228 Winter 2012 Manual (1)

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Chemistry 228

General Chemistry Lab 2

Winter 2012

Lab Coordinators:

Dr. Eric Sheagley [email protected]

Dr. Gwen Shusterman [email protected]

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Chemistry 228 General Chemistry Laboratory

SYLLABUS – Winter 2012 Lab Packet: All printed material for this lab will be available on Blackboard OR may be purchased at Smart Copy (1915 SW 6th Avenue).

Prelab Exercises: Prelab instructions are included in the lab packet. You should answer any questions presented and prepare for the weeks lab before your lab meeting. Pre-labs are due at the beginning of the lab period.

Materials: You will need chemical splash safety goggles. These are available from the chemistry stockroom (Room 280 SRTC) or at the campus bookstore. You will need a bound carbonless copy notebook (not loose paper) for recording data. You are responsible for all laboratory equipment checked out to you. If you break glassware, you will pay the replacement cost of the glassware.

Dress for lab: You must wear shoes that cover your entire foot, including the heel. They should fit up near your ankle; leather is preferred but any non-porous material is okay. Short shorts and short skirts are not allowed. Your clothing must cover your torso and legs down to your knees.

Grading: The laboratory is graded on a Pass/No Pass basis. An average of 75% of all points available in the lab is required to pass. Late Work: Laboratory reports are due at the beginning of the lab period following completion of the experiment. Lab reports should be typed. Late reports will be docked 5 points per day late. Attendance: Attendance in this lab is mandatory. YOU MUST ATTEND ALL SCHEDULED LABORATORY MEETINGS. If you are not able to attend lab you must notify your laboratory instructor as soon as possible. Students are responsible for completing the lab report for the missed lab. Data can be obtained from a lab partner or the lab TA. The made up work should be clearly labeled and indicate the origin of the data reported. Reports are due the class meeting following the syllabus deadline. In addition to completing the make-up lab you must make up the missed lab time. The make-up laboratory will not be the same lab you missed but will be a unique activity that will take place during week 10 of the quarter, during the regularly scheduled lab period. FAILURE TO DO BOTH WILL RESULT IN A NO PASS GRADE. If you miss two or more labs your grade will be a NO PASS. NOTE: If you are more than 15 minutes late to lab you will be marked late. Two late arrivals during the term will be counted as a missed lab. In addition, late students may be assigned to lab clean up duties at the conclusion of the lab period. If you are chronically late you will be given a NO PASS.

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Plagiarism: Experiments will be done in groups sharing the computer for data analysis and acquisition. You may compare data with other groups, but the content of your lab reports MUST be written individually. It will be considered an act of plagiarism if you borrow tables or graphs from another student (learning how to properly create a table or graph is an important skill, learn how to do it on your own!). You cannot paraphrase the internet, your book or any other source without the proper reference. Additionally, it will be considered an act of plagiarism if you borrow data without prior approval from your TA. There are additional resources online to help you avoid plagiarism. Please be sure to check http://www.lib.pdx.edu/instruction/survivalguide/writeandcitemain.htm or http://web.pdx.edu/~b5mg/plagweb.html, and feel free to discuss the issue with your TA or the lab coordinator. Depending on the severity of the offense(s), you will receive, at a minimum, a zero score for the report. Additionally, a report may be made to the Office of Student Affairs.

Grading Criteria Unless otherwise noted, every lab report is worth 90 points, including the prelab, notebook and technique. Each lab report will be graded according to the following point distribution:

Prelab: 10 points

Abstract: 10 points

Introduction: 10 points

Data: 10 points

Results: 15 points

Discussion: 15 points In addition to the above points each lab meeting will have an additional 20 points assigned on the following basis:

Notebook: 10 points These points are awarded by the TA based upon the quality of your lab notebook. Your TA will be looking to see that you are including a title, a statement of purpose, the procedures, data tables and that all data is present.

Lab technique: 10 points The basis for assigning these points includes (but is not limited to) general lab technique and methods, safety, general mannerism in lab and cleanliness.

Both of these criteria will be evaluated by your TA during each lab meeting. At the end of each lab you must check out with your TA so that he or she can assess your lab notebook and verify that you have cleaned your work area

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Chemistry 228 Winter 2012 – Schedule

Week 1 Check-in. Announcements and registration adjustments and Lab Safety View the Lab Safety Video. A link is available on D2L. Complete the quiz, also available on D2L, before returning week 2 or you will not be able to participate in the lab. Enthalpy of Neutralization of Phosphoric Acid Week 2 Monday is a holiday, the observance of Martin Luther King, Jr. Day. All labs canceled for the week. Week 3 (Enthalpy of Neutralization of Phosphoric Acid report due)

Enthalpy of Reaction and Hess's Law Week 4 (Hess’s Law report due. Please write your TA’s name on the report and deposit it in the mailbox located outside of the door to the chemistry office on the second floor of SRTC)

Gas Laws – Online Lab. All lab sections do not meet. Week 5 (Gas Laws – online report due) Decomposition of Hydrogen Peroxide Week 6 (Decomposition of Hydrogen Peroxide report due)

Vapor Pressure and Heat of Vaporization Week 7 (Vapor Pressure and Heat of Vaporization report due)

Using Freezing Point Depression to Find Molecular Weight Week 8 (Molecular Weight by Freezing Point Depression Report Due) Kinetics: Rate and Order of a Chemical Reaction Week 9 (Kinetics: Rate and Order of a Chemical Reaction report due) Chemical Equilibrium: Finding a Constant Kc Week 10 (Chemical Equilibrium: Finding a Constant Kc report due)

Make up Lab Check Out and

Labs are graded on a Pass/No Pass basis.

In order to pass the lab you must turn in every lab report.

More than one absence will result in a No Pass for the class.

You must receive 75% or greater of all the points available to pass.

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Do not copy your partners, friends, old lab reports. That is plagiarism!

Laboratory Safety Rules and

Procedures

Safety Rules The guidelines below are established for your and your classmates’ personal safety. Failure to adhere to the guidelines below will result in a loss of Lab Technique points. • Personal Protective Equipment (PPE) is used to protect you from serious injuries or illnesses

resulting from contact with chemical hazards in the laboratory. Spills and other accidents can occur when least expected. For this reason it is necessary to wear proper PPE. The PPE for student labs consist of goggles, gloves and clothing. Proper PPE is required for all students or they will be asked to leave the lab

•Goggles – Goggles must be worn whenever any experimental work is being done in the laboratory to protect the eyes against splashes. Only indirect-vented goggles are allowed in the student labs and should be worn at all times when any chemical is being used in the lab. These are for sale in the bookstore and stockroom. You should not wear contact lenses in a chemical laboratory. Chemical vapors may become trapped behind the lenses and cause eye damage. Some chemicals may dissolve “soft” contact lenses. The most important aspect of having the goggles fit comfortably is the proper adjustment of the strap length. Adjust the strap length so that the goggles fit comfortably securely and are not too tight. If you find that your goggles tend to fog, you can pick-up anti-fog tissue from the stockroom. • Gloves – Gloves should be worn to protect the hands from chemicals. Gloves are provided through your student fees and are located in the student labs. For health and safety reasons it is important to always remove at least one glove when leaving the student laboratory, this prevents things such as door handles from getting contaminated. • Clothing – Dress appropriately for laboratory work. You must wear shoes that cover your entire foot, including the heel. They should fit up near your ankle; leather is preferred but any non-porous material is okay. Your clothing must cover your torso and legs down to your knees. Short shorts, short skirts, tank tops and halter tops are not allowed.

• Eating, drinking and smoking are prohibited in the laboratory at ALL times. Wash your hands

after finishing lab work and refrain from quick trips to the hall to drink or eat during lab. If you take a break, be certain to remove gloves and wash hands before ingesting food or drink.

• Never work alone in the laboratory or in the absence of the instructor.

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• Headphones may not be worn in lab.

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Safety Procedures • Know location of safety equipment; fire extinguisher, fire blanket, first aid kit, safety

shower, eyewash fountain and all exits. • In case of fire or accident, call the instructor at once. • Small fires may be extinguished by wet towels.

• If a person’s clothing catches fire, roll the person in the fire blanket to extinguish the flames. • In case of a chemical spill on the body or clothing, stand under the safety shower and flood the affected area with water. Remove clothing to minimize contamination with the chemical. • If evacuation of the lab is necessary, leave through any door that is safe, or not obstructed; doors that lead to other labs may be the best choice. Leave the building by the nearest exit and meet your TA on the field next to Hoffmann Hall. This would also be the meeting place in the event of an earthquake or other emergency. It is good to know the nearest exits of your lab on the first day of class.

• Spilled chemicals must be cleaned up immediately. If the material is corrosive or

flammable, ask the instructor for assistance. If acids or bases are spilled on the floor or bench, neutralize with sodium bicarbonate, then dilute with water. Most other chemicals can be sponged off with water.

• Avoid contact with blood or bodily fluids. Notify the instructor or stockroom personnel if

ANY blood is spilled in the lab so that proper clean up and disposal procedures may be followed.

• If a mercury thermometer is broken, do not attempt to clean up yourself. Notify students

around you, so that mercury is not spread, then notify your lab instructor or stockroom personnel. The stockroom is equipped for proper clean up and disposal of mercury.

Laboratory Procedures and Protocol General Etiquette: • Leave all equipment and work areas as you would wish to find them. • Keep your lab bench area neat and free of spilled chemicals. Your book bag, coat, etc.,

should be kept in the designated area at the entrance to the lab, not at your bench.

• All chemical waste must be disposed of in proper containers. Proper disposal of chemicals is important for student safety and proper disposal. Putting chemicals into the wrong containers can lead to injury from unexpected chemical reactions. Mixing waste can also

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make it more difficult or expensive for PSU to dispose of them. Only chemicals should go into waste jars. Waste jars for each experiment will be provided in the lab. They will be labeled specifying which contents should be placed inside. It is important that you replace the lids to the waste containers. When done with the waste jar, make sure it is placed in a secondary container. Do not put anything down the sink unless you are explicitly told to dispose of it this way. Your instructor will provide specific disposal guidelines when needed. Following these guidelines assists us in lowering the environmental impact of the labs.

There are several locations for very specific waste.

i. Chemical waste – these containers are ONLY for chemical waste generated in the lab. They are each specifically labeled for each lab and waste type. READ THE LABELS.

ii. Contaminated paper waste – this is ONLY for paper towels used for clean-up of chemical spills.

iii. Broken glass – this is ONLY for broken glassware. iv. Gloves – this is ONLY for used gloves. v. Normal trash – this is for all other trash that is not chemically

contaminated, glass, or gloves.

• Clean your bench and equipment Clean all your glassware- dirty glassware is harder to clean later. Wash with water and detergent scrubbing with a brush as necessary. Rinse well with water. Do not dry glassware with compressed air, as it is frequently oily. The water and gas should be turned off and your equipment drawer locked.

• Clean the common areas before you leave the lab. Point deductions for the entire class

will be imposed if the instructor or stockroom is not satisfied. • Return any special equipment to its proper location or the stockroom.

Handling Chemicals: Obtaining reagents:

• Read the label CAREFULLY. The Chemicals are organized by experiment in secondary containment bins. Make sure the chemical name and concentration match what is required by the experiment!

• Do not take the reagents to your bench. • We recommend always picking up bottles by the label. If all students do this, then any

unnoticed spills when pouring will not cause possible problems for the next user. Remember to wear gloves while working with reagents.

• Do not put stoppers or lids from reagents down on the lab bench. They may become

contaminated. Be sure that the lids or stoppers are replaced.

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• Do not place your own pipet, dropper, or spatulas into the reagent jar. Pour a small amount into a beaker and measure from that. Please pour on the conservative side to minimize waste and cost of labs. You can always go back for more.

• Do not put any excess reagent back in the reagent jar. Treat it as waste and dispose of it

properly. • When weighing chemicals on the balances, never weigh directly onto the weighing pan.

Weigh into a weighing boat or beaker. Any spills on the balances MUST be cleaned up immediately. If you are unclear how to clean a spill, notify your instructor. The balances you are using are precision pieces of equipment and costs up to $4000.

• All chemicals should be treated as potentially hazardous and toxic. Never taste a chemical

or solution. When smelling a chemical, gently fan the vapors toward your nose. • Any chemicals that come in contact with your skin should be immediately washed with soap

and copious amounts of water.

Laboratory Procedures • Never pipet any liquid directly by mouth! Use a rubber bulb to draw liquid into the pipet. • Never weigh hot chemicals or equipment. • When heating a test tube, always use a test tube holder and be certain never to point the

open end of the test tube toward yourself or another person. • Handling glass tubing or thermometers: to insert glass tubing into a rubber stopper,

lubricate the glass tubing with a drop of glycerin, hold the tubing in your hand close to the hole, and keep all glass pieces wrapped in a towel while applying gentle pressure with a twisting motion.

• To prepare a dilute acid solution from concentrated acid, acid should be added slowly to

water with continuous stirring. This process is strongly exothermic, and adding water to acid may result in a dangerous, explosive spattering.

• Use the fume hood for all procedures that involve poisonous or objectionable gases or

vapors. • Never use an open flame and flammable liquids at the same time.

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Keeping a Lab Notebook In keeping a lab notebook, there are certain principles that should be followed. These boil down to being clear and complete in your entries in your lab notebook. There are also certain conventions for lab notebooks that are universally followed. High on this list are the following:

Use a notebook with pre-numbered pages Record entries in ink Keep entries reasonably neat and organized Never tear pages out of your lab notebook (other than the carbonless copy pages)

What Kind of Notebook Should I Use? For this class you must use a notebook with carbonless copy pages. General Guidelines

• Write your name on outside front of notebook • Use black ink, fine-tipped ball-point pen (this will photocopy clearly) • At the front of the notebook, leave a few pages for a Table of Contents • Each lab should have a brief introduction and description of procedure • Generally use only the right hand page for most text • Use facing left page for working graphs, manual calculation, and working notes • Prepare data tables in advance - with columns for calculated results and notes • Working graphs done in lab notebook to monitor progress

Usage and Structure The overriding principle for a lab notebook is to record in it all the pertinent information about your lab work. This boils down to clear descriptions of what you did and what you observed as a result. It is a working tool, and a reference for other researchers who might want to read your notebook and reproduce your work. (This applies to notebooks in learning laboratories: Your lab instructor may want to look at what you did in order to understand your results. This is often the case. So, it needs to be clear.) The word “clear” here is crucial. In order to be clear, data must be recorded in well-thought-out tables, clearly labeled. Descriptions of procedures must be clear and concise; to the point. You should record all your work in your lab notebook. That is the proper place for all lab planning and observations. Nothing should be recorded on odd scraps of paper, etc.

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Structure for your Lab Notebooks: For each lab in this class you should have the following sections in your lab notebook: Title Purpose Procedure and Observations It is also often helpful to include a Result section Note: When preparing your notebook for lab only write on the right hand page. Title: With your lab notebook laid open, on the right hand page write down the title of the experiment, and the date. In general, you will use the right-hand page for all your writing. The left-hand page is reserved for recording scratch work. Don’t use this space until you need to. One example of how to use the left-hand page: if your work requires simple calculations using your measurements, use the left-hand page to do the calculations. If unexpected results occur later, sometimes you can look back at your scratch work and discover the error. (“Oh, I subtracted wrong! We put in 10.5 grams of copper sulfate, not 9.5 like we thought!”) Better to discover the error after the fact than never to discover it at all. Purpose: Below the title, write the purpose of the experiment in one or two sentences. This section serves to remind you and notify the reader what the experiment is about. Procedure and Observations: This next section will be labeled Procedure and Observations. As the name suggests, write down what you actually do and what you observe. This section is where you should have pre-prepared tables for data collection. Set up this section by dividing the page into a right and left column. In the left hand column write your procedure and in the right column next to the procedure, record observations and data or measurements. Results and Discussion: You might want to include a final section that is labeled Results and Discussion. In this section, you would describe what results you got, what conclusions you have reached, ideas for continuing work, etc.

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An example of a prepared notebook follows.

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Writing Style in the Lab Notebook For certain entries in your lab notebook, such as the Introduction before each experiment, you should strive to write as logically and clearly as possible. It is also a good idea to write in the third person passive voice, to get into the habit, and so that in many cases you can copy entries from your lab notebook into your reports without the need for major revisions/rewrite. However, this is a working document. It is not expected that you write perfect prose in your notebook – it is a first draft. Just do the best you can. Also, as a working document, with many entries being written while an experiment is in progress (your observations) it is understood that many entries will be brief – but still record crucial observations. Example Notebook entry: “Added 10 mL of 1M HCl – solution turned red instantly; pcpt.↓ a few secs later→ clr soln.” When written into a lab report or journal article, this would be expanded a bit and made grammatically correct. “10 mL of 1.0 M HCl were added to the clear reaction mixture. This immediately resulted in a crimson solution, and a red precipitate formed a few seconds later, leaving a clear solution.” Adapted courtesy of Keith James

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Report Guidelines

For most experiments performed this term you will turn in a type written report (at the end of each lab you will find a summary of which sections to include in the report for that lab). The reports are due at the beginning of class the week following completion of the experiment. Below is a description of what should be included in each section. The sections are presented here in the order they should appear in your lab report. It is expected that you will complete each experiment and do the necessary calculations and analysis during the scheduled lab period each week. You may discuss the calculations and analysis with your lab mates. Your written lab report should be your own individual work!! The lab report sections should be complete but CONCISE. For most experiments this term, your report should be 1-2 pages long.

Writing Style You will write you reports using a formal scientific writing style. A lab report must be written in the third person, passive voice. Also, it must be in the past tense. It should not contain personal pronouns such as, “I”, “we” or “he” neither should it contain proper names of persons. Good: “50 mL of 1.0 M HCl were poured into a 125 mL Erlenmeyer flask”

Bad: “I poured 50 mL of hydrochloric acid into a flask.”

Also bad: “Joe Shmoe poured 50 mL of hydrochloric acid into a flask.” This is not the correct form of 3rd person. It includes Joe’s name.

Also bad: “We are going to put 50 mL of acid into the flask.” Uses future tense; also, “we”. After you write your report, there is one more thing to do before you print it and hand it in: Proofread it! Read it out loud. If is doesn’t sound right, it isn’t. Fix it. Then do it again until it is right. You will enjoy writing reports more if you take pride in what you hand in.

Abstract: This is like a condensed version of your lab report. It is a stand-alone document. Abstracts are, in fact, often published separately from the articles they describe. A library search of the literature generally involves reading abstracts. This is done with the aim to identify articles that need to be read in full, and eliminate many others whose abstract makes it clear that they are not relevant to the study at hand. So, the abstract needs to be brief, but complete. There are three questions that should be answered in any good abstract

1. What did you do? 2. How did you do it? 3. What did you find?

Even though it sequentially appears first, you should consider writing this part of the lab report after you have finished the remaining sections.

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Introduction: Here, you want to address WHY you did this experiment. Your introduction begins with a statement of the purpose of the experiment. You will do this again in the abstract, but remember the Abstract is a stand-alone document. What you say there will not count; you will find that as you write the report that you will be repeating yourself a bit. Next, provide any relevant background, to put the experiment into context. Include any key concepts, mathematical equations or chemical equations needed by the reader to understand your experiment. This means that your Introduction will often include some explanation of the theory behind the experiment. Don’t just write the equations, but provide information as to why they are relevant. You may consider writing your introduction with a central theme, such as density, types of chemical reactions….. Data: This is section is where your experimental data belong. In this section you would also include observations and descriptions of other pertinent events. This section is not where the calculations, interpretation and discussion of your results belong. (In published papers, a data section is usually not included, but, this is a class so this section will be included.) Tables

Whenever possible, data should be presented in the clearest format possible, usually in the form of a table. When you present your data in a table it is necessary to take the following into account. Number tables sequentially as they appear (Table 1, Table 2….). Be sure to refer the reader to view the tables in the text. Construct a descriptive table caption and place it above the table. Tables should include descriptive column headings, including units. Tables should not be divided across page boundaries

For a simple example, see Table 1.

Table 3: Mass of Water as Determined by a Pan Balance (+/- 0.01g): Here the volume of water delivered by a 10 mL volumetric pipet was determined utilizing the mass of water delivered and waters density (0.9980 g/ml).

Run # mass water weighed (g) Volume water (ml) 1 9.95 - 2 9.94 - 3 9.98 -

Average 9.96 9.98 Error +/- 0.02 +/- 0.02

Graphs

When graphical presentation of data is necessary, please prepare graphs using the following guidelines. Number figures sequentially as they appear (Figure 1, Figure 2….). In your writing, be sure to cite the tables in the text.

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Insert a caption below he graph that indicates what is being plotted on the y-axis vs. what is being plotted on the x-axis (always y vs. x)

Each axis should be clearly labeled, including units. Figures should not be divided across page boundaries Remove gridlines, titles and equations from the graph. If this information is pertinent, it

should be included in the caption. If the slope or intercept is necessary for other parts of the experiment, then place the

values in the caption with proper units.

For a simple example, see Figure 1.

Figure 1. A calibration curve for the absorbance at 470 nm of aqueous Allura red solutions as a function of the concentration. A best fit line was rendered resulting in a slope of 5.86 mM-1.

Results: The results section is where you should show sample calculations and report all of your results. For every type of calculation you should show one sample calculation. Each calculation should have a descriptive title, i.e. “Calculating the density of Coca-Cola”. The calculation section should be annotated. The annotation is provided to describe why each calculation is useful and relevant to the lab activity. The description should not be any longer than two or three sentences and should help you describe your results in your discussion section. Sample calculations may be written by hand attached as an appendix to your report. The results of all calculations should be summarized in a table where appropriate.

Calculating the density of Coca-Cola The volume (355 mL) and mass (394 g) of the contents of a can of coke had previously been determined above. The density is determined utilizing the relationship d=m/v (equation 1) which was explained in the introduction.

d = 394 g / 355 mL = 1.11 g/mL Discussion: In this section, you will discuss interpretations of the experimental results. This is where you get to present your thinking process. For any labs that have questions to answer, this is also where the answers get written up.

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The discussion is one of the most important parts of the lab report! It is your chance to show WHAT YOUR RESULTS ARE and that you UNDERSTAND what you did in the lab. This DOES NOT mean to include detailed procedures or that you need to re-explain your calculations in words. It DOES mean that a general description of the experiment can be useful in explaining your results and putting them in context. In this section you should also discuss error analysis. This does not necessarily mean trying to explain what went wrong. (Maybe nothing did go wrong!) It means discussing the limitations of your experiment. For example, if you are doing calorimetry in a coffee cup, and the cup feels warm to your hand, it means that some heat is escaping. Also, if you are reading a 5 degree temperature change with a thermometer that you can only read to the nearest 0.5 degree, there is a significant uncertainty in the exact magnitude of the temperature change. You could easily have a 10% error, or even more, and this needs to be taken into account. It at least needs to be mentioned, to show that you were aware of the issue. This is a limitation of the apparatus, not an error on your part. And, yes, if something did go wrong (your lab partner forgot to write down the exact molarity of your reagent), then that should go here, too, along with an explanation of how you attempted to correct for the error. (In this case, you may have had to re-do the experiment.) Adapted courtesy of Keith James.

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Example Lab Report Following is an example of a lab report prepared according to the previous report guidelines. Sample calculations can be written on a separate paper and attached to the report. Calibration of a 10 ml Volumetric Pipette a 10 Abstract: A 10 ml volumetric pipette was calibrated by determining the mass of water delivered by a pipette. A pipet was used to precisely deliver 10 mL of water. The mass of water was then converted to volume using the density of water. The volume of the pipette was determined to be 9.98 +/- 0.02 ml when the mass of water was determined on a pan balance and 9.998 +/- 0.002 ml when determined with an analytical balance. Introduction: A volumetric pipette is designed to deliver a stated volume of liquid; however, the actual amount of liquid any individual pipette delivers may vary slightly from this ideal stated volume. In order to determine the actual volume an individual pipette delivers, it will be calibrated. In this case, calibration refers to the comparison of the actual amount of liquid delivered by the pipette to the standard value of the pipette (10 ml). Because delivered volume is being calculated, another measurable quantity must be used to verify the volume delivered by the pipette. In this case, the relationship between mass and volume (density) will be used. Mass is an easily measurable quantity that can be determined with a high degree of accuracy due to the availability of electronic balances. Mass can then be converted to volume by the use of density. D (density) = m (mass)/ V (volume) The density of water at a variety of temperatures is readily available and will be used here to calibrate the volume of the pipette. Data: Diameter of beaker: 3.9 cm +/- 0.1 cm Mass of water evaporated in 60 seconds: 0.0016g +/- 0.0002g Temp of water: 20.5 ºC +/- 0.2 ºC Density of water: 0.9980 g/ml Table 1: Mass Determined by Pan Balance (+/- 0.01g)

Run # mass beaker (g) mass beaker + water (g) 1 27.88 37.83 2 27.88 37.82 3 27.88 37.86

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Table 2: Mass Determined by Analytical Balance (+/- 0.0001g) Run # mass beaker (g) mass beaker + water (g) t(transfer) t(weigh)

1 27.2349 36.5618 2:29:00 2:30:30 2 27.2348 36.7813 2:32:00 2:33:20 3 27.2335 36.8251 2:41:20 2:42:30

Results: Calculation of the volume of water: In this calculation, the average mass of water for the three

trials, shown in Table 1, as determined by the pan balance was divided by the know density of water at 20.5 ºC. The data are summarized in Table 3.

Volume = 9.96 g / 0.9980 g/mL = 9.98 mL

Table 3: Mass of Water as Determined by a Pan Balance (+/- 0.01g): Here the volume of water delivered by a 10 mL volumetric pipet was determined utilizing the mass of water delivered and waters density (0.9980 g/ml).

Run # mass water weighed (g) Volume water (ml) 1 9.95 - 2 9.94 - 3 9.98 -

Average 9.96 9.98 Error +/- 0.02 +/- 0.02 Calculation for the mass evaporated: To correct for evaporation of water in the time it takes to

measure the mass of the water delivered by the volumetric pipet, the mass of water that evaporated was estimated. The rate of evaporation of water in the 50 mL beaker in 60 seconds: 0.0016g. The data are summarized in Table 4.

Mass evaporated = rate of evaporation x time of evaporation

= (0.0016 g/ 60 s) x 90 s = 0.0024 g Calculation of the mass transferred: The mass of water initially transferred was the sum of the mass of water evaporated and the mass of water present at the time of weighing (Table 2). The data are summarized in Table 4.

Mass transferred = mass water weighed + mass (transferred) = 9.9769 + 0.0024 = 9.9793 g

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Table 4: Mass of Water as Determined by an Analytical Balance (+/-0.0001 g): ): Here the volume of water delivered by a 10 mL volumetric pipet was determined utilizing the mass of water delivered and the density water (0.9980 g/ml). A correction was added to account for the water that evaporated during the measurement.

Run # mass water weighed (g) t(evap) (s) mass (evap) (g) mass (transferred) (g) Volume (ml)1 9.9769 90 0.0024 9.9793 - 2 9.9735 80 0.0021 9.9757 - 3 9.9756 70 0.0019 9.9775 -

Average 9.978 9.998 Error +/- 0.002 +/- 0.002 Discussion: The mass of water delivered by a 10 ml volumetric pipette was determined on both a pan balance and an analytical balance (Tables 1 and 2 respectively). The mass of water was then converted to volume using the density of water. In the case of the analytical balance, the rate of evaporation of water (which is a systematic error) was taken into consideration. In this case the mass of water that evaporated from the time the water was delivered to the beaker to the time of weighing was added to the weighed mass of water delivered by the pipette. This correction was not necessary when the pan balance was used since the accuracy of the pan balance is +/- 0.01 g and the evaporation rate of water under experimental conditions was found to be 2.7 x 10-5 g/s. The use of the analytical balance increased both the precision and the accuracy of the calculated volume of the pipette (9.98 +/- 0.02 ml with the pan balance and 9.998 +/- 0.002 ml with the analytical balance, see tables 3 and 4). The improvement in the results can easily be seen by the percent error which was calculated to be 0.2 % with the pan balance and 0.02 % with the analytical balance. The largest source of error in this experiment most likely came from the difficulty in accurately filling the pipette to the mark with water which introduced random error into the experiment.

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General Chemistry Lab Report Checklist

General _____ Have you listed your name, partner's name, a descriptive lab title and date? _____ Did you use spellchecker? _____ Is your report written in passive third person voice (you did not use the words I,

we, they, etc.) _____ Is proper tense is maintained within sections? _____ Have you correctly written your chemical formula and names correctly? _____ Were correct subscripts, superscripts, and symbols are used? _____ Did you separate the numbers from their units (0.25 mL was added…. not 0.25mL

was added)? _____ Did you check significant figures? _____ Do your numbers include leading zeros (0.25 mL was added…. not .25 mL was

added)? _____ Did you make sure that you did not start a sentence with a number? _____ Are your references cited in one official style? _____ Have you made a citations whenever ideas from outside? _____ All subjects and verbs are in agreement? _____ Did you make sure that there are no run-on sentences or fragments?

Abstract

The abstract is a condensed summary of the report's findings. Abstracts are often written last. They should be clear, concise, and self-contained and, in the context of this lab, approximately three sentences long.

_____What did you do? (Identify the rationale behind the investigation)? _____How did you do it (summarize the procedure, without using specific steps)? _____Present the important findings numerically including error statistics?

Introduction

The introduction will provide the reader information on what you are doing why you did it and critical background information necessary in understanding the methods and results of your experiment.

_____Did you include a statement of purpose? _____Is there sufficient background so that the reader can understand what you did? _____Are necessary equations, chemical or mathematical, included?

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Data This section should give only the data and observations from the lab, without results

_____Are your data tables properly formatted? _____Are your calculations, either attached as an appendix, or typed neatly into the data

section? _____Are your figures and tables numbered sequentially and referred to in the text. Table

captions above and figure captions below. Tables and figures are not broken over multiple pages

_____Are the axes on your graphs formatted properly with labels? _____Are all graphs and tables accompanied by a written description relating the same

information to the reader?

Results: We will be treating this section as a calculational section. This is where you will be showing all calculations along with a written description as to how the calculations were carried out and what the result of the calculation is and how it relates to the lab. Your readers must easily find your results in order to evaluate and interpret them.

_____Are calculations accompanied by text explaining the both the method of calculation and results of the calculation?

_____Units? Significant Figures? _____Is a straight forward presentation of the results of your experiment included in

either a table or in text? _____Can your key results be understood by a reader without reliance on figures and

tables?

Discussion: In this section, you will discuss interpretations of the experimental results. It will be necessary to describe your results, cite tables or figures. It should include a general description of the experiment to put the results into context.

_____Can your key results and discussion be understood by a reader without reliance on figures and tables?

_____Are key results highlighted and carefully explained? _____Did you make logical deductions based on the results (usually questions are given

in the lab manual to help this)? _____Have you discussed sources of error or ambiguities in the data? _____Did you confirm all relationships that were stated in purpose or abstract? _____Do your conclusions clearly contribute to the understanding of the overall

problem?

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Chemistry 228

Pre-Lab: Enthalpy of Neutralization of Phosphoric Acid

Part A Answer the following questions in your lab notebook (be sure to show work for any calculations):

1. A neutralization reaction was carried out in a calorimeter. The temperature of the solution rose from 20.0 °C to 25.6 °C. Is this reaction endothermic or exothermic?

2. A neutralization reaction was carried out in a calorimeter. The change in temperature (∆T)

of the solution was 5.6 °C and the mass of the solution was 100.0 g. Calculate the amount of heat energy gained by the solution (qsol). Use 4.18 J/(g•°C) as the specific heat, Cs, of the solution.

3. What is the value of qreaction for the neutralization reaction described in number 2?

4. How many moles of phosphoric acid are contained in 50.0 mL of 0.60 M H3PO4?

5. What is the value of Hreaction (in kJ/mol phosphoric acid) if 50.0 mL of 0.60 M H3PO4 was

used in the reaction described in number 2? Part B Prepare your lab notebook for the lab. This includes stating the purpose of the experiment, summarizing the procedure in a bulleted format (be sure to include space for observations) and preparing any tables necessary for data collection.

At the start of your lab, remove the copies of the pages where you completed the above work from your lab notebook and turn them into your TA.

25

The Enthalpy of Neutralization of Phosphoric Acid

OBJECTIVES

In this experiment, you will

Measure the temperature change of the reaction between solutions of sodium hydroxide and phosphoric acid.

Calculate the enthalpy, ΔH, of neutralization of phosphoric acid. Compare your calculated enthalpy of neutralization with the accepted value. Calculate the enthalpy, ΔH, of neutralization per ionizable hydrogen for phosphoric acid.

INTRODUCTION

A great deal can be learned by conducting an acid-base reaction as a titration. In addition, acid-base reactions can be observed and measured thermodynamically. In this case, the reaction is carried out in a calorimeter. If the temperature of the reaction is measured precisely, the enthalpy of neutralization of an acid by a base (or vice versa) can be determined. In this experiment, you will react phosphoric acid with sodium hydroxide.

You will use a Styrofoam cup nested in a beaker as a calorimeter, as shown in Figure 1. For purposes of this experiment, you may assume that the heat loss to the calorimeter and the surrounding air is negligible. Phosphoric acid will be the limiting reactant in this experiment, and you will accordingly be determining the enthalpy, ΔH, of neutralization of the acid. Selecting a limiting reactant helps ensure that the temperature measurements and subsequent calculations are as precise as possible.

Pages 246-248 and 257-258 in your text will provide background information.

Figure 1

MATERIALS NEEDED

Vernier computer interface Temperature Probe PROCEDURE

1. Obtain and wear goggles. It is best to conduct this experiment in a well-ventilated room.

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2. Connect a Temperature Probe to Channel 1 of the Vernier computer interface.

3. Start the Logger Pro program on your computer. Open the file “Lab 1 Phosphoric” from the Chemistry 228 folder.

4. Nest a Styrofoam cup in a 250 mL beaker as shown in Figure 1. Measure out 50.0 mL of 0.60 M H3PO4 solution into the foam cup. CAUTION: Handle the phosphoric acid with care. It can cause painful burns if it comes in contact with the skin.

5. Use a utility clamp to suspend the Temperature Probe from a ring stand (see Figure 1). Lower the Temperature Probe into the phosphoric acid solution.

6. Measure out 50.0 mL of 1.85 M NaOH solution in a graduated cylinder and transfer it to a 250 mL beaker. CAUTION: Sodium hydroxide solution is caustic. Avoid spilling it on your skin or clothing.

7. Conduct the experiment.

a. Click to begin the data collection and obtain the initial temperature of the H3PO4 solution.

b. After you have recorded three or four readings at the same temperature, add the 50.0 mL of NaOH solution to the Styrofoam cup all at once. Use a glass stirring rod to stir the reaction mixture gently and thoroughly.

c. Data will be collected for 10 minutes. You may terminate the trial early by clicking , if the temperature readings are no longer changing.

d. Click the Statistics button, . The minimum and maximum temperatures are listed in the statistics box on the graph. If the minimum temperature is not a suitable initial temperature, examine the graph and determine the initial temperature.

e. Record the initial and maximum temperatures for Trial 1. f. Close the Statistics box by clicking the X in the corner of the box.

8. Rinse and dry the Temperature Probe, Styrofoam cup, and stirring rod. Dispose of the solution as directed.

9. Repeat Steps 4–8 to conduct a second trial. If directed, conduct a third trial. Print a copy of the graph of the second trial to include with your data and analysis.

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DATA TABLE

Trial 1 Trial 2 Trial 3

Maximum temperature (°C)

Initial temperature (°C)

Temperature change (∆T)

DATA ANALYSIS

6. Write the balanced equation for the reaction of phosphoric acid and sodium hydroxide. 7. Use the equation below to calculate the amount of heat energy gained by the solution (qsol).

In determining the mass, m, of the solution use 1.11 g/mL for the density (be sure to use the total volume of the solution after the acid and base are mixed). The change in temperature (∆T) is a directional change where ∆T = Tf –Ti. Use 4.18 J/(g•°C) as the specific heat, Cs, of the solution.

qsol = Cs m ∆T

8. The heat calculated above represents the heat gained by the solution (the solution being predominantly water). Since we are interested in the heat of neutralization of phosphoric acid we need the heat transfer associated with the reaction (qrxn). If the solution gained heat, the reaction must have given off heat. This relationship can be expressed by the following equation:

qsol = -qrxn

9. Determine the number of moles of phosphoric acid used in the reaction. Use the moles of phosphoric acid along with qrxn to determine the enthalpy change, ∆H, for the reaction in terms of kJ/mol of phosphoric acid. This is your experimental value of ∆H.

∆H = qrxn/moles H3PO4

10. The accepted value for the ∆H of neutralization for phosphoric acid is -156.44 kJ/mol.

Calculate the percent error in your experimental value.

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The Enthalpy of Neutralization of Phosphoric Acid Lab Report Your report for this lab should include the following sections: Abstract:

Your abstract should be written individually Introduction: Include why you did this experiment, relevant background, and general equations. Data:

Include your data table

Results: Report your calculated value of ∆H of neutralization for phosphoric acid and the value of ∆H of neutralization per ionizable hydrogen in phosphoric acid. Include a results table with results from each trial and an average value for the ∆H of neutralization for phosphoric acid. Report the percent error for the ∆H of neutralization for phosphoric acid. Be sure to attach hand written sample calculations to the back of your report.

Discussion: Discuss the experiment and any possible sources of error

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Chemistry 228

Pre-Lab: Hess’s Law

Part A Answer the following questions in your lab notebook (be sure to show work for any calculations): 1. What is the formula that relates the temperature change observed in a substance with the energy released or absorbed? 2. When you measure a temperature rise during a chemical reaction, is the reaction endothermic or exothermic? 3. The enthalpy of the reaction for the reaction of calcium oxide with hydrochloric acid is exothermic. Will the reverse reaction have a positive or negative H? 4. Hess’s Law allows us to combine reactions to determine the heat of reaction for a net reaction that has not been measured. For the reactions described in the lab, the second reaction is difficult to measure as written. You will measure the heat of reaction for the reverse reaction. How will you use the measurement in the Hess’s Law calculation? Part B Prepare your lab notebook for the lab. This includes stating the purpose of the experiment, summarizing the procedure in a bulleted format (be sure to include space for observations) and preparing any tables necessary for data collection.


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Enthalpy of Reaction and Hess's Law

Introduction In this experiment you will be finding the enthalpy of formation for MgO(s) using an indirect method. Remember, according to Hess's Law (see your textbook for more details), if two or more reactions can be added to give a net reaction, H° for the net reaction is simply the sum of the H°''s for the reactions which are added. Consider the following three reactions: 1) Mg(s) + 2 H+(aq) ---> Mg2+(aq) + H2(g) H°1 2) Mg2+(aq) + H2O(l) ---> MgO(s) + 2 H+(aq) H°2 3) H2(g) + 1/2 O2(g) ---> H2O(l) H°3 4) Mg(s) + 1/2 O2(g) ---> MgO(s) H°4 You will determine the heat of reaction for reactions 1 and 2 experimentally, then use the known value of the enthalpy of formation of water (H°3 = -285.9 kJ) to calculate H°4 which is the enthalpy of formation of MgO. Be aware that equation (2) is the reverse of the reaction you actually run and measure. (Note: the enthalpy of formation of MgO cannot easily be measured.) Experimental Procedure Obtain a coffee cup calorimeter from the stockroom. Make sure the cup is clean and dry. a) Using a weighing boat, weigh out a sample containing between 0.45 - 0.55 grams

of magnesium. Put 50. mL of 1.0 M HCl into the calorimeter and measure the temperature until it stabilizes. Record this reading as your initial temperature. Then add the Mg and replace lid. Stir vigorously as the metal dissolves (use a stir rod, not the thermometer), and record the temperature every 30 seconds until it is approximately constant for two minutes. Record your data in your lab notebook. Use a format that allows you to readily identify which experiment you are recording.

Conduct another trial as above. Make sure the calorimeter is clean and mostly dry

before repeating the experiment.

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b) Repeat the above procedure, this time replacing Mg with MgO. (Use a clean, dry calorimeter.) You should use a molar equivalent of MgO (24.3 g Mg is the molar equivalent of 40.3 g MgO, why?, your measurement should be within 5%) Be certain all the MgO dissolves, this will require vigorous stirring!!

Conduct another trial as above. Calculations To relate heats of reactions (in energy units of Joules) with temperature differences we use:

q = m x S x T

For the reactions above, it is a good approximation to take specific heat of the solution to be the specific heat of water, S = 4.184 J/g-°C. For mass, because you are using the specific heat of pure water, use the mass of the water only, not the combined mass of water and solute. Calculate q for reaction 1 and 2. Report H°rxn for reaction 1 and 2, be certain to use units of kJ/mol. Calculate H°4

Caution: Wear your goggles at all times. HCl is a strong acid. Hydrogen gas is flammable. Do not use any open flames in the lab.

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Hess’s Law Lab Report: Your report for this lab should include the following sections: Abstract:

Your abstract should be written individually

Introduction: Include why you did this experiment, relevant background, and general equations. Data:

Prepare a data table that includes the initial and final temperatures for each trial Report the mass of Mg and MgO used in each trial

Results: Prepare a results table showing the calculated Hrxn for each trial and averages for reactions one and two and the value of H for reaction four Be sure to attach hand written sample calculations to the back of your report


As part of your discussion, answer the following questions: 1. For an exothermic reaction, does the temperature observed rise or fall? 2. For an exothermic reaction, is H° positive or negative? 3. Is reaction 1 endothermic or exothermic? reaction 2? 4. In this lab, you measure the quantity, q. How is this different

from H°rxn, (remember the definitions and units)? Answer the following question and attach it to your report: Further Analysis An alloy (a metal mixture) containing magnesium and another metal, that does not react with hydrochloric acid, needs to analyzed. You are asked to determine the percent magnesium in the alloy. a) Describe the procedure you would use to determine the percent magnesium in the alloy, and b) if the sample were 30% magnesium calculate the heat evolved if a 5 gram sample were analyzed in that manner.

33

Deriving the Gas Laws Using Computer Simulations

Introduction According to the kinetic molecular theory, gases are in constant and random motion with enough kinetic energy such that they rarely interact with one another. When gas particles collide with the walls of a container, they rebound with no apparent loss of energy. These characteristics describe an "Ideal Gas." Experimental evidence suggests that many common gases making up air behave in this manner when studied at temperatures well above their boiling points. We are constantly being exposed to the behavior of gases. Each time we pump up a tire, blow up a balloon, use a spray can, or experience the cooling of gases as they escape from a gas storage container, we are reminded of how gases behave with changes in temperature (T), volume (V), pressure (P), or number of particles (n). The behavior of gases has been scientifically investigated starting with Robert Boyle's work in 1662, followed by Jacques Charles' (1787) and Joseph Gay-Lussac's work (1802). Together these studies led to the so called "Gas Laws" which relate volume (V), pressure (P), temperature (T) and numbers of particles of gas (n). In a scientific manner, one can derive the mathematical relationships that exist between these variables by holding two of the variables constant, changing one and monitoring the effect on the fourth variable. To derive the relationships, you will be using an interactive research-based simulation produced by the PhET project at the University of Colorado. PROCEDURE 1: Pressure Volume Relationship

1. Go to the Physics Education Technology from the University of Colorado at: http://phet.colorado.edu/new/simulations/sims.php?sim=Gas_Properties

2. Click the RUN NOW button under the Gas Properties Simulation window (highlighted in green).

3. Play around with the simulator and see what sorts of tools are available to you to analyze

the behaviors of gases. Qualitatively get a feel for the relationships that exist between the four variables that describe gases: P, V, n and T. If you ever get to a point that you need to reset the simulator, you can always hit the reset button at the bottom right of the screen.

4. If you have not already done so, on the lower right side of the screen, click on the

RESET button.

5. On the right side of the screen, click on the MEASURMENT TOOLS button. Next, click on the RULER option to activate the ruler.

34

CLICK HERE

6. In the upper right hand corner, click on the TEMPERATURE button under the

Constant Parameter heading. This will hold temperature constant while allowing you to observe the relationship between pressure and volume.

CLICK HERE

7. Using the mouse and the right button, drag the ruler into a position that will allow you to measure the length of the container.

8. Using the mouse and the right button, grab hold of the man pushing against the container

and expand the length of the container so that it measures 9.0 cm. Record this as your initial length (the height of the box will remain 5.0 cm and the width of the box will remain 5.0 cm)

35

GRAB AND DRAG

9. Using the mouse and the right button, grab hold of the pump handle and inject one cycles worth of gas into the chamber by pulling the handle up then pushing it back down.

MOVE UP THEN DOWN

10. Once the pressure has somewhat stabilized, record your pressure value for the chamber length of 9.0 cm. This will represent your initial pressure in atmospheres.

36

PRESSURE (ATM)

11. Using the mouse and right button, grab hold of the man pushing on the container and decrease the length of the container to approximately 8.0 cm. Once the pressure has stabilized (again, this may take a short period of time to happen), record the new pressure for a length of 8.0 cm.

PUSH IN

12. Repeat step 9 for approximate lengths of 7.0 cm, 6.0 cm, 5.0 cm, 4.0 cm, 3.0 cm, and 2.0 cm (you will probably not get to exactly 2 cm). For each trial, record the length value and resulting pressure value in a properly labeled data table.

37

13. Record any qualitative observations on the behavior of the gas molecules as the volume decreases.

14. Click the RESET button to remove all the gas particles from the chamber before moving

on to the next section. PROCEDURE 2: Volume Temperature Relationship Devise an experiment using the simulator in which you can elucidate the relationship between Temperature and the Volume of a gas. Collect and record your data over a wide range of temperatures, 0-600 K, in a properly labeled table Procedure 3: Temperature Pressure Relationship Devise an experiment using the simulator in which you can elucidate the relationship between Temperature and the Pressure of a gas. Collect and record your data over a wide range of temperatures, 0-600 K, in a properly labeled table Procedure 4: Pressure Quantity Relationship Devise an experiment using the simulator in which you can elucidate the relationship between Quantity and the Pressure of a gas. Collect and record your data over a wide range of number of molecules in a properly labeled table Analysis: For this lab, you will need to submit neat labeled data tables for each procedure. You must also submit a graphical representation for each relationship. Be sure to label each axis and include a title for each graph (Please see the information on pages 15 and 16 of this lab manual). I suggest that you utilize Microsoft Excel or some other comparable spreadsheet software to produce your tables and graphs. Along with the graphs and tables for each procedure, answer completely the questions below that correlate with each section. Analysis: Procedure 1: Pressure Volume Relationship

1. Graphically represent the Pressure (atm) Volume (cm3) relationship with volume on the x-axis.

2. Graphically represent the Pressure (atm) and Inverse Volume, 1/V (cm-3) relationship with 1/V on the x axis

3. Identify the mathematical relationship that exists between pressure and volume, when temperature and quantity are held constant, as being directly proportional or inversely proportional. Explain your answer and write an equation that relates pressure and volume to a constant.

4. Why were you asked to graph pressure and the inverse of volume? 5. Calculate the slope of the line for your pressure vs. 1/volume graph. What does this

number represent? Would you expect it to be the same for all gases? Explain your answer.

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Analysis Questions: Procedure 2: Volume Temperature Relationship 6. Graphically represent the Temperature (K) Volume (cm3) relationship. 7. Identify the mathematical relationship that exists between volume and temperature, when

pressure and quantity are held constant, as being directly proportional or inversely proportional. Explain your answer and write an equation that relates volume and temperature to a constant.

8. Calculate the slope of the line for your temperature vs. volume graph. What does this number represent? Would you expect it to be the same for all gases? Explain your answer.

Analysis Questions: Procedure 3: Temperature Pressure Relationship

9. Graphically represent the Temperature (K) Pressure (atm) relationship. Make sure the axis that represents temperature includes a range from 0 K to 600 K.

10. Identify the mathematical relationship that exists between pressure and temperature, when volume and quantity are held constant, as being directly proportional or inversely proportional. Explain your answer and write an equation that relates pressure and temperature to a constant.

11. Calculate the slope of the line for your temperature vs. pressure graph. What does this number represent? Would you expect it to be the same for all gases? Explain your answer.

12. Explain the effects of temperature on molecular motion. Using this explanation, explain why both pressure and volume can decrease with decreasing temperature.

13. Absolute zero is theorized to be the temperature that all molecular motion stops. Based on this, what would you predict to be the pressure and volume of a gas sample whose temperature is decreased to absolute zero? Explain.

Analysis Questions: Procedure 4: Pressure Quantity Relationship

14. Graphically represent the Quantity (number of molecules) Pressure (atm) relationship 15. Describe the impact of increasing the number of molecules (or moles) of a gas on the

pressure of a gas sample. Would you expect this trend to be the same for all gases? Explain your answer.

16. Based on your previous observations, predict the impact of changing the number of moles of a gas sample on the volume of the gas sample (if pressure and temperature are held constant) and on the temperature of a gas sample (if pressure and volume are held constant). Explain your answer.

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Chemistry 228

Pre-Lab: Decomposition of Hydrogen Peroxide


1. What is Dalton’s law of partial pressure? 2. A mixture of three gasses (A, B and C) has a total pressure of 849 torr and the partial

pressure of A is 57 torr and the partial pressure of B is 573 torr. What is the partial pressure of C?

3. A gas has a volume of 94 mL, a pressure of 743 torr and a temperature of 20 °C.

Calculate the number of moles of gas present. Assume ideal behavior of the gas.

4. If 0.00946 moles of O2 gas is collected from the decomposition of hydrogen peroxide, how many moles of hydrogen peroxide were reacted?

Part B Prepare your lab notebook for the lab. This includes stating the purpose of the experiment, summarizing the procedure in a bulleted format (be sure to include space for observations) and preparing any tables necessary for data collection.


40

Decomposition of Hydrogen Peroxide OBJECTIVES Decompose hydrogen peroxide using KI as a catalyst Measure the volume of oxygen gas generated through the decomposition reaction Illustrate Dalton’s Law of partial pressure Determine the number of moles of oxygen gas produced using the ideal gas law

Determine the percent hydrogen peroxide in an aqueous solution

INTRODUCTION Hydrogen peroxide spontaneously decomposes to form oxygen gas according to the following equation: 2 H2O2 (aq) → 2 H2O (l) + O2 (g) This process usually occurs very slowly. Many different compounds or ions are capable of acting as catalysts increasing the rate of the reaction. Here, potassium iodide (KI) will be used as a catalyst to make the reaction produce products rapidly enough to study the reaction in the lab. The apparatus we will use to collect oxygen gas in this experiment is shown in figure 1. Hydrogen peroxide will be placed in the Erlenmyer flask. The catalyst (KI) is located in the syringe and can be added to the Erlenmyer flask to initiate the reaction. As the reaction proceeds oxygen gas will be produced in the Erlenmyer flask and travel through the tubing. The gas will be collected in the graduated cylinder. The graduated cylinder is initially filled with water. As the gas enters the cylinder it displaces water allowing the volume of the gas to be measured.

Figure 1

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MATERIALS NEEDED 125 mL Erlenmyer flask

100 mL graduated cylinder Rubber stopper with adaptors, tubing and syringe PROCEEDURE

1. Place an 125 mL Erlenmyer flask on a balance and tare the scale. Add approximately 5 g

of hydrogen peroxide solution into the Erlenmyer flask. Record the actual mass used. Obtain a ring stand and clamp the flask as shown in figure 1. Place the rubber stopper tightly in the flask (this should be air tight).

2. Place approximately 400 mL of water in an 800 mL beaker. 3. Completely fill a 100 mL graduated cylinder with water. Cover the cylinder with

parafilm and invert the cylinder in the 800 mL beaker. Carefully clamp the cylinder in place such that the opening of the cylinder is below the surface of the water in the beaker. Remove the parafilm and carefully place the end of the tubing just inside the graduated cylinder as shown in figure 1. The graduated cylinder should be completely filled with water. If there is a small amount of air present in the cylinder record the volume. If there is more than 10 mL of air in the cylinder, you will need to redo the setup.

4. In a small beaker obtain a small amount (approximately 10 mL) of 0.5 M KI. Draw up 3 mL of the KI solution into the syringe. Attach the syringe to the adaptor in the rubber stopper.

5. Initiate the reaction by depressing the stopper on the syringe and adding the KI to the hydrogen peroxide.

6. Allow the reaction to proceed until no further production of oxygen gas is observed (around 10 to 15 minutes).

7. Measure and record the temperature of the water. 8. Record the final level of the water in the graduated cylinder. Be sure to record your

measurement to 2 decimal places. 9. Repeat the above procedure two more times for a total of three trials. At least two of

your trials should agree well with one another. DATA ANALYSIS

1. Determine the pressure of the oxygen gas:

Because the oxygen gas was collected over water some of the gas collected is water vapor. The gas collected is therefore a mixture of both oxygen and water. The total pressure of the gas is the sum of the pressures exerted by the oxygen gas and water vapor. To determine the pressure of oxygen gas we must apply Dalton’s law of partial pressure.

PTot = pO2 + pH2O

In other words, you can find the pressure of oxygen gas by subtracting the partial pressure of water at the temperature of the water (also known as the vapor pressure of water) from the total pressure (or atmospheric pressure). Your TA will provide the

42

current barometric pressure. A table of the vapor pressure of water at various temperatures follows.

Table 1: Vapor pressure of water at various temperatures

Temperature Vapor

Pressure TemperatureVapor

Pressure TemperatureVapor

Pressure

(oC) (torr) (oC) (torr) (oC) (torr) 15 12.8 21 18.6 27 26.7 16 13.6 22 19.8 28 28.3 17 14.5 23 21.1 29 30 18 15.5 24 22.4 30 31.8 19 16.5 25 23.8 31 33.7 20 17.5 26 25.2 32 35.7

2. Determine the volume of the oxygen gas:

When the reaction was initiated, 3 mL of KI solution was added. This volume needs to be subtracted from the volume of gas collected. If your initial volume of gas was not zero, this must also be taken into consideration. VO2

= Vfinal – Vinitial – 3 mL

3. Calculate the number of moles of oxygen gas generated:

Now that the pressure, volume and temperature of the gas are known, the moles of gas can be calculated using the ideal gas law. The temperature of the gas will be considered to be the same temperature as the water temperature measured during the experiment. PV = nRT

4. Calculate the amount of H2O2: Using the balanced equation, calculate the number of moles of H2O2 present in the initial solution. Calculate the molar mass of H2O2 and determine the grams of H2O2 present in the initial solution.

5. Calculate the mass percent of hydrogen peroxide: Using the mass of H2O2 calculated above and the initial mass of the H2O2 solution, calculate the mass percent H2O2 in the initial solution.

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Hydrogen Peroxide Lab Report: Your report for this lab should include the following sections: Abstract:



Include a data table with data from all 3 trials

Results: Include a results table with the mass percent of hydrogen peroxide from each trial Be sure to attach hand written sample calculations to the back of your report


44

Chemistry 228

Pre-Lab: Vapor Pressure and Heat of Vaporization


1. When using the equation P1/T1 = P2/T2 to relate temperature and pressure of a gas, what must be held constant?

2. A sample of gas is held in a capped flask. At 25 °C the pressure is 693 mmHg. What is

the pressure of the gas at 37 °C?

3. If the heat of vaporization of water is 40.7 kJ/mol, how much energy is required to

vaporize 5.0 g of liquid water at 100 °C? 4. Would you expect most of the components in a perfume to have a low or high vapor

pressure? Explain. Part B Prepare your lab notebook for the lab. This includes stating the purpose of the experiment, summarizing the procedure in a bulleted format (be sure to include space for observations) and preparing any tables necessary for data collection.


45

Vapor Pressure and Heat of Vaporization

When a volatile liquid is placed in a container, and the container is sealed tightly, a portion of the liquid will evaporate. The newly formed gas molecules exert pressure in the container, while some of the gas condenses back into the liquid state. If the temperature inside the container is held constant, then at some point a physical equilibrium will be reached. At this equilibrium, the rate of condensation is equal to the rate of evaporation. The pressure at equilibrium is called vapor pressure, and will remain constant as long as the temperature in the container does not change.

In mathematical terms, the relationship between the vapor pressure of a liquid and temperature is described in the Clausius-Clayperon equation,

CTR

HP vap

1

ln

where ln P is the natural logarithm of the vapor pressure, ΔHvap is the heat of vaporization, R is the universal gas constant (8.31 J/mol•K), T is the temperature (in Kelvin) and C is a constant not related to heat capacity. Thus, the Clausius-Clayperon equation not only describes how vapor pressure is affected by temperature, but it relates these factors to the heat of vaporization of a liquid. ΔHvap is the amount of energy required to cause the vaporization of one mole of liquid at constant pressure.

In this experiment, you will introduce a specific volume of a volatile liquid into a closed vessel, and measure the pressure in the vessel at several different temperatures. By analyzing your measurements, you will be able to calculate the ΔHvap of the liquid.

OBJECTIVES


Measure the pressure inside a sealed vessel containing a volatile liquid over a range of temperatures.

Determine the relationship between pressure and temperature of the volatile liquid. Calculate the heat of vaporization of the liquid.

Figure 1

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MATERIALS

Vernier Gas Pressure Sensor Temperature Probe rubber stopper assembly plastic tubing with two connectors

PROCEDURE

1. Obtain and wear goggles. CAUTION: The alcohol used in this experiment is flammable and poisonous. Avoid inhaling the vapors. Avoid contact with your skin or clothing. Be sure that there are no open flames in the room during the experiment. Notify your teacher immediately if an accident occurs.

2. Use a hot plate to heat ~200 mL of water in a 400 mL beaker.

3. Prepare a room temperature water bath in an 800 mL beaker. The bath should be deep enough to completely cover the gas level in the 125 mL Erlenmeyer flask.

4. Connect a Gas Pressure Sensor to Channel 1 of the Vernier computer interface. Connect a Temperature Probe to Channel 2 of the interface.

5. Start the Logger Pro program on your computer. Open the file “Lab 4 Vapor Pressure” from the Chemistry 228 folder.

6. Use the clear tubing to connect the white rubber stopper to the Gas Pressure Sensor. (About one-half turn of the fittings will secure the tubing tightly.) Twist the white stopper snugly into the neck of the Erlenmeyer flask to avoid losing any of the gas that will be produced as the liquid evaporates (see Figure 1). Important: Open the valve on the white stopper.

7. Your first measurement will be of the pressure of the air in the flask and the room temperature. Place the Temperature Probe near the flask. When the pressure and temperature readings stabilize, record these values.

8. Condition the Erlenmeyer flask and the sensors to the water bath.

a. Place the Temperature Probe in the room temperature water bath. b. Place the Erlenmeyer flask in the water bath. Hold the flask down into the water bath to

the bottom of the white stopper. c. After 30 seconds, close the valve on the white stopper.

9. Obtain a small amount of ethanol. Draw 3 mL of ethanol into the 20 mL syringe that is part of the Gas Pressure Sensor accessories. Thread the syringe onto the valve on the white stopper (see Figure 1).

10. Add ethanol to the flask.

a. Open the valve below the syringe containing the 3 mL of ethanol. b. Push down on the plunger of the syringe to inject the ethanol. c. Quickly pull the plunger back to the 3-mL mark. Close the valve below the syringe. d. Carefully remove the syringe from the stopper so that the stopper is not moved.

11. Gently rotate the flask in the water bath for a few seconds, using a motion similar to slowly stirring a cup of coffee or tea, to accelerate the evaporation of the ethanol.

47

12. Monitor and collect temperature and pressure data.

a. Click to begin data collection. b. Hold the flask steady once again. c. Monitor the pressure and temperature readings. d. When the readings stabilize, click . e. Record these values in your notebook.

13. Add a small amount of hot water, from the beaker on the hot plate, to warm the water bath by 3–5°C. Use a spoon or a dipper to transfer the hot water. Stir the water bath slowly with the Temperature Probe. Monitor the pressure and temperature readings. When the readings stabilize, click . Record these values in your notebook.

14. Repeat Step 13 until you have completed five total trials. Add enough hot water for each trial so that the temperature of the water bath increases by 3-5°C, but do not warm the water bath beyond 40°C because the pressure increase may pop the stopper out of the flask. If you must remove some of the water in the bath, do it carefully so as not to disturb the flask.

15. After you have recorded the fifth set of readings, open the valve to release the pressure in the flask. Remove the flask from the water bath and take the stopper off the flask. Dispose of the ethanol as directed.

16. Click to end the data collection. Record the pressure readings, as Ptotal, and the temperature readings in your data table.

17. Do not exit the Logger Pro program until you have completed 1–4 of the Data Analysis section.

DATA TABLE

Initial Trial 1 Trial 2 Trial 3 Trial 4 Trial 5

Ptotal (kPa)

Pair (kPa)

Pvap (kPa)

Temperature (°C)

48

DATA ANALYSIS

1. The Pair for Trials 2-5 must be calculated because the temperatures were increased. As you

warmed the flask, the air in the flask exerted pressure that you must calculate. Use the gas law relationship shown below to complete the calculations. Remember that all gas law calculations require Kelvin temperature. Use the Pair from Trial 1 as P1 and the Kelvin temperature of Trial 1 as T1.

2

2

1

1

T

P

T

P

3. Calculate and record the Pvap for each trial by subtracting Pair from Ptotal.

4. Prepare and print a graph of Pvap (y-axis) vs. Celsius temperature (x-axis).

a. Disconnect your Gas Pressure Sensor and Temperature Probe from the interface. b. Choose New from the File menu. An empty graph and table will be created in Logger Pro. c. Double-click on the x-axis heading in the table, enter a name and unit, then enter the five

values for temperature (°C) from your data table above. d. Double-click on the y-axis heading in the table, enter a name and unit, then enter the five

values for vapor pressure from your data table above. e. Does the plot follow the expected trend of the effect of temperature on vapor pressure?

Explain.

5. In order to determine the heat of vaporization, ΔHvap, you will first need to plot the natural log of Pvap vs. the reciprocal of absolute temperature.

a. Choose New Calculated Column from the Data menu. b. Create a column ln vapor pressure. c. Create a second column, reciprocal of absolute temperature, 1/(Temperature (°C) + 273). d. On the displayed graph, click on the respective axes, and then select ln vapor pressure to

plot on the y-axis, and reciprocal of absolute temperature to plot on the x-axis. Autoscale the graph, if necessary.

e. Calculate the linear regression (best-fit line) equation for this graph. Calculate ΔHvap from the slope of the linear regression.

f. Prepare and print a second graph.

6. The accepted value of the ΔHvap of ethanol is 42.32 kJ/mol. Compare your experimentally determined value of ΔHvap with the accepted value.

49

Vapor Pressure and Heat of Vaporization Lab Report: Your report for this lab should include the following sections: Abstract:



Include your data table

Results: Include a copy of your graph of Pvap vs. T (°C) Include a copy of your graph of ln Pvap vs. 1/T (K) Report your value for Hvap of ethanol Report the percent error in your calculated value of Hvap for ethanol Be sure to attach hand written sample calculations to the back of your report


50

Chemistry 228

Pre-Lab: Freezing Point Depression

Part A Answer the following questions in your lab notebook (be sure to show work for any calculations): 1. What is a colligative property? 2. Give the equation for freezing point depression and indicate the units for each term in the expression. 3. A measurement of the freezing temperature of a solution allows you to calculate the concentration of the solution. What else do you need to measure to determine the molar mass of the solid added to the solvent? 4. A student adds 1.504 g of a solid to 25.0 mL of water. The freezing temperature is measured to be -1.20 °C. What is the molality of the solution? 5. What is the molar mass of the solid above? Part B Prepare your lab notebook for the lab. This includes stating the purpose of the experiment, summarizing the procedure in a bulleted format (be sure to include space for observations) and preparing any tables necessary for data collection.


51

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 (8.28°C-kg/mol for t-butanol in this experiment1 ), 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, t-butanol, C4H10O. You will then add a known mass of aspirin, to a known mass of t-butanol, and determine the lowering of the freezing temperature of the solution. By measuring the freezing point depression, T, and the mass of aspirin, you can use the formula above to find the molar mass of the aspirin solute, in g/mol.

OBJECTIVES


Determine the freezing temperature of pure t-butanol. Determine the freezing temperature of a solution of aspirin and t-butanol. Examine the freezing curves for each. Calculate the experimental molar mass of aspirin. Compare it to the accepted molar mass for aspirin.

Figure 1

MATERIALS

Temperature Probe Copper stirrer

1 “The Computer-Based Laboratory,” Journal of Chemical Education: Software, 1988, Vol. 1A, No. 2, p. 73.

52

PROCEDURE

1. Obtain and wear goggles.

2. Connect the Temperature Probe to the computer interface. Prepare the computer for data collection by opening the file “15 Freezing Pt Depression” from the Chemistry with Computers folder.

Part I Freezing Temperature of Pure T-butanol

3. Add about 175 mL of tap water 250 mL beaker. Using a hot plate warm the water to a temperature of about 35 °C. The t-butanol used in this experiment is flammable. Do not use Bunsen burners during this lab.

4. Weigh a CLEAN, DRY test tube. It can be propped in a plastic 250 mL beaker to facilitate measuring; this is useful when the tube is not empty. Add 3 mL of t-butanol to your test tube using a dry Pasteur pipette. Do this carefully so that you do not get any t-butanol on the upper portion of the test tube. Warm the t-butanol to 35 °C for three minutes.

5. Insert the Temperature Probe into the hot t-butanol. 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 to begin data collection.

6. Prepare a large beaker, 400-600 mL with an ice water bath. Fill beaker 2/3 full with tap water. Lower the test tube into the water bath. Add ice to your water bath to bring the temperature down. Make sure the water level outside the test tube is higher than the t-butanol level inside the test tube.

7. With a very slight up and down motion with the copper stirrer, continuously stir the t-butanol during the cooling.

8. Continue with the experiment until data collection has stopped (10 minute run). Use the hot water bath to melt the probe out of the solid t-butanol. Do not attempt to pull the probe out—this might damage it. Carefully wipe any excess t-butanol liquid from the probe with a paper towel or tissue. Weigh the test tube and t- butanol.

9. To determine the freezing temperature of pure t-butanol, 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 t-butanol. Close the statistics box.

Part II Freezing Temperature of a Solution of Aspirin and T-butanol

10. Store your data by choosing Store Latest Run from the Experiment menu. Hide the curve from your first run by clicking on the vertical axis label and unchecking the appropriate box. Click . Repeat steps 5-10 so that you have two trials of the freezing point t-butanol.

11. Measure out approximately 0.3 grams of aspirin into a weighing boat. Using a funnel, add the aspirin to the t- butanol already in the 4” test tube. The purpose of the funnel is to prevent aspirin from sticking to the inside wall of the test tube where it would be difficult (or impossible) to get into solution. Be careful not to get the funnel stem into the solvent! If you do, you will have to dump the solvent, clean and dry your apparatus, and start all over by weighing out a new portion of t-butanol. Determine the mass of the aspirin by weighing the test tube, solvent and aspirin. It may take several minutes for the aspirin to dissolve. Heat the

53

test tube gently with hot water and agitate very gently (be careful not to splash) until dissolution is complete. 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 aspirin-t-butanol solution, you need to determine the temperature at which the mixture initially started to freeze. Unlike pure t-butanol, cooling a mixture of aspirin and t-butanol 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 your data table.

13. Discard your solution and repeat from step 11 with a new sample of t-butanol and aspirin.

14. To print a graph of temperature vs. time showing all data runs:

a. Click on the vertical-axis label of the graph. To display both temperature runs, click More, and check the Run 1 and Latest Temperature boxes. Click .

b. Label both curves by choosing Text Annotation from the Insert menu, and typing “T-butanol” (or “Aspirin-t-butanol mixture”) in the edit box. Then drag each box to a position on or near its respective curve.

c. Print the graph.

PROCESSING THE DATA (METHOD 1)

1. Determine the difference in freezing temperatures, t, between the pure t-butanol (t1) and the mixture of t-butanol and aspirin (t2). Use the formula, t = t1 - t2.

2. Calculate molality (m), in mol/kg, using the formula, t = Kf • m (Kf = 8.28°C-kg/mol for t-butanol).

3. Calculate moles of aspirin solute, using the answer in Step 2 (in mol/kg) and the mass (in kg) of t-butanol solvent.

4. Calculate the experimental molar mass of aspirin, in g/mol. Use the original mass of aspirin from your data table, and the moles of aspirin you found in the previous step.

5. Compare your experimentally determined molar mass of aspirin with the known value.

6. Calculate the percent error.

Time

Freezing Point

54

PROCESSING THE DATA (METHOD 2)

Here is another method that can be used to determine the freezing temperature from your data in Part II. With a graph of the Part II data displayed, use this procedure:

1. Move the mouse pointer to the initial part of the cooling curve, where the temperature has an initial rapid decrease (before freezing occurred). Press the mouse button and hold it down as you drag across the linear region of this steep temperature decrease.

2. Click on the Linear Fit button, .

3. Now press the mouse button and drag over the next linear region of the curve (the gently sloping section of the curve where freezing took place). Press the mouse button and hold it down as you drag only this linear region of the curve.

4. Click again. The graph should now have two regression lines displayed.

5. Choose Interpolate from the Analyze menu. Move the mouse pointer left to the point where the two regression lines intersect. When the small circles on each cursor line overlap each other at the intersection, the temperatures shown in either examine box should be equal to the freezing temperature for the aspirin-t-butanol mixture.

6. Use the temperature to calculate T and your molar mass for aspirin. Compare your results from the two methods.

55

Freezing Point Depression Lab Report: Your report for this lab should include the following sections: Abstract:



Include a data table with all necessary mass measurements Include graphs for the freezing of t-butanol and t-butanol-aspirin solution

Results: Report the freezing point of pure t-butanol Report your calculated molar mass of aspirin Determine the percent error in your calculated molar mass (the actual molar mass of aspirin is 180.2 g/mol) Calculate the percent error in your determined molar mass of aspirin Be sure to attach hand written sample calculations to the back of your report


56

Chemistry 228

Pre-Lab: Kinetics


1. Write the general form of the rate law for the reaction you will be studying this week.

2. A first order reaction has a rate constant of 2.90 x 10-4 s-1. Calculate the half-life for this

reaction.

3. What is the overall order of a reaction that has the following rate law? Rate = [A]2[B]

4. For a reaction where the general form of the rate law is rate = [A]m[B]n, the following data were collected. What is the order of the reaction with respect to A? What is the order of the reaction with respect to B?

Initial Rate [A] [B] 0.01 M/s 0.025 M 0.025 M 0.01 M/s 0.025 M 0.050 M 0.09 M/s 0.075 M 0.025 M



57

The Rate and Order of a Chemical Reaction

OBJECTIVES


Conduct the reaction of KI and FeCl3 using various concentrations of reactants. Determine the order of the reaction in KI and FeCl3. Determine the rate law expression for the reaction.

INTRODUCTION

A basic kinetic study of a chemical reaction often involves conducting the reaction at varying concentrations of reactants. In this way, you can determine the order of the reaction in each species, and determine a rate law expression. Once you select a reaction to examine, you must decide how to follow the reaction by measuring some parameter that changes regularly as time passes, such as temperature, pH, pressure, conductance, or absorbance of light.

In this experiment you will conduct the reaction between solutions of potassium iodide and iron (III) chloride. The reaction equation is shown below, in ionic form.

2 I– (aq) + 2 Fe3+ (aq) → I2 (aq) + 2 Fe2+ (aq)

As this reaction proceeds, it undergoes a color change that can be precisely measured by a Colorimeter (see Figure 1). By carefully varying the concentrations of the reactants, you will determine the effect each reactant has on the rate of the reaction, and consequently the order of the reaction. From this information, you will write a rate law expression for the reaction.

Figure 1

MATERIALS

computer Vernier Colorimeter plastic cuvettes

58

PROCEDURE


2. Connect a Colorimeter to Channel 1 of the Vernier computer interface.

3. Start the Logger Pro program on your computer. Open the file “30b. Crystal Violet” from the Chemistry with Computers folder.

4. Set up and calibrate the Colorimeter.

f. Prepare a blank by filling an empty cuvette ¾ full with distilled water. Place the blank in the cuvette slot of the Colorimeter and close the lid.

g. If your Colorimeter has a CAL button, set the wavelength on the Colorimeter to 430 nm, press the CAL button, and proceed directly to Step 5. If your Colorimeter does not have a CAL button, continue with this step to calibrate your Colorimeter.

h. Choose Calibrate CH1: Colorimeter from the Experiment menu, then click . i. Turn the wavelength knob on the Colorimeter to the “0% T” position. j. Type 0 in the edit box. k. When the displayed voltage reading for Reading 1 stabilizes, click . l. Turn the knob of the Colorimeter to the Blue LED position (470 nm). m. Type 100 in the edit box. n. When the voltage reading for Reading 2 stabilizes, click , then click .

5. Obtain the materials you will need to conduct this experiment.

Three 25 mL graduated cylinders. Approximately 100 mL of 0.020 M KI solution in a 100 mL beaker. Approximately 100 mL of 0.020 M FeCl3 solution in a separate 100 mL beaker.

CAUTION: The FeCl3 solution in this experiment is prepared in 0.1 M HCl and should be handled with care.

Approximately 60 mL of distilled water in a third 100 mL beaker.

6. During this experiment you will conduct 5 trials. This step describes the process for conducting the trials using the Trial 1 volumes. When you repeat this process, use the correct volume for each trial based on the table below.

Trial FeCl3 (mL) KI (mL) H2O (mL)

1 10.0 10.0 0.0

2 10.0 5.0 5.0

3 5.0 10.0 5.0

4 7.5 5.0 7.5

5 5.0 7.5 7.5

o. Measure 10.0 mL of FeCl3 solution using a graduated cylinder and pour it into a large test

tube. p. Measure 10.0 mL of KI solution using a graduated cylinder. q. Prepare a clean cuvette.

r. Add the 10.0 mL of KI solution to the test tube containing 10.0 mL of FeCl3 solution. Cover the end of the test tube with your thumb and quickly invert to mix.

s. Within 15 seconds of mixing the two solutions, fill the cuvette ¾ full with the mixture. Wipe the outside of the cuvette with a tissue, place it in the Colorimeter, and close the lid and begin collecting absorbance data. The timing of this step is imperative to receiving

59

useful data; practice several times with water before attempting with the KI and FeCl3 solutions.

7. Click to begin collecting absorbance data. Data will be gathered for 2 minutes.

Observe the progress of the reaction in the beaker.

8. When the data collection is complete, carefully remove the cuvette from the Colorimeter. Dispose of the contents of the beaker and cuvette as directed. Rinse and clean the beakers and the cuvette for the next trial.

9. Examine the graph of the first trial. On the toolbar, select the Slope button. Slide the cursor to the initial time point. This tool will determine the initial slope and thus approximate the initial rate of the reaction. Record the slope as the initial rate of the Trial 1 reaction.

10. Repeat Steps 6–9 to conduct Trials 2–5. When you complete Step 9, use the same technique to analyze Trials 2–5 that you used to analyze Trial 1. Note: You will skip Step 6c in Trials 2–5.

DATA ANALYSIS

1. Calculate the initial molar concentration of FeCl3 and KI for each reaction and prepare a data table containing the concentrations of each reaction and the initial reaction rate.

2. What is the order of the reaction in FeCl3 and KI? 3. Write the rate law expression for the reaction.

60

Kinetics Lab Report: Your report for this lab should include the following sections: Abstract:


Introduction: Include why you did this experiment, relevant background, and general equations.

Data:

Include your data table with the initial concentrations of each reactant and the initial rate of each reaction Include graphs for each of your kinetic trials

Results: Report the order of the reaction with respect to each reactant State the rate law for the reaction Be sure to attach hand written sample calculations to the back of your report

Discussion: Discuss the experiment and any possible sources of error Explain how you determined the order of each reactant

As part of your discussion, answer the following question: 1. Is it possible to calculate the rate constant, k, from your data? If so, calculate the rate constant. If not, explain why not.

61

Chemistry 228

Pre-Lab: Chemical Equilibrium


1. Write the equilibrium constant expression for the experiment you will be studying this week.

2. If the equilibrium values of Fe3+, SCN- and FeSCN2+ are 9.5 x 10-4 M, 3.6 x 10-4 M and

5.7 x 10-5 M respectively, what is the value of Kc?

3. Write the general form of the dilution equation.

4. A solution is prepared by adding 18 mL of 0.200 M Fe(NO3)3 and 2 mL of 0.0020 M KSCN. Calculate the initial concentrations of Fe3+ and SCN- in the solution.



62

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) FeSCN2+(aq)

iron(III) thiocyanate thiocyanoiron(III)

When Fe3+ and SCN- are combined, a dynamic equilibrium is established between these two

ions and the FeSCN2+ ion. In order to calculate the equilibrium constant, Keq, for the reaction, it is necessary to know the concentrations of all the ions at equilibrium. In this experiment four separate equilibrium systems, or trials, containing different concentrations of these three ions (Fe3+, SCN-, and FeSCN2+) will be determined experimentally. The values for these equilibrium concentrations will be substituted into the equilibrium constant expression to see if Keq is indeed constant despite varied initial concentrations for the reactants. The Keq, is determined by using the Law of Mass Action.

aA + bB cC + dD

This equation gives the equilibrium constant expression of:

Keq = [C]c[D]d/[A]a[B]b

In order to determine the equilibrium concentrations for the three ions a standard solution needs

to be prepared. 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 initial

[Fe3+] in the standard solution is 900 times larger than [SCN-]i. According to LeChatelier's principle, which states that when a system in dynamic equilibrium is disturbed, the system responds so as to minimize the disturbance and return the system to a state of equilibrium. This high initial concentration Fe3+ ions on the left side of the equation forces the reaction far to the right, using up nearly 100% of the initial SCN– ions. Using stoichiometry and the balanced

equation, it is assumed that for every mole of FeSCN2+ produced, one mole of SCN– is used up. Thus since nearly all of the SCN- ions are consumed in order to minimize the disturbance, the

product’s concentration, [FeSCN2+]std, at equilibrium is assumed to be equal to the [SCN–]i.

Since the reaction produces the FeSCN2+ ions and this ion transmits the color red, the solution’s absorbance of blue light can be measured through the use of a colorimeter (see Figure 1). Because the red solutions absorb blue light very well, the blue LED setting on the Colorimeter is used. The computer-interfaced Colorimeter measures the amount of blue light absorbed by the colored solutions (absorbance, A).

63

Figure 1 Figure 2 According to Beer’s Law, there is a direct relationship between a solution’s concentration and its absorbance. In this case the concentration is the FeSCN2+ ion and the absorbance is blue light (470 nm). In other words, as the concentration of FeSCN2+ increases so will the absorbance of

blue light (see Figure 2). The concentration of FeSCN2+ for any of the equilibrium systems, trials 1-4, can be found by comparing the absorbance of each equilibrium system, Aeq, to the absorbance of the standard solution, Astd, according to the following equation:

[FeSCN2+]std/Astd = [FeSCN2+]eq/ Aeq

Since the concentration of [FeSCN2+]std is known and the all of the absorbances for the equilibrium solutions and the standard are measured and recorded all that needs to be done is to solve for the unknown.

[FeSCN2+]eq = Aeq Astd

X [FeSCN2+]std

Knowing the [FeSCN2+]eq allows you to determine the concentrations of the other two ions at

equilibrium. For each mole of FeSCN2+ ions produced, one less mole of Fe3+ and SCN- ions will be found in the solution (see the 1:1 ratio of coefficients in the equation on the previous

page). At equilibrium the [Fe3+] and [SCN-] can be determined according to the following equations:

[Fe3+]eq = [Fe3+]i – [FeSCN2+]eq

[SCN–]eq = [SCN–]i – [FeSCN2+]eq

Knowing the values of [Fe3+]eq, [SCN–]eq, and [FeSCN2+]eq, you can now calculate the value of Kc, the equilibrium constant.

64

OBJECTIVE

In this experiment, you will determine the equilibrium constant, Kc, for the following chemical reaction:

Fe3+(aq) + SCN–(aq) FeSCN2+(aq)

iron(III) thiocyanate thiocyanoiron(III)

MATERIALS

Vernier Colorimeter 1 plastic cuvette five 20 150 mm test tubes pipet bulb or pipet pump Serological pipet

PROCEDURE


2. Label four 20 150 mm 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 FeSCN2+ by pipetting 18 mL of 0.200 M Fe(NO3)3 into a

20 150 mm test tube labeled “5”. Pipet 2 mL of 0.0020 M KSCN into the same test tube. Stir thoroughly.

4. Connect the Colorimeter to the computer interface. Prepare the computer for data collection by opening the file “Lab 8 Equilibrium” from the Chemistry 228 folder of Logger Pro

5. Prepare a blank by filling a cuvette 3/4 full with distilled water. To correctly use a Colorimeter 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.

65

All solutions should be free of bubbles. Always position the cuvette with its reference mark facing toward the white reference mark

at the top of the cuvette slot on the Colorimeter. 6. Calibrate the Colorimeter.

a. Open the Colorimeter lid. b. Holding the cuvette by the upper edges, place it in the cuvette slot of the Colorimeter.

Close the lid. c. If your Colorimeter has a CAL button, Press the < or > button on the Colorimeter to select

a wavelength of 470 nm (Blue) for this experiment. Press the CAL button until the red LED begins to flash. Then release the CAL button. When the LED stops flashing, the calibration is complete. Proceed directly to Step 7. If your Colorimeter does not have a CAL button, continue with this step to calibrate your Colorimeter.

First Calibration Point

d. Choose Calibrate CH1: Colorimeter (%T) from the Experiment menu and then click .

e. Turn the wavelength knob on the Colorimeter to the “0% T” position. f. Type “0” in the edit box. g. When the displayed voltage reading for Reading 1 stabilizes, click .

Second Calibration Point

h. Turn the knob of the Colorimeter to the Blue LED position (470 nm). i. Type “100” in the edit box. j. When the displayed voltage reading for Reading 2 stabilizes, click , then click

. 7. You are now ready to collect absorbance data for the four equilibrium systems and the

standard solution.

a. Click to begin data collection. b. Empty the water from the cuvette. Rinse it twice with ~1 mL portions of the Test Tube 1

solution. c. Wipe the outside of the cuvette with a tissue and then place the cuvette in the Colorimeter.

After closing the lid, wait for the absorbance value displayed in the meter to stabilize. Then click , type “1” (the trial number) in edit box, and press the ENTER key.

d. Discard the cuvette contents as directed by your teacher. Rinse the cuvette twice with the Test Tube 2 solution and fill the cuvette 3/4 full. Follow the Step-c procedure to find the absorbance of this solution. Type “2” in the edit box and press ENTER.

e. Repeat the Step-d procedure to find the absorbance of the solutions in Test Tubes 3, 4, and 5 (the standard solution).

f. From the table, record the absorbance values for each of the five trials in your data table. g. Dispose of all solutions as directed by your instructor.

66

PROCESSING THE DATA

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 (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 mLtotal mL (0.0020 M)

In Test Tube 1, [SCN–]i = (2 mL / 10 mL)(0.0020 M) = 0.00040 M. Calculate this for the other three test tubes.

4. [FeSCN2+]eq is calculated using the formula:

[FeSCN2+]eq = Aeq Astd

[FeSCN2+]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 [FeSCN2+]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 – [FeSCN2+]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?

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Equilibrium Lab Report: Your report for this lab should include the following sections: Abstract:



Include your data table with initial concentrations of each reactant for each trial

Results: Include a results table with calculated Keq values for each trial and an average value for Keq Be sure to attach hand written sample calculations to the back of your report

Discussion: Discuss the experiment and any possible sources of error In addition, answer the following question as part of your report:

1. How are you Keq values to each other? Are they close enough to justify the assertion that an equilibrium constant is constant?

2. What factors could have led to variations in Keq between trials.

Documents

228 Winter 2012 Manual (1)