CHEM115 General Chemistry I Dr. Myton Class meets MTWR at 11:00 am

CHEM115 General Chemistry I

Dr. Myton

Class meets MTWR at 11:00 am

CHEM115 General Chemistry I

David M. Myton, Ph.D. (Dr. Myton)• CRW327 [email protected]• Pronto: dmyton• Professor of Chemistry

mailto:[email protected]

Introductions

• This class is not a direct competition

• Introduce your self to your neighbor Name Hometown Major

• Class survey Biology Chemistry Criminalistics Engineering Fisheries & Wildlife Geology Other

SYLLABUS

• Blackboard – get there through Anchor Access Pronto – IM with voice Documents, including slide summaries writing exercises

• http://edugen.wiley.com – homework, textbook, video office hours, student solutions

• Supplemental Instruction (SI)

• Bring to class: active chapter of text, calculator, i-clicker, lecture notes

http://edugen.wiley.com/

i-Clicker

• Unique student serial numbers• Credit given for participation and accuracy• Bring to EVERY class

Strongly Disagree--5EDisagree--4D

Neutral--3CAgreeFalseNo2B

Strongly AgreeTrue Yes1A

Essential High School Science Content: Properties of Matter

C1.1B Evaluate the uncertainties or validity of scientific conclusions using an understanding of sources of measurement error, the challenges of controlling variables, accuracy of data analysis, logic of argument, logic of experimental design, and/or the dependence on underlying assumptions.C2.2B Describe the various states of matter in terms of the motion and arrangement of the molecules (atoms)making up the substance.C4.2A Name simple binary compounds using their formulae.C4.2B Given the name, write the formula of simple binary compounds.C4.3A Recognize that substances that are solid at room temperature have stronger attractive forces than liquids at room temperature, which have stronger attractive forces than gases at room temperature.C4.3B Recognize that solids have a more ordered, regular arrangement of their particles than liquids and that liquids are more ordered than gases.


C4.8A Identify the location, relative mass, and charge for electrons, protons, and neutrons.C4.8B Describe the atom as mostly empty space with an extremely small, dense nucleus consisting of the protonsand neutrons and an electron cloud surrounding the nucleus.C4.8C Recognize that protons repel each other and that a strong force needs to be present to keep the nucleusintact.C4.8D Give the number of electrons and protons present if the fluoride ion has a -1 charge.C4.9A Identify elements with similar chemical and physical properties using the periodic table.C4.10A List the number of protons, neutrons, and electrons for any given ion or isotope.C4.10B Recognize that an element always contains the same number of protons.


C5.2A Balance simple chemical equations applying the conservation of matter.C5.2B Distinguish between chemical and physical changes in terms of the properties of the reactants and products.C5.2C Draw pictures to distinguish the relationships between atoms in physical and chemical changes.C5.4A Compare the energy required to raise the temperature of one gram of aluminum and one gram of waterthe same number of degrees.C5.5A Predict if the bonding between two atoms of different elements will be primarily ionic or covalent.C5.4B Predict the formula for binary compounds of main group elements.

Chemistry & Environmental Sciences Student Organization

1st Meeting: Wednesday September 2 at NOON in Crawford Hall Upstairs Lobby (w/ pizza)

1st Function: Camping at Muskellunge Lake Leaving Soo Saturday Sept 5 at NOON, returning Monday Sept 7 mid-day. Campsites and dinner provided by club

Chapter 1: Fundamental Concepts and Units of Measurement:

Learning Objectives• Upon completion of the chapter, the student should:• Know how chemistry fits into the sciences and everyday life.• Understand the difference between chemical reactions and

physical changes.• Understand the Law of Conservation of Energy.• Be able to convert between ºF, ºC and K.• Know the difference between precision and accuracy.• Have a basic understanding of significant figures.• Know the basic SI units.• Be able to convert between calories and joules.• Be able to determine the density, mass, or volume of a

substance when given two of these three variables.

Cumulative • accumulative: increasing by successive

addition; "the benefits are cumulative"; "the eventual accumulative effect of these substances" wordnet.princeton.edu/perl/webwn

• Incorporating all data up to the present en.wiktionary.org/wiki/cumulative

http://www.google.com/url?sa=X&start=4&oi=define&q=http://wordnet.princeton.edu/perl/webwn%3Fs%3Dcumulative&usg=AFQjCNERfTh767FdoHGgI7ovEprIyHlfOg

http://www.google.com/url?sa=X&start=5&oi=define&q=http://en.wiktionary.org/wiki/cumulative&usg=AFQjCNGL5rGfyaZQd3GFRe2oFeSXqLWIfg

http://www.google.com/url?sa=X&start=5&oi=define&q=http://en.wiktionary.org/wiki/cumulative&usg=AFQjCNGL5rGfyaZQd3GFRe2oFeSXqLWIfg

“Parfaits are delicious” - donkey

Juggling?

C:\Documents and Settings\dmyton.LSSU\Desktop\dmyton Documents\Courses\115 08F\Juggling_video.mp4

"Scientific knowledge is a body of statements of varying degrees of certainty -- some most unsure, some nearly sure, but none absolutely certain ... Now, we scientists are used to this, and we take it for granted that it is perfectly consistent to be unsure, that it is possible to live and not know." Richard Feynman (1918-1988)

Nobel Prize in Physics, 1965

15

Chapter 1:Fundamental Concepts and Units of Measurement

Brady & Senese 5th Ed

A Chemist’s ViewA Chemist’s View

2 H2(g) + O2 (g) 2 H2O(g)

Macroscopic

Symbolic

Particulate (Molecular)

01m11vd1.mov

1.1. Chemistry is important for anyone studying the sciences 17

Chemistry and the Sciences

• Chemistry- the study of the composition of matter and its transformations

• Matter- anything that takes up space and has mass

• Chemical reaction- change that results from the interaction of matter.

1.2. The scientific method helps us build models of nature 18

Scientific Method : Getting Started

Observe a Phenomenon-accurately describe something we see, taste, feel,

smell or hear

Pose A Question To Explain The Phenomenon

Form a Hypothesis-a tentative explanation of the

phenomenon


Scientific Method: Testing the Hypothesis

Experiment to Prove or Disprove Hypothesis

If experiment proves hypothesis,

form theory (theoretical model)

If experiment disproves hypothesis,

Pose new question or hypothesis

Continue experimentation. If results form pattern, considered

a law

It is the mark of an educated mind to be able to entertain a thought without accepting it. Aristotle (384 BC - 322 BC)

Question:

Which statement is a hypothesis?a: Objects on Earth are attracted by

gravity.b: When pushed off the table, my

chemistry book will fall to the floor.c: Opposite charges repel each other.d: Mass can be converted into energy.


Scientific Method Case Study: The Process of Growth

• A child sees that a seed, when planted in soil, watered, and exposed to sunlight, grows to form a flower. He concludes that all living things require sunlight, water, and burial in soil to grow.

• Build a case for rebuttal using the scientific method.


Your Turn!

Which of the following is not a hypothesis for the observed plant growth?A. soil is necessary to all growth

B. light is essential to growth of the seed

C. water is required to allow growth

D. plants grow to a greater height if they receive fertilizer

E. none of the above


Your Turn!

A chicken egg is buried, left in the sun, and watered. A second egg is left above the soil, watered and left in the sun. Would this prove that soil is necessary to growth?

A. Yes

B. No


The Scientific Method- Evaluating The Data

A theory is an explanation (based on well-tested, internally consistent experimental results) about why the phenomenon may occur it should explain currently available data It should be as simple as possible It should clearly show underlying connections It should accurately predict future behaviors


The Scientific Method is Cyclical

Question:

Which describes a tested explanation of behavior of nature?

a: a scientific law b: a theory c: a hypothesis d: empirical facts

http://www.amazon.com/gp/product/images/0716779390/sr=8-1/qid=1179368334/ref=dp_image_0/103-9725723-6451808?ie=UTF8&n=283155&s=books&qid=1179368334&sr=8-1


Atomic Theory Helps Us Visualize Matter

• Air inflates a balloon air must be composed of matter the matter is colliding with the walls

of the container.

• A leaf floats on water’s surface water is composed of particles that

occupy space

• A leaf falls through air, but rests on water’s surface particles are closer in liquid than in

gases


Models Helps Us Visualize Matter


1.3. Matter is Composed of Elements, Compounds, and Mixtures 31

Changes in Matter

• Chemical change- a process that results in the formation of a new substance

• Evidence? Formation of a new solid, new liquid, new gas, temperature change, or an unexpected color change

• Physical change- a process that results in no new substance, but that may change the state of those present, or the proportions

Question:

What properties change when a substance undergoes a chemical reaction?

a: physical propertiesb: chemical propertiesc: both chemical and physical propertiesd: neither chemical nor physical properties



Learning Check: Chemical Or Physical Change?

Chemical Physical

Magnesium burns when heated in a flame

Magnesium metal tarnishes in air

Magnesium metal melts at 922K

Grape Kool-aid lightens when water is added


Your Turn!

Which of the following is not a chemical change?A. a match burns in air

B. ice melts in air

C. an aluminum door whitens in air

D. all of these

E. none of these


• Matter can be classified (figure 1.10 p12):

(Atom)(Molecule orFormula unit)

Elements are substances that cannot be decomposed by chemical means into simpler substances

chemical symbol Most are one or two letters First letter is always capitalized All remaining letters are

lowercase Names and chemical symbols of

the elements are listed on the inside front cover of the book

Elements

• Learn the name, spelling and symbol for elements #1-30, Au, Ag, Hg, Pb, Br, I

K

Nitrogen

Cl

Copper

Sodium

SymbolName


Atomic naming

Sb antimony stibium K potassium kalium

Cu copper cuprum Sn tin stannum

Au Gold aurum Na sodium natrium

Ag silver argentum W tungsten wolfram

Fe iron ferrum Hg mercury hydragyrum

Pb lead plumbum


What Is A Compound?

• Compounds - formed from two or more atoms of different elements combined in a fixed proportion

• Have different characteristics than the elements that compose them

• Can be broken down into elements by some chemical changes

A A MOLECULEMOLECULE is the smallest is the smallest

unit of a unit of a compoundcompound that retains that retains thethe chemical characteristics of the chemical characteristics of the compoundcompound..

MOLECULARMOLECULAR FORMULA FORMULA

H2O

C8H10N4O2 - caffeine

01m06an1.mov


Mixtures

• mixtures consist of varying amounts of two or more elements or compounds

• Homogeneous mixtures or “solutions”- have the same properties throughout the sample Brass, tap water

• Heterogeneous mixtures- consist of two or more phases Salad dressing, Coca-Cola ™


Learning Check: Classification

Sand Ice

(H2O)

Flour Table Salt (NaCl)

Pure

Element

Compound

Molecule

Heterogeneous Mix

Homogeneous Mix

44

Classification Of Matter By State

Classification by state is based on packing, motion, and shape Solids have fixed shape and volume Liquids have fixed volume, but take the container shape Gases have to expand to fill the shape and volume of

the container

1.4. Properties of matter can be classified in different ways 45

Properties Of Matter

• Chemical properties describe the behavior of the matter that leads to the formation of a new substance: the "reactivity" of the substance

• Physical properties can be observed about the matter alone, without changing the composition


Learning Check: Chemical or Physical Property?

Chemical Physical

Magnesium metal is grey

Magnesium metal tarnishes in air

Magnesium metal melts at 922K

Magnesium reacts violently with hydrochloric acid


Your Turn!

Which of the following is a chemical property?

A. water is colorless

B. water reacts violently with solid Na metal

C. water dissolves table salt

D. all of these

E. none of these


Question:

Intensive properties are Independent of the quantity of material present

Which is an extensive physical property?

a: mass b: melting point c: reactivity with water d: temperature


1.5 Measurements are essential to describe properties 49

Measurements are Observations

• Qualitative observations are non-numerical-- ask “what” or “how” or “why”

• Quantitative observations are numerical--ask “how much” and are also called measurements

• This course is general chemistry with quantitative analysis


Your turn!

Which of the following is a quantitative observation?

A. the height of the plant

B. the mass of water added

C. the temperature of the day

D. all of the above

E. none of the above


Measurement

http://images.google.com/imgres?imgurl=http://www.chime.ucla.edu/measurement/images/tape.jpg&imgrefurl=http://www.chime.ucla.edu/measurement/AboutEXPORT.htm&h=197&w=180&sz=4&hl=en&start=36&tbnid=sMtma56lYWgqbM:&tbnh=104&tbnw=95&prev=/images%3Fq%3Dmeasurement%26start%3D20%26ndsp%3D20%26svnum%3D10%26hl%3Den%26lr%3D%26safe%3Dactive%26rls%3Dcom.microsoft:en-US%26sa%3DN


• Always involve a comparison

• Require units

• Involve numbers that are inexact (estimated). This uncertainty is due to the limitations of the observer and the instruments used

• In science, all digits in a measurement up to and including the first estimated digit are recorded

Measurements:

Which types of numbers are considered “exact?” Below are the general rules.1. Conversions between units within the English System are exact.

e.g. 12 in = 1 ft or 12 in/1 ft (In this conversion, 12 and 1 are both exact.) 2. Conversions between units within the Metric System are exact.

e.g. 1 m = 100 cm or 1 m/100 cm (In this conversion, 1 and 100 are both exact.) 3. Conversions between English and Metric system are generally NOT exact. Exceptions will

be pointed out to you.e.g. 1 in = 2.54 cm exactly (1 and 2.54 are both exact.)e.g. 454 g = 1 lb or 454 g/1 lb (454 has 3 sig. fig., but 1 is exact.)

4. “Per” means out of exactly one.

e.g. 45 miles per hour means 45 mi = 1 hr or 45 mi/1 hr. (45 has 2 sig. fig. but 1 is exactly one.)

5. “Percent” means out of exactly one hundred.

e.g. 25.9% means 25.9 out of exactly 100 or 25.9/100 (25.9 has 3 sig. fig., but 100 is exact.) 6. Counting numbers are exact. Sometimes it is hard to decide whether a number is a “counting

number” or not. In most cases it would be obvious. Ask when in doubt.e.g.There are 5 students in the room. (5 would be an exact number because you cannot have a

fraction of a student in the room.)


Measurements and units

• In the U.S., we use the Imperial (USCS) System (United States Customary System units)

• The scientific community (and most of the world) uses the metric system

• Variations in the metric system exist, thus a standard system is used: International System of Units (SI)

• SI units we will use now: Length (m) Mass (kg) Time (s) Temperature

(K)


Mass- Matter Content

USCS: oz (avdp.), lb, T

Metric: g

SI: kg


Length

USCS: in, ft, yd, mi

Metric: L, cm3

SI: m


Volume-bulk

• measured directly, using equipment for volumetric measure

• calculated using dimensional (length) information and appropriate formulas. 1 cm3= 1mL

• USCS: fl. oz., pt., qt., gal

• Metric: L, cm3

• SI: m3


Temperature

• USCS: °F• Metric: °C• SI: K

C 1

K1C15.273CK

tT

F32C5F9

CF

tt


Your Turn!

Which of the following is the lowest temperature?

A. 300. K

B. 16 ºC

C. 55 ºF

D. they are the same


Question


1.6. Measurements always contain some uncertainty 62

• Because each measurement involves an estimate, measurements always have error.

• Record all measured numbers, including the first estimated digit

• These digits are called significant digits or significant figures

• Exact numbers have infinite significant digits

Measurement Error


Significant Digits In A Measurement Are Limited By Instrument Precision

• Using the first thermometer, the temperature is 21.3 ºC (3 significant digits)

• Using the more precise (second) thermometer, the temperature is 21.32 ºC (4 significant digits)

Measurement Exercise

• Complete the following table

• Convert the area in mm2 to in2 and from in2 to mm2 and record.

• Note that 2.54 cm = 1 in

mm in

Length

Width

Area

Converted area:


• Errors-inherent error due to the equipment or procedure Changing volume due to thermal expansion or contraction

(temperature changes) Improperly calibrated equipment procedural design allows variable measurements

• Mistakes-blunders that you know that you have made. Do not use these data Spillage Incomplete procedures Reading scales incorrectly Using the measuring device incorrectly

Errors Arise From A Number Of Sources Including:


Reducing Error:

• Errors can often be detected by making repeated measurements

• Error can be reduced by calibrating equipment

• The average or mean reduces data variations: it helps find a central value


• An accurate measurement is close to the true or correct value, a “hole-in-one”

• A precise measurement is close to the average of a series of repeated measurements

• When calibrated instruments are used properly, the greater the number of significant figures, the greater is the degree of precision for a given measurement

Accuracy vs. Precision

Measurements with lots of scatter are probably considered: a: not accurate and not precise. b: not accurate, but precise. c: not precise, but accurate. d: precise and accurate.

Question

Question

Nonzero digits in a measured number are always significant

Zeros must be considered more carefully: Zeros between significant digits are significant Zeros to the right of the decimal point are

always counted as significant Zeros to the left of the first nonzero digit are

never counted as significant Zeros at the end of a number without a decimal

point are assumed not to be significant

How many significant figures?

200.0

12100

21000

1010

200.001

0.98700

0.000012

2.200002

125

1.25e2

1.250e-3

125000

0.00125

6.2303e23

6 230 300 000 000

• Measurements limit the precision of the results calculated from them

• Rules for combining measurements depend on the type of operation performed: Multiplication and division

The number of significant figures in the answer should not be greater than the number of significant figures in the least precise measurement.

figs.) sig. (2

13

figs.) sig. (2figs.) sig. (4 figs.) sig. (3

0.642.751 3.14

Addition and SubtractionThe answer should have the same number of

decimal places as the quantity with the fewest number of decimal places

3.247 3 decimal places 41.36 2 decimal places +125.2 1 decimal place 169.8 answer rounded to 1 decimal place

Note: Remember that some numbers are exact. Numbers that come from definitions or direct counts have no uncertainty and can be assumed to contain an infinite number of significant figures.

Question

Question

What is the proper way to report the sum 1.150 m + 3.3 m? a: 4.45 m b: 4.4 m c: 4.5 m d: 4.450 m

Question

Dimensional Analysis

• We plan to bike the 14.1 miles from the university to Brimley State Park What is the distance in km given that 0.6215 mi = 1 km?

• What is the distance in meters?


• We plan to bike the 14.1 miles from the university to Brimley State Park If we travel at a rate of 17 miles per hour how many minutes will the trip take?


• If we drive the 14.1 miles from the university to Brimley State Park and get 22.4 mpg in our car, what will the cost of the trip be if gas is $2.56/gal?

1.7 Units can be converted using the factor-label method 81

USCS Unit Conversions

Mass Volume Distance16 oz. (avdp.) = 1 lb.2000 lb. = 1 T.

3 tsp. = 1 Tbsp.16 Tbsp. = 1 c.2 c. = 1 pt.2 pt. = 1 qt.4 qt. = 1 gal.8 fl. oz. = 1 c.

12 in. = 1 ft.3 ft. = 1 yd.1760 yd. = 1 mi.


USCS to Metric Metric to USCS Length 1 in. = 2.54 cm 1 m = 39.37 in

1 yd = 0.9144 m 1 km = 0.6215 mi1 mi = 1.609 km

Mass 1 lb = 453.6 g 1 kg = 2.205 lb1 oz = 28.35 g

Volume 1 gal = 3.785 L 1 L = 1.0567 qt1 qt = 946.4 mL1 oz (fluid) = 29.6 mL

It is also useful to know that 1 mL = 1 cm3=1 cc

USCS And Metric Units Are Related Using “Critical Links”


Building Conversion Factors in Unit Conversions

1. Write the number to be converted as a fraction (with units)

2. Identify the target units3. Are the starting units in the same system as the

target? If not, you will need a critical link. USCS→USCS Conversions: Write down the

conversion factors from smallest to largest . metric →metric conversions: Write down the

definitions of all prefixed units.


Building Conversion factors (cont).

4. Use the form of the conversion factor that allows the units to cancel--they must be on opposite levels of the fraction to cancel.

5. Continue adding conversion factors until the units match the target units.

2nd Check- are all units written on the page two times? If so, you have enough info to start the problem.

What conversion factors would be necessary to convert miles per gallon to km per liter? a: (1.6 km / 1 mi) and (1 gal / 3.79 L) b: (1 mi / 1.6 km) and (1 gal / 3.79 L) c: (1.6 km / 1 mi) and (3.79 L / 1 gal) d: (1 mi / 1.6 km) and (3.79 L / 1 gal)

Question

What is the volume of a box that measures 50 cm by 1.0 m by 2000 mm? a: 100,000 m3

b: 10,000 m3

c: 100 m3

d: 1.0 m3

e: 1 m3

Question

If there are exactly 2.54 cm in one inch , what is the volume of a cube 1 foot on each side in units of cubic centimeters?

Chalcopyrite, the principle ore of copper (Cu) contains 34.63 percent Cu by mass. How many grams of Cu can be obtained from 5.11e3 kg of the ore?

If a bracelet is made of silver and copper with a mass of 38.9 g contains 32.3 g of silver, what is the percentage of silver and of copper. How many grams of copper are in a 50 gram bracelet?

Density (d) is an intensive property defined as the ratio

of an objects mass (m) to volume (v), d = m/v

Density(g/cm3)Water 1.00Aluminum 2.70Iron 7.86Gold 19.3Air 0.0012

•characteristic of pure substances at a specified temperature •Since most substances expand when heated, densities decrease when heated.•units : g/L for gases and g/mL for solids and liquids.

Problem Solving• The density of ethanol, a colorless liquid that is

commonly known as grain alcohol, is 0.798 g/mL. Calculate the mass of 17.4 mL of the liquid

• Calculate the volume of 17.4 g of the liquid

A student pipets 25.0 0mL of isopropyl alcohol into an empty flask weighing 35.182 g. She finds the mass of the flask + alcohol is 54.707 g. Calculate the density of the alcohol.

The density of a solution of sulfuric acid is 1.285 g/cm3, and it is 38.08% acid by mass. What volume of the acid solution in mL do you need to supply 125 g of sulfuric acid?

A solution of sugar in water has a density of 1.05 g/cm3. If you have 250. mL of the solution and if the solution is 8.1% by weight sugar, how many grams of sugar are in the solution.

Air at room temperature has a density of about 0.0012 g/cm3. What is the mass of 1.0 L of air? a: 1.2 g b: 12 g c: 0.0012 g d: none of these

Question

1.8. Density is a useful intensive property 96

Learning Check:

A crash sounds from the lab- a large vial of mercury has fallen from a broken shelf. We call the hazardous materials team to report the spill, about 2.0 quarts of mercury. They ask for the mass- what is it? (hint: d=13.69g/mL & 1 L = 1.0567 qt)

Amethyst is a colored form of the mineral quartz in which the purple color comes from traces of the element manganese. To determine the density of amethyst, you take a stone having a mass of 15.25 g and place it in a 100. mL graduated cylinder containing 45.0 mL of water. On adding the stone the water surface rises to the 50.8 mL mark. What is the density of amethyst?