, mL Tap water 3 20 X 150 mm test tubes

NO

2

SO

CO

2

2

2H

H

H

H

CO

NO

NO

SO

2

3

3

3

2

Page 182

PROCEDURE

1. Think safety continuously!

2. Label the three clean medium (20 x 150 mm) test tubes with the formula of the solid they will contain: “NaHCO3”, “NaNO2”, and “NaHSO3”. Label the three special (for this experiment) clean glass sample vials with the formula of the gas that will be dissolved in the 6 mL water that they will contain: “CO2”, “NO2” and “SO2”. You can use a 150-mL beaker to support the test tubes.

3. Using a clean metal scoop, place some solid NaHCO3.in the tube labeled “NaHCO3”. Add enough NaHCO3 to fill the curved bottom end of the test tube.

4. Repeat the Step 3 procedure to add solid NaNO2 and NaHSO3 to the other two appropriately labeled test tubes. CAUTION: Avoid inhaling dust from these solids.

5. Fill the clean 250 mL beaker about half full with clean tap water.

6. Carefully unscrew the little plastic bottle on the pH sensor and carefully place it upright on the lab bench so that the storage solution does not spill. The glass bulb at the end of the pH sensor is very fragile, so be careful not to hit it with anything. Carefully slip the plastic bottle cap off from the cylindrical wall of the pH sensor. Wash the pH sensor by squirting deionized water from a wash bottle onto the shaft and around the bulb. A 600 mL beaker serves as a good waste container to collect the wash water as it runs off the pH sensor. After washing, place the pH sensor in the 250 mL half full beaker of clean tap water. Screw the lid on the plastic bottle and place the 00 rubber stopper in the hole to prevent spilling the storage solution. Set this plastic bottle, with lid on and rubber stopper in the hole, in a safe place to prevent spilling until the end of the experiment, when you have finished with the pH sensor and will store it for future use.

7. Using a 10 mL graduated cylinder, measure 6.0 mL of clean tap water into the clean sample vial labeled “CO2”.

8. With the pH sensor connected to CH1 on the side of the LoggerPro interface, calibrate the pH meter using pH 7 and pH 4 buffers. Go to the “Experiment” tab on the menu and click on “Calibrate”. Click “Calibrate Now”. Immerse the pH meter into the pH 7 buffer, enter 7.00, and hit “Keep”. Do the same for the pH 4 calibration.

9. Prepare the computer for data collection by opening by clicking on “File” in the upper left corner of the menu bar, select “open”, double click on “Chemistry with Computers”, double click on “Experiment 22 Acid Rain”, open (or double click) “Experiment 22 pH Sensor.MBL”. A screen appears with the vertical axis of the graph scaled with pH from 0 to 10 pH units. The horizontal axis has time scaled from 0 to 100 seconds. Change the 100 seconds time duration to 1000 seconds, by going to the “Experiment” tab on the menu and clicking on “Data Collection”. Change the time interval for collection to 1000 seconds. Check to see that the input display at the lower right in the meter window shows a pH value between 6 and 9 for the water.

Safety First! Wear goggles. Avoid ingesting dust from solids and breathing gasses produced. The glass bulb at the tip of the pH sensor is fragile, so please do not hit it.


10. Remove the pH sensor from the 250 mL beaker of tap water, and rinse it with deionized water, allowing the rinse water to collect in the 600 mL beaker waste container. Transfer the pH sensor to the sample vial (labeled “CO2”) with the 6 mL of tap water. The pH sensor in the sample vial is very unstable vertically and tips over, so place the sample vial, with pH sensor inserted, into a 400 mL beaker to provide vertical stability. Very gently stir the pH sensor in the 6 mL tap water in the sample vial. Note that in a minute or so the pH in the meter display window is becoming somewhat constant.

11. Click on “Collect” and start collecting pH data on the tap water sample

12. Using the 2 mL plastic dropper provided with the reagent bottle 1.0 M HCl, add about 1 mL (1/2 full dropper, about 20 drops) of 1.0 M HCl to the test tube with the solid NaHCO3. Try to wash powder stuck on the walls of the test tube to the bottom as you add the 1.0 M HCl. Fizzing is observed as carbon dioxide, CO2, is generated in this tube. Squeeze the bulb of a clean 2 mL plastic dropper and insert the tip and stem of the dropper into the test tube and release the squeezing as the bulb end gets close to the open end of the test tube. Note that the bulb of the 2 mL plastic dropper catches at the mouth of the test tube, preventing the dropper from touching the solution at the bottom of the test tube and providing somewhat of a seal at the open end of the test tube. The dropper must not touch the solution at the bottom of the test tube. Very gently swirl the test tube that now contains NaHCO3 and HCl and dropper. Place the test tube with dropper in the 150-mL beaker, to prevent spillage.

13. With the 2 mL plastic dropper still inserted into the test tube, squeeze the bulb of the dropper completely flat several times to completely fill the dropper with the CO2 gas that is in the test tube. You want to collect CO2 gas only in the dropper and not any of the solution.

14. Without depressing the bulb. gently, quickly, remove the 2 mL plastic dropper from the test tube, wipe the stem clean with tissue and insert the stem of the 2 mL dropper into the sample vial with the pH sensor and 6 mL tap water. Push the plastic dropper into the sample vial along the side wall of the pH sensor, using the stem (not the bulb) to push the tip of the dropper through the tap water to the bottom of the sample vial.

15. Very slowly completely squeeze the bulb of the dropper so that the CO2 gas slowly bubbles through the tap water. Do this 8 to 10 times so that the tap water drawn up into the dropper is repeatedly contacted with the residual CO2 gas in the dropper.

16. Hold the bulb on the plastic 2 mL dropper squeezed tightly as you withdraw the stem of the dropper from the sample vial, so that no solution remains in the stopper. Very slowly, gently stir the solution (tap water with gas bubbled through) in the glass sample vial with the pH sensor. Look at the graph which has been monitoring the pH during Steps #9 through #14. Observe the constant pH of the tap water only at the earlier times, the decrease in pH as the CO2 gas was bubbled through and the falling value of the pH with the slow gentle stirring of the solution. Continue to follow the falling pH with very gentle stirring until the pH becomes constant (may be 6 to 8 minutes or more).

17. When the pH becomes constant, you may click on “Stop” to stop collecting data, if the 1000 second time has not already elapsed.

18. Record on your data sheet the best constant initial value of the pH of the tap water just before you inserted the dropper with CO2 gas (initial pH value, before CO2 was

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added) and the best constant final value for the pH after adding the CO2 gas, followed with the gentle stirring (final pH value after CO2 was added and pH stabilized). You will find using the “x=?” button below the menu at the top of the screen is a good help in determining the pH values that you must record from the constant regions on the graph. If you wish to examine these values another way, click the Statistics button, , and examine the minimum and maximum values in the pH box displayed on the graph. Make certain that you have recorded the initial and final pH values in your data table.

19. After recording your data, remove the pH sensor from the glass sample vial, rinse it with deionized water (collecting the waste water in your 600 mL waste container) and return it to the 250 mL beaker half full of tap water. Set aside the CO2 labeled glass sample vial and “NaHCO3” labeled test tube for washing later.

20. To gather data on sulfur dioxide, SO2, start again like at Step #7, using the 10 mL graduated cylinder to measure 6.0 mL of clean tap water into the clean sample vial labeled “SO2”.

21. Repeat Steps similar to #10 through #19, with this time adding about 1 mL of 1.0 M HCl to the test tube containing solid NaHSO3, and appropriately using “NaHSO3“ in place of “NaHCO3 and “SO2“ in place of “CO2“. Of course, SO2 is generated in this test tube. Do not breathe the SO2 gas.

22. To gather data on nitrogen dioxide, NO2, start again like at Step #7, using the 10 mL graduated cylinder to measure 6.0 mL of clean tap water into the clean sample vial labeled “NO2”.

23. Repeat Steps similar to #10 through #19, with this time adding about 1 mL of 1.0 M HCl to the test tube containing solid NaNO2, and appropriately using “NaNO2“ in place of “NaHCO3 and “NO2“ in place of “CO2“. Of course, NO2 is generated in this test tube. Do not breathe the NO2 gas

24. If time permits, repeat the entire experiment. Appropriately discard wastes to protect the environment as you clean the test tubes and sample vials with tap water, followed by deionized water rinse. Then start again with NaHCO3 as in step #3.

25. To finish this experiment: Rinse the pH sensor with deionized water, unscrew the cap on the plastic storage bottle, take out the 00 rubber stopper, slip the cap carefully onto the cylindrical wall of the pH sensor (be careful of the glass bulb!) and then screw the plastic bottle with the storage solution onto the pH sensor with cap so that the glass bulb is submersed in the storage solution. Be sure the cap is screwed on tightly to prevent leakage of storage solution. If some loss of storage solution has occurred so that the bulb cannot be submersed in the storage solution, please take the sensor and bottle to the instructor for special handling. The pH sensor is permanently damaged if it is not stored under the storage solution.

26. Appropriately discard wastes to protect the environment. Wash and rinse other glass apparatus and return the special glass sampling vials.


DATA





Gas Initial pH Final pH Change in pH (pH)

CO2

NO2

SO2

For repeated experiments:

Gas Initial pH Final pH Change in pH (pH)

CO2

NO2

SO2

CALCULATIONS AND DATA ANALYSIS

Show the following calculations and answer the questions in your notebook.

1. For each of the three gases, calculate the change in pH (pH), by subtracting the final

pH from the initial pH. Record these values in the Data and Calculations table.

2. In this experiment, which gas caused the smallest drop in pH?

3. Which gas (or gases) caused the largest drop in pH?

4. Coal from western states such as Montana and Wyoming is known to have a lower percentage of sulfur impurities than coal found in the eastern United States. How would burning low-sulfur coal lower the level of acidity in rainfall? Use specific information about gases and acids to answer the question.

5. High temperatures in the automobile engine cause nitrogen and oxygen gases from the air to combine to form nitrogen oxides. What two acids in acid rain result from the nitrogen oxides in automobile exhaust?

6. Which gas and resulting acid in this experiment would cause rainfall in unpolluted air to have a pH value less than 7 (sometimes as low as 5.6)?

7. Why would acidity levels usually be lower (pH higher) in actual rainfall than the acidity levels you observed in this experiment? Rainfall in the United States generally has a pH between 4.5 and 6.0.

Page 186


PRE-LAB Name_________________


Experiment 20 Acid Rain

CO2 Worksheet

1. How many gallons of gas do you use per year?

I put ______(x)_ gallons of gas every _______(y)__ days. 365x/y = ______ gallons per

year.

Convert from gallons to mL to determine mL/year. (1 gallon = 3.785 L)

2. Follow the steps to calculate how much CO2 results in one year in moles, kg, L, and

classrooms from your personal gas consumption.

a. Assume gasoline is C8H18. Write a balanced equation for the combustion reaction of

gasoline in air.

b. Determine moles of CO2 from your mL of gasoline in #1 using the balanced equation.

(dC8H18 = 0.75 g/mL)

c. Determine kg of CO2 from b.

Page 188

d. Determine Volume (in L) using PV = nRT (assume 20oC and 0.75 atm).

e. Determine the number of classrooms this volume of CO2 would completely fill

(assume 1200 sq.ft. x 10ft room).

3. The US uses about 21 million barrels of oil per day. The world used about 86 million

barrels of oil per day. Calculate how many kg and L of CO2 are produced in the world

each day from gasoline. (Neglect all natural gas, coal and biofuels burned and any

energy used to process long chains into shorter chain hydrocarbons) (1 barrel = 42

gallons).


Experiment 21

Chemical Equilibrium: Finding a Constant, Kc

The purpose of this lab is to experimentally determine the equilibrium constant, Kc, for the following chemical reaction:

Fe3+(aq) + SCN

–(aq) FeSCN

2+(aq)

iron(III) thiocyanate thiocyanoiron(III)

When Fe3+ and SCN- are combined, equilibrium is established between these two ions

and the FeSCN2+ ion. In order to calculate Kc for the reaction, it is necessary to know the

concentrations of all ions at equilibrium: [FeSCN2+]eq, [SCN

–]eq, and [Fe

3+]eq. You will

prepare four equilibrium systems containing different concentrations of these three ions.

The equilibrium concentrations of the three ions will then be experimentally determined.

These values will be substituted into the equilibrium constant expression to see if Kc is

indeed constant.

In order to determine [FeSCN2+

]eq, you will use a Vernier spectrometer. The FeSCN2+

ion produces solutions with a red color. The computer-interfaced spectrometer measures

the amount of light absorbed by the colored solutions (absorbance, A). By comparing the

absorbance of each equilibrium system, Aeq, to the absorbance of a standard solution,

Astd, you can determine [FeSCN2+]eq. The standard solution has a known FeSCN

2+

concentration.

To prepare the standard solution, a very large concentration of Fe3+

will be added to a

small initial concentration of SCN– (hereafter referred to as [SCN

–]i. The [Fe

3+] in the

standard solution is 100 times larger than [Fe3+] in the equilibrium mixtures. According

to Le Chatelier's principle, this high concentration forces the reaction far to the right,

using up nearly 100% of the SCN– ions. According to the balanced equation, for every

one mole of SCN–

reacted, one mole of FeSCN2+ is produced. Thus [FeSCN

2+]std is

assumed to be equal to [SCN–]i.

Assuming [FeSCN

2+] and absorbance are related directly (Beer's Law), the concentration of FeSCN

2+ for any of the equilibrium systems can be found by:

[FeSCN2+]eq =

Aeq

Astd X [FeSCN

2+]std

Knowing the [FeSCN

2+]eq allows you to determine the concentrations of the other two ions at equilibrium. For each mole of FeSCN

2+ ions produced, one less mole of Fe3+ ions

will be found in the solution (see the 1:1 ratio of coefficients in the equation on the previous page). The [Fe

3+] can be determined by:

[Fe3+]eq = [Fe

3+]i – [FeSCN2+]eq

Page 190

Because one mole of SCN- is used up for each mole of FeSCN2+ ions produced, [SCN

–]eq

can be determined by:

[SCN–]eq = [SCN

–]i – [FeSCN

2+]eq

Knowing the values of [Fe

3+]eq, [SCN–]eq, and [FeSCN

2+]eq, you can now calculate the value of Kc, the equilibrium constant.

Materials

Windows PC 0.0020 M KSCN Vernier computer interface 0.0020 M Fe(NO3)3 (in 1.0 M HNO3) Logger Pro 0.200 M Fe(NO3)3 (in 1.0 M HNO3) Vernier Spectrometer four pipets 1 plastic cuvette pipet bulb or pipet pump five 20 X 150 mm test tubes three 100-mL beakers thermometer tissues (preferably lint-free)

PROCEDURE

1. Obtain and wear goggles.

2. Label four large test tubes 1-4. Pour about 30 mL of 0.0020 M Fe(NO3)3 into a clean, dry 100-mL beaker. Pipet 5.0 mL of this solution into each of the four labeled test tubes. Use a pipet pump or bulb to pipet all solutions. CAUTION: Fe(NO3)3 solutions in this experiment are prepared in 1.0 M HNO3 and should be handled with care. Pour about 25 mL of the 0.0020 M KSCN into another clean, dry 100-mL beaker. Pipet 2, 3, 4 and 5 mL of this solution into Test Tubes 1-4, respectively. Obtain about 25 mL of distilled water in a 100-mL beaker. Then pipet 3, 2, 1 and 0 mL of distilled water into Test Tubes 1-4, respectively, to bring the total volume of each test tube to 10 mL. Mix each solution thoroughly with a stirring rod. Be sure to clean and dry the stirring rod after each mixing. Measure and record the temperature of one of the above solutions to use as the temperature for the equilibrium constant, Kc. Volumes added to each test tube are summarized below:

Test Tube

Number Fe(NO3)3

(mL) KSCN (mL)

H2O (mL)

1 5 2 3

2 5 3 2

3 5 4 1

4 5 5 0

3. Prepare a standard solution of FeSCN

2+ by pipetting 18 mL of 0.200 M Fe(NO3)3 into a large test tube labeled “5”. Pipet 2 mL of 0.0020 M KSCN into the same test tube. Stir thoroughly.

4. Prepare a blank by filling a cuvette 3/4 full with distilled water. To correctly use a cuvette, remember:

All cuvettes should be wiped clean and dry on the outside with a tissue.

Handle cuvettes only by the top edge of the ribbed sides.

All solutions should be free of bubbles.


6. Calibrating and Determining the Maximum Wavelength of Absorbance (max) for the FeSCN

2+ solution.

a. Prepare the Vernier Spectrometer by plugging in the USB cable and opening the

Logger Pro software. If the software doesn’t immediately recognize the Spectrometer,

choose Connect Interface Spectrometer Scan for Spectrometers from the

Experiment menu. Allow the Spectrometer to warm up for 3 minutes before taking

readings.

b. Calibrate the spectrometer by choosing Calibrate Spectrometer from the

Experiment menu. Follow the instructions from the dialog box to complete the

calibration using your blank cuvette. You will be asked to insert the blank cuvette into

the cuvette slot. Insert it in such a way that the spectrometer light goes through the

smooth sides and not the ribbed sides of the cuvette. Click “Ok.”

d. Rinse the sample cuvette with two or three small portions of the standard FeSCN2+

solution in test tube #5. Fill the cuvette about ¾ full with this the #5 standard solution.

Place the cuvette in the slot. Click “Collect”. An absorbance curve should appear on the

screen. After viewing the absorbance curve, hit “Stop”.

e. Click on the “Configure Spectrometer Data Collection” icon, located on the right hand

side of the toolbar to open the display. (The button looks like a rainbow graph.) Click

Abs. vs. Concentration (under Set Collection Mode). The wavelength of the maximum

absorbance will be automatically selected. Double check that ~470 nm (max) is the

only wavelength that is selected. Click “Ok” to close the display.

f. Record the absorbance value for standard solution #5. The last digit may fluctuate so

do your best to find the average.

7. You are now ready to collect absorbance data for the four equilibrium systems.

a. Fill a cuvette about ¾ full with the equilibrium solution in Test Tube 1. Insert the

cuvette into the slot. Record the absorbance. Repeat this step for each of the

solutions in Test Tubes 2 – 4.

b. Dispose of all solutions as directed by your instructor.

Page 192

DATA





Absorbance Trial 1

_______

Trial 2

_______

Trial 3

_______

Trial 4

_______

Absorbance of standard (Trial 5)

_______

Temperature

_______ °C

Kc expression Kc =

[Fe3+

]i

[SCN–]i

[FeSCN2+

]eq

[Fe3+

]eq

[SCN–]eq

Kc value

Average of Kc values

Kc = ________ at ________°C


CALCULATIONS/DATA ANALYSIS

1. Write the Kc expression for the reaction in the Data and Calculation table.

2. Calculate the initial concentration of Fe3+

, based on the dilution that results from adding KSCN solution and water to the original 0.0020 M Fe(NO3)3 solution. See Step 2 of the procedure for the volume of each substance used in Trials 1-4. Calculate [Fe3+]i using the equation:

[Fe3+

]i = Fe(NO3)3 mL

total mL X (0.0020 M)

This should be the same for all four test tubes.

3. Calculate the initial concentration of SCN–, based on its dilution by Fe(NO3)3 and

water:

[SCN–]i =

KSCN mL

total mL X (0.0020 M)

In Test Tube 1, [SCN–]i = (2 mL / 10 mL)(.0020 M) = .00040 M. Calculate this for

the other three test tubes.

4. [FeSCN2+

]eq is calculated using the formula:

[FeSCN2+

]eq = Aeq

Astd X [FeSCN

2+]std

where Aeq and Astd are the absorbance values for the equilibrium and standard test

tubes, respectively, and [FeSCN2+

]std = (1/10)(0.0020) = 0.00020 M. Calculate [FeSCN

2+]eq for each of the four trials.

5. [Fe3+

]eq: Calculate the concentration of Fe3+

at equilibrium for Trials 1-4 using the equation:

[Fe3+

]eq = [Fe3+

]i – [FeSCN2+

]eq

6. [SCN–]eq: Calculate the concentration of SCN- at equilibrium for Trials 1-4 using the

equation: [SCN

–]eq = [SCN

–]i – [FeSCN

2+]eq

7. Calculate Kc for Trials 1-4. Be sure to show the Kc expression and the values

substituted in for each of these calculations.

8. Using your four calculated Kc values, determine an average value for Kc. How constant were your Kc values?

Page 194


Experiment 24

Measuring Sulfur Dioxide in Wine

The problem of preserving food is as old as civilization. Centuries ago, people learned

that the fumes of burning sulfur inhibited the browning and rotting of fruits and

vegetables. In the 19th century, the practice of using sulfur dioxide to preserve meat and

fish became widespread, to the point that it became of concern to health scientists who

pressed for regulations on its use.

Adding sulfur dioxide to food can be done either by exposing the material directly to the

fumes of burning sulfur, by dipping the food in a solution containing sulfite ions, or by

adding sulfites directly to the food, as is done in winemaking. In the latter two cases, acid

in the food reacts with the sulfites to produce free sulfur dioxide. Sodium hydrogen

sulfite (sodium bisulfite, NaHSO3) was commonly used for a long time, because it is

more stable than sodium sulfite (Na2SO3) and produces, gram for gram, more SO2.

Another form of the sulfite ion, the metabisulfite ion (S2O52-

), is even more stable and

produces slightly more SO2. However, the concern about sodium in foods has led to the

more common use of the potassium salt when additions are made directly to the food.

24.1 SULFUR DIOXIDE AND HEALTH

Food that has been freshly or heavily sulfited may cause sneezing and mild shortness of

breath. In an attempt to learn the extent of sulfur dioxide as a health problem, first rats

and dogs, and then humans, were fed increasingly larger doses of sulfites. After three

generations of rats which were fed drinking water containing 750 parts per million (ppm)

of sulfites, no abnormal effects were observed. At levels of 4-6 grams a day added to

their diet, humans would suffer abdominal pain and vomiting, but no other symptoms and

no permanent effects. On the other hand, some individuals seem to be allergic to sulfites,

and thus concern about their use remains. Regulations of the United States government

state that finished wines should not contain more than 350 ppm of sulfur dioxide.

24.2 THE COMMERCIAL IMPORTANCE OF SULFITES IN WINE

Wine has been described as one of the steps on the way to making vinegar. Until recently,

with the introduction of millipore filtering techniques, it was impossible to make wine

that would last very long unless sulfur dioxide was added, because of the difficulty in

removing every last yeast cell that might renew fermentation in wine that had been

sweetened, or every microbe that might lead to the conversion of alcohol to vinegar

(acetic acid). In addition to its low toxicity, sulfur dioxide has the advantage of

possessing more than one preservative action. It kills or inhibits microbes, such as

bacteria and molds. It deactivates the enzymes which are released from damaged fruit and

vegetables and are responsible for browning and discoloration. And, to a small extent, it

combines with oxygen that has entered from the air and prevents oxidation that might

otherwise affect flavor.

Page 196

24.3 SULFUR DIOXIDE LEVELS IN WINE

In winemaking, a quantity of sulfur dioxide is added to

the freshly crushed grapes to kill yeasts and molds that

are naturally present on the grapes. The yeasts that are

added to actually make the wine have been bred over the

years to tolerate high SO2 concentrations, typically

80-100 milligrams per liter (or, as it is often expressed,

80-100 per million (ppm)). During the rest parts of the

fermentation and finishing process, and during bulk

storage, the concentration needs to be maintained at 30-50 mg/L or higher to protect the

wine from deterioration by air and microbes. Finally, at the time of bottling, additional

sulfur dioxide is added to attain a level of 30-40 mg/L. In the case of dessert wines which

are likely to be consumed more slowly after the bottle is opened, or any wine which will

probably be used more casually, the sulfur dioxide level added to the bottle might need to

be 70-80 mg/L.

Once the sulfur dioxide has been added, it is at the mercy of environmental conditions. It

will vary with the pH of the wine, the storage temperature, ethanol (alcohol) content,

micronutrient levels, and the sanitary condition of the wine. Much of it bonds with other

chemicals in the wine, such as aldehydes, glucose and ketones, and in this form it is

“fixed”, that is, unavailable to function as a preservative. Some of the remaining “free”

sulfur dioxide evaporates directly into the atmosphere and is lost. Some is oxidized to

sulfates. As a result of these interactions, the level of sulfur dioxide in newly bottled wine

quickly declines to 20-30 mg/L. From that point, it gradually continues to dissipate, and

in older wines may persist in concentrations of 5-10 mg/L.

On the other hand, while a certain level of sulfur dioxide is necessary to preserve the

wine, the winemaker must be careful to regulate the amount present in his or her creation.

A certain amount of it will dissipate during aging, and the winemaker must be able to

trace its disappearance in order to know how much to add. Too much of it will alter the

desired aroma and flavor, or interfere with aging. Some delicate wines might be ruined by

a level of only 25 mg/L. Most individuals can recognize its odor at levels of 15-40 mg/L,

and for that reason many wine drinkers prefer to let their wines “breathe” before they

drink them by pouring them from the bottle into a decanter to allow some of the gas to

escape. Some critical judges can detect free sulfur dioxide at concentrations of 5-10

mg/L.

Furthermore, the laws of different countries limit the amount of sulfur dioxide that is

allowable in wine, and the commercial winemaker must be careful not to exceed the legal

limits.

24.4 THE CHEMISTRY OF SULFUR DIOXIDE

Sulfur dioxide dissolved in water can exist as sulfur dioxide molecules (SO2), or it can


form the hydrogen sulfite ion (HSO3-) and the sulfite ion (SO3

2-). There is an equilibrium

among these various forms of sulfur dioxide, depending on the amounts present, the pH,

and the temperature.

Gaseous sulfur dioxide, which is in equilibrium with dissolved sulfur dioxide, is easily

lost by dissipation into the atmosphere, thus presenting an inherent source of error in any

attempt to measure SO2. The higher the temperature, the quicker the gas escapes.

SO2 (aq) ⇋ SO2 (g)

Sulfur dioxide reacts with water molecules to form a weak acid, the hydrogen sulfite ion:

SO2 (aq) + H2O ⇋ H+ + HSO3

-

Hydrogen sulfite ions dissociate into sulfite ions and metabisulfite ions:

HSO3- ⇋ H

+ + SO3

2-

2 HSO3- ⇋ H2O + S2O5

2-

In an acid environment, the equilibrium is shifted to the left.

Thus, at a low pH, SO2 molecules predominate, as illustrated

by the graph at the right. At higher pH levels, SO32-

increases.

In between, HSO3- is the prevalent species.

24.5 MEASURING SULFUR DIOXIDE IN WINE

For reasons of health, economics, aesthetics, and compliance with government

regulations, it is important for the successful winemaker to monitor carefully the sulfur

dioxide content of his or her wine. While the measurement of sulfur dioxide in a pure,

uncomplicated sulfite solution is comparatively easy, the situation is quite different in

wine. Wine is a complex mixture of hundreds of chemical compounds present in varying

proportions. In wine, sulfur dioxide reacts with acetaldehyde (CH3CHO) to form

acetaldehyde--hydroxysulfonate:

CH3CHO + HSO3- ⇋ CH3CHOHSO3

-

It also reacts with certain sugars, (such as glucose, C6H12O6), organic acids (such as

pyruvic acid, CH3COCOOH) and phenolic compounds (which contain the phenol

molecule, C6H5OH). Sulfur dioxide bound with these molecules is not readily measured.

Page 198

But the winemaker needs to measure the total sulfur dioxide content -- both the free,

volatile sulfur dioxide, and the fixed or bound.

The free sulfur dioxide can be measured by titrating the wine with a standard solution of

iodine. Although elemental iodine is not very soluble in water, it can be made more

soluble by adding iodide ion to the solution. The iodine enters the solution as the triiodide

ion, I3-:

I2 + I- I3

-

Free sulfur dioxide reacts with water, to form the sulfite ion:

SO2 + H2O 2H

+ + SO3

2-

The brown triiodide ion oxidizes the sulfite, and becomes the colorless iodide ion:

H2O +

SO32-

+ I3- SO4

2- + 3I

- + 2H

+

When starch is added to the mixture, the slightest trace of triiodide ion at the endpoint

produces a conspicuous, deep blue-black color, the color of a starch-iodine complex.

However, iodine also oxidizes polyphenols, a class of compounds which are always

present in wines, but especially in red wines. Since this reaction also consumes some

iodine and produces an artificially high reading, sulfuric acid is added to the wine to

reduce the interaction between iodine and polyphenols.

After the free sulfur dioxide is measured in this way, the fixed sulfur dioxide can be

released by adding sodium hydroxide. The mixture then is reacidified, and the titration is

continued.

However, while the theory is simple, the practice is elusive. Because of its volatility,

some free sulfur dioxide is likely to escape as soon as the bottle is opened, when the

sample is measured out, and during the actual titration. During the titration some sulfur

dioxide may react with atmospheric oxygen. Some chemical other than sulfur dioxide,

such as sugars, aldehydes, ascorbic acid (Vitamin C), and the polyphenols mentioned

above, react with iodine and produce an artificially high measurement. The presence of

other chemicals in the wine causes the easily recognizable starch-iodine complex to

decompose, and the endpoint quickly fades, leading the chemist to puzzle about whether

he or she has in fact reached the endpoint. Thus, the analysis of a complex mixture is

subject to numerous sources of experimental error, which must be controlled by the

chemist’s experience and knowledge of the reactions involved. This is where the chemist

must be an artist as well as a technician.


24. 6 ANALYZING A POTASSIUM METABISULFITE SOLUTION

Obtain about 50 mL of a standardized 0.002 M I3- (aq) solution in your clean, dry 100-

mL beaker. Cover the beaker with a large square of Parafilm to decrease air movement

over the solution, which will increase the loss of iodine to the atmosphere. Record the

precise molarity of the solution.

Clean a buret and rinse it with very small portions of the 0.002 M triiodide solution to

displace any water in the buret. Fill the buret with triiodide solution and drain out just

enough to fill the buret tip and bring the level of the solution just below the ‘0’ mark.

Pour about 75 mL of a potassium rnetabisulfite (K2S2O5) solution of known concentration

into your clean, 150-mL beaker. After reviewing (and practicing, if necessary) correct

pipetting techniques, use a volumetric pipet to transfer 25.00 mL of the solution into a

clean (but not necessarily dry) 250-mL conical flask. Use your wash bottle to rinse down

the inside of the flask with a minimum of demineralized water.

Because the metabisulfite ion tends to be oxidized by air, you must know what you are

doing and proceed quickly and efficiently.

S2O52-

+ H2O ⇋ 2H+ + 2SO3

2-

2SO32-

+ O2 2SO42-

It is probably better not to use a stirring bar and mechanical stirrer during this titration.

Instead, swirl the flask by hand.

Read the initial level of triiodide solution in the buret. Add 2 droppersful (4-5 mL) of 1%

starch indicator to the flask. Rinse the inside of the flask with a small amount of water.

Control your buret with one hand. Hold the titration flask with the other. When you

begin adding the triiodide solution to the metabisulfite, you will notice a blue-black color

appearing at the moment of contact of the two solutions.

This is the starch-iodine complex which will immediately disappear as metabisulfite ions

react with the iodine. Swirl the titration mixture after each addition of triiodide just until

the blue-black color disappears. When you notice that the color of the complex begins to

linger before disappearing, add iodine solution very cautiously until the addition of one

drop, or even less than one drop, causes the entire mixture to remain a pale blue in color.

At this point, rinse down the inner walls of the flask and any triiodide solution clinging to

the buret tip. If the color persists for 10 seconds or more, record the final level of

triiodide solution.

Calculate the molar concentration of S2O52-

, using the equation.

3 H2O + S2O52-

+ 2 I3- 2 SO4

2- + 6 I

- + 6 H

+

At this point, you may want to check with your instructor to learn if your determination is

reasonable. Then repeat the titration two more times and average your results.

Page 200

Next, obtain a sample of potassium metabisulfite solution of unknown concentration and

repeat the titration.

24.7 MEASURING THE FREE SULFUR DIOXIDE IN WINE

This time, it will not be as easy to ascertain the endpoint of the titration, because of the

tendency of the blue-black color to fade as the equilibrium shifts and iodine in the

complex is reclaimed by competing reactions in the mixture.

You will titrate two samples of the same wine. The purpose of the first sample will be to

ascertain semi-quantitatively the approximate volume of triiodide required to reach an

endpoint. You will use the second sample to titrate more quickly by adding most of the

triiodide all at once, then more meticulously approaching the final endpoint.

From your instructor, obtain your sample of wine. Using a clean, dry pipet, or one that

has been previously rinsed with the same sample of wine, transfer 50.00 mL of wine into

a clean flask. Wash down the inside walls of the flask with demineralized water.

Add 10 mL of 6 M H2SO4, which will inhibit the oxidation of polyphenols by iodine.

Then add 2 droppersful of starch indicator.

Begin titrating as before. To speed things up, you can add about 0.5mL of iodine at a

time, swirling the flask vigorously after each addition. At first, the appearance and

clearing of the blue-black color will be similar to that which you observed when titrating

the known and unknown metabisulfite solutions. However, as you approach the endpoint

this time, you will notice that the blue-black color will change to a pale pink color just

before the mixture clears. At this point, begin adding one drop of triiodide at a time, then

swirl the flask just until the colors fade back to the original straw color. Continue to do

this until the pink persists for 5 seconds. You may stop titrating when you decide that the

pink color persists for about 5 seconds, regardless of whether it subsequently fades away.

Record your final buret reading.

Now add enough 6 M NaOH to change the mixture from a pale yellow to a deep yellow

color, probably 10-30 mL, or 5-15 droppersful. The pH of the mixture should be about

13. The NaOH will hydrolyze compounds to which SO2 is bound, freeing the SO2.

Stopper the flask, and allow it to stand for about 15 minutes, or until you are ready to

return to it.

Refill your buret. At this point, you may acidify your second 50.00-mL sample of wine

and add the starch indicator. Read the initial level on your buret, and run in all but about

one milliliter of the volume of triiodide solution required by the first sample. The

sample will probably turn a deep blue-black color, but should fade to clear as soon as you

swirl it. Then titrate drop by drop until the transitory blue-black color is replaced by the

pink hue that persists for 5 seconds, no longer. Don’t forget to record your final buret

reading. This volume of titrant will allow you to calculate the free SO2 in the wine, using

the equations:


.

H2O + SO32-

+ I3- SO4

2- + 3I

- + 3H

+

SO2 + H2O H+ + HSO3

-

Now add as much 6 M NaOH as you used in the first sample. Stopper the second sample

and allow it to stand for 15 minutes or until you are ready to return to it.

24.8 MEASURING THE FIXED SULFUR DIOXIDE IN WINE

Continue the analysis of your first sample. Add enough 6 M H2SO4 to return the color of

the mixture to the original pale yellow color, probably 5-10 mL, or 2-4 droppersful. If it

doesn’t, add a few milliliters more acid, or enough to return to the pale yellow. The pH of

the mixture should be about 3. You don’t need to add more starch indicator. Read your

buret, then titrate as you did before, looking for the pale pink hue that replaces the blue-

black and persists for 5 seconds. Take a final reading.

Since you will need about the same volume of triiodide to titrate your second sample,

check to be sure you have enough triiodide in the buret. After the second sample has

stood for at least 15 minutes, repeat the titration more carefully with your second sample.

This volume of titrant will allow you to calculate the fixed or bound SO2 in the wine.

24.9 CALCULATIONS

1. Determine the number of moles of free SO2 that reacted with the volume of

triiodide for both samples of wine.

2. Determine the molarity of the free SO2 in both samples of wine.

3. The molarity tells you the number of moles of SO2 per liter of wine, or moles per 1000

mL. Enologists commonly express the concentration of SO2 not in moles but in grams.

Convert moles per liter to grams per liter.

4. Furthermore, enologists express the concentration of SO2 not in parts per

thousand (ppt) but in parts per million (ppm). “Grams per liter” is the same as “grams per

1,000 mL, or ppt. Express the concentration of SO2 in parts per million, either as grams

per million parts of wine or as milligrams per liter of wine.

5. Determine the concentration of fixed SO2 in the sample of wine.

6. Express the total SO2 content of the wine in ppm.

Page 202

Acknowledgements

This laboratory experiment was developed with the generous assistance of the

technicians from the wine laboratory at Wente Bros. Winery, Livermore, California:

John Jeffray

Maria Coburn

Brad Buehler

The chemistry staff at Las Positas College thanks the Wente family for helping us to

further the academic development of our students by means of this interesting and

practical application of chemistry in action.

* * * * * * * * * *

Jim Adams, 5-24-95

Las Positas College, Livermore, CA

All Rights Reserved


RELATED REACTIONS

Sulfite Equilibria

(1) SO2 (aq) ⇋ SO2 (g)

(2) SO2 (aq) + H2O ⇋ H2SO3 (aq) ⇋ H+ (aq) + HSO3

- (aq)

(3) HSO3- (aq) ⇋ H+

(aq) + SO32-

(aq)

Metabisulfite Equilibria

(4) S2O52-

(aq) + H2O ⇋ 2 HSO3- (aq) ⇋ 2 H

+ (aq) + 2 SO3

2- (aq)

Iodine-sulfite Reactions

(5) 3 H2O (l) + S2O52-

(aq) + 2 I3- 2 SO4

2- (aq) + 6 I

- (aq) + 6 H

+ (aq)

(6) H2O (l) + SO32-

(aq) + I3- (aq) SO4

2- (aq) + 3 I

- (aq) + 2 H

+ (aq)

Iodine Reactions

(7) I2 (aq) + I- (aq) I3

- (aq)

(8) I2 (aq) ⇋ I2 (g)

(9) I2 (aq) + H2O ⇋ HIO + H+ + I

-

(10) 2 HIO 2 H+ + I

- (aq) + O2 (g)

(11) 3HIO + 3 OH- 2 I

- + IO3

- + 3 H2O

(12) 4 I- + 4H

+ + O2 2 I2 + 2 H2O

Sulfite Complexes

(13)

CH3

C

O

H HSO3- CH

3C

OH

SO3-

H

+

Page 204

THE CHEMICAL COMPOSITION OF GRAPES AND WINE

Ripe wine grapes are about 71 - 80 % water. Dissolved in that water are various sugars (mainly glucose, fructose, and

sucrose) to an extent of 10 - 25 % by weight, depending on the ripeness of the fruit. Another 2-10 % consists of

organic acids (such as tartaric, malic, succinic, citric, lactic and acetic), about 2-3 % minerals (calcium, phosphorus,

iron, aluminum, magnesium, manganese, iron, sodium and potassium), various anions (chloride, silicate, carbonate,

sulfate, phosphate), and 0.5-1% other compounds (proteins, ails, vitamins, esters, enzymes, starch, pectins, gums,

tannins and pigments). Furthermore, not all grapes are identical. Different varieties of grapes have their own distinctive

flavor components, present in very small amounts, some of which are more desirable than others in the production of

flavorful wine.

The addition of yeast to crushed grapes stimulates the living yeast cells to begin metabolizing the sugars in the juice,

which is called “must”. The bulk of this metabolism is the conversion of sugars to ethanol (ethyl alcohol), carbon

dioxide, and heat. However, the yeast cells also release small amounts of other metabolic products, and numerous other

chemical reactions take place during the fermentation, depending on the strain of yeast, environmental conditions, and

the presence of air, bacteria, and other chemical and biological factors present on the grapes and on the equipment.

Some of these additional reactions and factors are encouraged in order to produce a beverage that is pleasant to the

taste, while others are carefully inhibited or controlled.

Thus, to say that wine is a mixture, not a pure substance, is an understatement. The following is a list of most of the

chemicals that have been detected in wine. Some of these affect any attempt to measure accurately the sulfur dioxide

content of wine.

Hydrocarbons

Myrcene Alcohols

methanol

ethanol

1-propanol

2-propanol

1 –butanol

2-butanol

2- methylpropanol

2-methyl-2-propanol

3-methylthiopropanol

(-)2,3-butanediol

furfuryl alcohol

isopentyl alcohol

2-methyl-1-butanol

1 -pentanol

2-pentanol

cis-3-hexen-1-ol

1-hexanol

2-hexanol

3-methyl-1-pentanol

4-methyl-1-pentanol

benzyl alcohol

1 -heptanol

2-heptanol

2-phenethanol

p-hydroxyphenylethanol

1 -octanol

1 -nonanol

2-nonanol

glycerol

2-indolylethanol

2-bornanol

Linalool

Citranellol

1-decanol

2-decanol

1 –undecanol

1-dodecanol

2-dodecanol

-terpineol

3,7-dimethyl-1,5,7-octatrien-3-ol

2-ethylhexanol

trans-2-hexen-1-ol

trans-3-hexen-1-ol

3-octanol

(±)2,3-butanediol

1-octen-3-ol

1-phenylethanol

nerol

geraniol

4-terpineol

farnesol Aldehydes

acetaldehyde

propionaldehyde

2-methylpropanal

2-furaldehyde

Isovaleraldehyde

Hexenal

2-hexenal

Benzaldehyde

Vanillin

Cinnamaldehyde

phenylacetaldehyde Ketones

-ionone

-ionone

2-nonanone

3-hydroxy-2-butanone

2,3-butanedione

acetone

2-butanone Acids

formic

oxalic

acetic

pyruvic

3-hydroxypyruvic


propionic

lactic

succinic

malic

butyric

isobutyric

tartaric

citric

2-oxoisovaleric

4-(methylthio)-2-oxobutyric

glutaric

2-methylbutyric

3-methylbutyric

2-hydroxy-3-methylbutyric

pentanoic

tricarballylic

3-methyl-2-oxopentanoic

4-methyl-2-oxopentanoic

hexanoic

4-methylpentanoic

2-hydroxyhexanoic

2-hydroxy-4-methylpentanoic

benzoic

4-hydroxybenzoic

salicylic

protocatechuic

gallic

heptanoic

phenylacetic

vanillic

octanoic

3-hydroxyoctanoic

p-hydroxyphenylpyruvic

1-phenyllactic

syringic

azelaic

nonanoic

9-decenoic

decanoic

undecanoic

dodecanoic

tetradecanoic

pentadecanoic

hexadecanoic

heptadecanoic

2-furoic

2-oxoglutaric Esters

ethyl formate

ethyl acetate

ethyl propionate

propyl acetate

isopropyl acetate

ethyl 3-hydroxypropionate

ethyl lactate

1,3-propanediol monoacetate

ethyl acetoacetate

ethyl 4-oxobutyrate

ethyl acid succinate

diethyl oxalate

ethyl acid malate

ethyl acid tartrate

butyl acetate

ethyl butyrate

ethyl isobutyrate

isobutyl acetate

ethyl pyruvate

ethyl 3-hydroxybutyrate

ethyl 4-hydroxybutyrate

ethyl isovalerate

ethyl valerate

pentyl acetate

isopentyl acetate

2-methylbutyl acetate

ethyl 2-hydroxy-3-methylbutyrate

methyl salicylate

diethyl succinate

diethyl malate

diethyl tartrate

ethyl hexanoate

hexyl acetate

isobutyl isobutyrate

ethyl-2-hydroxy-4-methylpentanoate

isopentyl lactate

2-phenethyl formate

ethyl heptanoate

isobutyl valerate

dimethyl phthalate

2-phenethyl acetate

ethyl octanoate

hexyl butyrate

hexyl butyrate

isobutyl hexanoate

isopentyl isovalerate

ethyl isopentyl succinate

ethyl 3-phenyllactate

ethyl nonanoate

hexyl valerate

hexyl isovalerate

isopentyt hexanoate

2-methylbutyl hexanoate

propyl octanoate

diethyl phthalate

2-phenethyl butyrate

ethyl 9-decenoate

ethyl decanoate

hexyl hexanoate

isobutyl octanoate

ethyl undecanoate

isopentyl octanoate

(-)2,3-butanediol monoacetate

(±)2,3-butanediol monoacetate

2-phenethyl hexanoate

diisopentyl succinate

ethyldodecanoate

hexyl octanoate

isobutyl decanoate

isopentyl decanoate

2-methylbutyl pentanoate

2-phenethyl octanoate

ethyl tetradecanoate

Page 206

ethyl pentadecanoate

isopentyl dodecanoate

2-methylbutyl dodecanoate

ethyl hexadecanoate

isopentyl tetradecanoate

2-methylbutyl tetradecanoate

ethyl methylmalate

2,6,6-trimethyl-2-vinyl-4-acetoxytetrahydropyran Carbohydrates

glucose

fructose

mannitol Lactones

-butyrolactone

4-acetyl-4-hyroxybutyric acid- -lactone

pantolactone

4,5-dihydroxyhexanoic acid- -lactone

4-carboethoxy-4-hydroxybutyric acid- -lactone

trans-4-hydroxy-3-methyloctanoic acid- -lactone

2-methyl-4-hydroxybutyric acid- -lactone

4-hydroxypentanoic acid- -lactone

4-hydroxyhexanoic acid- -lactone

4-hydroxynonanoic- -lactone

2-vinyl-2-methyltetrahydrofuran-5-one Acetals

2,4,5-trimethyl-1,3-dioxolane

1,1-diethoxyethane

2.4-dimethyl-5-ethyl-1,3-dioxolane

1-ethoxy-1-propoxyethane

1,1-diethoxypropane

1,1-dipropoxyethane

1,1-diethoxybutane

1,1-diethoxy-2-methylpropane

1-ethoxy-1 (3-methylbutoxy)ethane

1-ethoxy-1 (2-methylbutoxy)ethane

1,1-diethoxy-3-methylbutane

1,1-diethoxy-2-methylbutane

1-ethoxy-1-(2-phenethoxy)ethane

1,1-diisopentoxyethane

1,1-di(2-methylbutoxy)ethane

1 -isopentoxy-1 -(2-methylbutoxy)ethane

1-isopentoxy-1-(2-phenothoxy)ethane

1 -(2-methylbutoxy)-1 -(2-phenthoxy)ethane Phenols

Phenol

3-methylphenol

p-ethylphenol

2-methylphenol

p-vinylphenol

1-naphthol

acetovanillon

tyrosol

p-vinylguaiacol

Nitrogen-containing compounds

methylamine

dimethylamine

ethylamine

ethanolamine

propylamine

isopropylamine

pyrrolidine

N-ethylacetamide

butylamine

isobutylamine

putrescine

histamine

pentylamine

isopentylamine

hexylamine

methyl anthranilate

4-carboethoxy-4-aminobutyric acid- -lactam

N-acetylalanine ethyl ester

N-isopentylacetamide

2-phenethylamine

alanine

arginine

glutamic acid

proline

tyramine

N-(2-phenethyl)acetamide

N-isobutylacetamide

N-(2-methylbutyl)acetamide

N-3-(methylthio)propylacetamide Miscellaneous

Methanethiol

ethanethiol

dimethylsulfide

dimethyldisulfide

diethylsulfide

diethyldisulfide

diallylsulfide

butylethylsulfide

diisopropylsulfide

diisopropyldisulfide

di(3-methylbutyl)sulfide

cis-linalool oxide

trans-linalool oxide

2,6,6-trimethyl-2-vinyl-4-hydroxytetrahydropyran

nerol oxide

cis-roseoxide

trans-roseoxide

2-methylthiophane-3-one phthalide Gases

carbon dioxide

sulfur dioxide

hydrogen sulfide


Lab Report 24

DATA





24.6 Analyzing a Potassium Metabisulfite Solution

a. Accepted concentration of known K2S2O5 (aq) (as M) _______________

b. Volume of K2S2O5 (aq) taken (in mL) _______________

c. Concentration of standardized I3- (aq) solution (as M) _______________

d. Titration data for a sulfite solution of known concentration:

Trial 1 2 3 4

Final volume (mL)

Starting volume (mL)

Delivered volume (mL)

e. Observations:

f. Sources and sizes of errors:

g. Calculations: Molarity of known K2S2O5 (aq) based on volume of I3- (aq):

Trial 1 2 3 4

Moles of I3- (aq)

Moles of S2O52-

(aq)

M of S2O52-

(aq)

h. Average experimental value of M of S2O52-

(aq) ____________

Page 208

i. Spread: high - low

x 100% = __________average

j. Compare your experimental value to the accepted concentration of S2O52-

(aq):

average experimental M - accepted M x 100% = __________

accepted M

l. Code number of unknown K2S2O5 (aq) ___________________

m. Volume of K2S2O5 (aq) taken (in mL) ___________________

n. Concentration of standardized I3- (aq) solution (as M) ___________________

o. Titration data for a sulfite solution of unknown concentration:

Trial 1 2 3 4

Final volume (mL)



p. Observations:

q. Sources and sizes of errors:

r. Ca1cuations: Molarity of unknown K2S2O5 (aq) based on volume of I3- (aq):

Trial 1 2 3 4

Moles of I3- (aq)

Moles of S2O52-

(aq)

M of S2O52-

(aq)

s. Average experimental value of M of S2O52-

(aq): ____________

t. Spread: high - low

x 100% = __________average


24.7 Measuring the Free Sulfur Dioxide in Wine

a. Name of the wine _______________________________________________

b. Volume of wine taken each time (in mL) __________________

c. Concentration of standardized I3- (aq) solution (as M) __________________

Acidify the sample of wine with 10 mL of 6 M H2SO4

d. Titration data for free SO2:

Sample

Trial 1 2

Final volume (mL)



e. Observations:

f. Sources and sizes of errors:

g. Calculations: Molarity of free SO2 based on volume of I3- (aq):

Trial 1 2

Moles of I3- (aq)

Moles of HSO3- (aq)

M HSO3- (aq)

M of SO2 (aq)

mg SO2/L

Page 210

24.8 Measuring the Fixed Sulfur Dioxide in Wine

Hydrolyze with 10-30 mL of 6 M NaOH

Stopper and let stand for 15 minutes.

Re-acidify with 5-l0 mL 6M H2SO4

a. Titration data for fixed SO2:

Sample

Trial 1 2

Final volume (mL)



b. Observations:

c. Sources and sizes of errors:

d. Calculations: Molarity of fixed SO2 based on volume of I3- (aq):

Trial 1 2

Moles of I3- (aq)

Moles of HSO3- (aq)

M HSO3- (aq)

M of SO2 (aq)

mg SO2/L

e. Total sulfites (add the free and fixed SO2 determined for Sample 2):

mg SO2/L

Parts SO2 per million (ppm)


Reference Material

Lab Report Format

Periodic Table

Errors, Precision and Accuracy

Treatment of Experimental Data

Statistics and Uncertainty in the Laboratory

Names, Formulas and Oxidation Numbers of Some Common Ions

Net Ionic Equations

Solubility Tables (4 versions)

Colors of Ions in Aqueous Solutions

Common Oxidation States of Six Elements Important in Redox Chemistry

Activity Series of Metals and Nonmetals

Acids and Bases

Properties of Water: Density and Vapor Pressure

Las Positas College

Las Positas College

General Chemistry

Ansell/Deleray

Lab Reports

(Information which might be included)

-Experiment Title: (brief, descriptive)

-Experimenter’s Name (include lab partners)

-Date: Place:

-Purpose: (Objectives, Anticipated results, 1-2 complete sentences)

-Experimental procedures and Equipment Set Up: (brief outline, sketches and diagrams, may include by

reference, another experimenter should be able to repeat your experiment from your report.)

-Equipment: (name, type, number, accuracy, precision, limitations, etc.)

-Outside Data and References: (e.g. density or boiling points found in the Merck Index)

- Notes, Potential Problem areas, Reminders, Safety Cautions

- Ambient Conditions: IF RELEVANT (time start, time end, temperature, humidity, barometric pressure,

changes in conditions during experiment.) Record on data sheet during experiment.

- Experimental data: (Prepare data tables in advance, if possible)

(Record in INK on data sheet during experiment)

-Observations, Notes: (Record in INK during the experiment, in detail!)

----------------------------------------------------------------------------------------------------------------------------- --

-Calculations: (may be done during or after experiment, pencil okay)

-Discussion – (thoughtful explanation, analysis, discussion of results)

-Compare with theoretical/other experimental, expected, unanticipated)

-Sources of error, magnitudes of error, effects of error: (record during experiment)

-Conclusion, Recommendations: (quality of results, brief status of results and technique,

recommendations for improvements/changes in experiment, future experiments suggested,

applications.)

NOTES:

There are many ways to set up an experimental record. A record style should be appropriate for the

specific experiment.

The reports should be brief, concise and complete—one/several word notes or descriptions are desirable.

Set up and do as much as you can before starting experimental work (i.e. before lab period). Especially in a

crowded lab, completed pre-lab work is essential for safety, time.

Use ink for recording all data, sources of error, observations (pencil is okay for calculations). Immediately

record all data directly on data sheet, NOT on scratch paper, paper towels or the palm of your hand!

Record data, observations, sources of error in duplicate (using recorded lab notebooks.) Always keep a

copy for your records.

If a piece of data is recorded incorrectly, a single thin line (no erasure or obliterations) and a brief note

describing the error may be used to delete such data. Do not throw away or destroy any

data.

The experimental reports and-post experiment assignments are due at the beginning of your next lab

meeting following the scheduled completion of the experiment.

R1

Las P

osita

s Col

lege

Peri

odic

Tab

le o

f the

Ele

men

ts18

1 IAV

IIIA

11 H

1.00

82 IIA

13 IIIA

14 IVA

15 VA

16 VIA

17 VII

A

2 He

4.00

3

23 Li

6.94

1

4 Be

9.01

2

5 B10

.81

6 C12

.01

7 N14

.01

8 O16

.00

9 F19

.00

10 Ne

20.1

8

311 N

a22

.99

12 Mg

24.3

13 IIIB

4 IVB

5 VB

6 VIB

7V

IIB

8V

IIIB

9V

IIIB

10V

IIIB

11 IB12 IIB

13 Al

26.9

8

14 Si28

.09

15 P30

.97

16 S32

.07

17 Cl

35.4

5

18 Ar

39.9

5

419 K

39.1

0

20 Ca

40.0

8

21 Sc 44.9

6

22 Ti

47.8

8

23 V50

.94

24 Cr

52.0

0

25 Mn

54.9

4

26 Fe 55.8

5

27 Co

58.9

3

28 Ni

58.6

9

29 Cu

63.5

5

30 Zn 65.3

9

31 Ga

69.7

2

32 Ge

72.6

1

33 As

74.9

2

34 Se 78.9

6

35 Br

79.9

0

36 Kr

83.8

0

537 R

b85

.47

38 Sr 87.6

2

39 Y88

.91

40 Zr 91.2

2

41 Nb

92.9

1

42 Mo

95.9

4

43 Tc

(98)

44 Ru

101.

1

45 Rh

102.

9

46 Pd 106.

4

47 Ag

107.

9

48 Cd

112.

4

49 In 114.

8

50 Sn 118.

7

51 Sb 121.

8

52 Te

127.

6

53 I12

6.9

54 Xe

131.

3

655 C

s13

2.9

56 Ba

137.

3

57 La

138.

9

72 Hf

178.

5

73 Ta

180.

9

74 W 183.

9

75 Re

186.

2

76 Os

190.

2

77 Ir19

2.2

78 Pt 195.

1

79 Au

197.

0

80 Hg

200.

6

81 Tl

204.

4

82 Pb 207.

2

83 Bi

209.

0

84 Po (209

)

85 At

(210

)

86 Rn

(222

)

787 Fr (223

)

88 Ra

(226

)

89 Ac

(227

)

104

Rf

(261

)

105

Db

(262

)

106

Sg (263

)

107

Bh

(262

)

108

Hs

(265

)

109

Mt

(266

)

110

Ds

(269

)

111

Rg

(272

)

112

58 Ce

140.

1

59 Pr 140.

9

60 Nd

144.

2

61 Pm (145

)

62 Sm 150.

4

63 Eu

152.

0

64 Gd

157.

3

65 Tb

158.

9

66 Dy

162.

5

67 Ho

164.

9

68 Er

167.

3

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Las Positas College Chemistry 30A, 31 JB—7/30/01

DECISION TREE: Should a formula be written as ions?

Is the substance an acid (formula begins with H), a base (contains hydroxide ion), a salt (cation anion) or other?

If it is an acid, go to A. If it is a base, go to B. If it is a salt, go to S. If it is an other, do not write it as ions. Example: Zn stays Zn

(Note: There are three substances that decompose yielding gases when formed in chemical reactions: H2CO3, H2SO3 and NH4OH.)

A. Is it a strong acid: H2SO4(aq), HClO4(aq), HNO3(aq), HCl(aq), HBr(aq), or HI(aq)?

If yes, write it as ions. Example: H2SO4(aq) becomes 2 H+ + SO42-

If no, do not write it as ions. Example: HNO2(aq) stays HNO2

B. Is it a strong base: NaOH, LiOH, KOH, Ba(OH)2, or Ca(OH)2?

If yes, write it as ions. Example: NaOH becomes Na+ + OH-

If no, do not write it as ions. Example: Fe(OH)3 stays Fe(OH)3(Note: Most hydroxides not listed above are weak or nonelectrolytes because they are insoluble in water—always check this out when writing net ionic equations.)

S. Is it a soluble salt? Consult a table of solubilities.

If soluble, write it as ions. Example: MgCl2 becomes Mg2+ + 2 Cl-

If not soluble, do not write it as ions. Example: AgCl stays AgCl(s)

R18Las Positas College Chemistry 1A 10/15/94 Juliet Bryson

COMMON OXIDATION STATES OF SIX ELEMENTS IMPORTANT IN REDOX CHEMISTRY

Much of the redox chemistry in this course can be derived from about six elements. The six are Cl, Cr, Mn, N, O, and S. The common oxidation states of each element, in turn, will be displayed on a vertical number scale, with compounds and ions listed that exhibit each particular oxidation number. In general, compounds at the very highest oxidation level are likely to be oxidizing agents. (Certainly not reducing agents. Why?). Compounds at the very bottom are likely to be reducing agents. (Certainly not oxidizing agents. Why?) Compounds exhibiting intermediate oxidation states can go either way—it depends on the other reagent involved. If that reagent is a stronger reducing agent (RA), the first will become the oxidizing agent (OA), and vice versa.

Chlorine (Bromine and iodine behave similarly) Oxidation No Chlorine Compound Characteristics +7 (Cl2O7), HClO4, ClO4- Cl2O7 is unstable.

HClO4 is a strong OA, reduced to Cl- +5 HClO3, ClO3- strong OA, reduced to Cl- +3 HClO2, ClO2- good OA, reduced to Cl- +1 (Cl2O), HClO, ClO-, OCl- good OA, reduced to Cl- 0 Cl2 good OA, reduced to Cl- -1 Cl-

ChromiumOxidation No Chromium Compound Characteristics

acidic basic +6 Cr2O72- CrO42-

(orange) (yellow) both are strong OA’s

+3 Cr3+ Cr(OH)4-

(green or violet) (green) Cr(OH)3 is amphoteric (soluble in either acid or base)

0 Cr the metal

OxygenOxidation No Oxygen Compound Characteristics 0 O2 the ‘original’ and most abundant OA -1 H2O2, HO2- OA or RA -2 H2O, OH-

R19

ManganeseOxidation No Manganese Compound Characteristics +7 MnO4- (purple)

(permanganate ion) strong OA, reduced to a) Mn2+ in acid b) MnO2 in neutral or slightly basic c) MnO42- in strongly basic

+6 MnO42- (green) (manganate ion)

easily reduced to MnO2

+4 MnO2 brown solid, not soluble +2 Mn2+ pale pink to colorless 0 Mn the metal

NitrogenOxidation No Nitrogen Compound Characteristics +5 N2O5, HNO3, NO3- strong OA, reduced to:

a) NO2 in conc acid (> 8 M) b) NO in dil acid (<6 M), but can go all

the way to NH3 with strong OAs +4 NO2 (N2O4) brown gas +3 (N2O3), HNO2, NO2- only NO2- is stable

active as OA or RA +2 NO colorless but oxidized by O2 in air to NO2

(brown) +1 N2O supports combustion about like oxygen 0 N2 -3 NH3, NH4+ RA, but not commonly used in this

capacity in Chem 1A

SulfurOxidation No Sulfur Compound Characteristics +6 SO3, H2SO4, HSO4-, SO42- conc acid is strong OA +4 SO2, H2SO3, HSO3-, SO32- OA or RA +2 (ave!)

S2O32-

(thiosulfate ion) used as RA in analytical chemistry

0 S yellowish powder -2 H2S, HS-, S2- strong RA; usually oxidized to S

SOME COMMON REDUCING AGENTS 1) The metals: oxidized to their positive ions; e.g. Sn oxidized to Sn2+ or Zn oxidized to Zn2+

2) Ions in which the metal has another higher oxidation state; e.g. Sn2+ oxidized to Sn4+ or Fe2+

oxidized to Fe3+ or Hg22+ oxidized to Hg2+

3) Carbon and organic compounds may be oxidized to other organic compounds and to CO2 and H2O; e.g. C (coke, much used in industry) oxidized to CO or CO2 or CH3CH2OH (an alcohol, ethanol) oxidized to CH3CHO (an aldehyde, ethanal) and then further oxidized to CH3COOH ( an organic acid, ethanoic acid or acetic acid)

R20ACTIVITY SERIES OF METALS

The metals are listed in order of decreasing strength as reducing agents.

Li best RAK

Very active with H2O Ba or acids. H2 formed. Ca Oxides reduced by electrolysis, but

Na not by H2 or CO. Mg Al

Active with acids or Mn with steam. H2 formed. Zn Oxides reduced by C

Cr or Al; not by H2 or CO. Fe Cd

Ni Less active with acids. Sn H2 formed. Pb Oxides reduced by heating H2 with H2 or CO. Active only with stronger Cu oxidizing acids, as HNO3. Sb No H2 formed. Bi

Hg Oxides reduced to metal Ag (decomposed) by heat alone.

React only with aqua regia: Au HCl+HNO3, 3:1 Pt poorest RA

ACTIVITY SERIES OF NONMETALS

The nonmetals are listed in order of decreasing strength as oxidizing agents.

F strongest OA Cl O Br I S less strong

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Chemistry 1A: General Chemistry - Las Positas...Chemistry 1A