WELCOME TO CHEM 1WITH MRS. KAUR

Learn about each other Learn how to be a good student and a

professional employee Learn general chemistry General course info in syllabus My website – use it! First two chapters will be FAST – they are a

review! Look at slide: tell me what you can about it.

1- 1

Tell me all you can about this picture:

1- 2

1- 3

Keys to the Study of Chemistry

Chapter 1

1- 4

HOW TO LEARN CHEMISTRY?

Think of it as a foreign languageMemorize the new words, names of elements and

compounds, definitionsUse flash cards Learn how to be a scientistAlso question everything you read or hear - even me

or the textbooks. Scientific exploration begins with questions, and leads

to more questions.Read “Study tips and how to take a test” in your packet

of handouts.

1- 5

Is the study of matter,

its properties,

the changes that matter undergoes,

and

the energy associated with these changes.

1- 6

Definitions

Chemical Properties those which the substance shows as it interacts with, or transforms into, other substances such as flammability, corrosiveness

Matter anything that has mass and volume -the “stuff” of the universe: books, planets, trees, professors, students

Composition the types and amounts of simpler substances that make up a sample of matter

Properties the characteristics that give each substance a unique identity

Physical Properties those which the substance shows by itself without interacting with another substance such as color, melting point, boiling point, density

1- 7

Figure 1.1

A Physical change B Chemical change

The distinction between physical and chemical change.

1- 8

Some Characteristic Properties of CopperTable 1.1 (4th ed.)

Physical Properties Chemical Properties

reddish brown, metallic luster

easily shaped into sheets(malleable) and wires

(ductile)

good conductor of heatand electricity

density = 8.95 g/cm3

melting point = 10830C

boiling point = 25700C

slowly forms a basic blue-greensulfate in moist air

reacts with nitric acid and sulfuric acid

slowly form a deep-bluesolution in aqueous ammonia

1- 9

Figure 1.2

The physical states of matter.

What are the phase changes? There are six - see how many you can remember…

1- 10

REVIEW OF DEFINITIONS YOU SHOULD ALREADY KNOW:Phases of matter: solid, liquid, and gas, and

phase changes between them

Solutions of matter: s, l or g mixtures. Solute dissolved in a solvent = a solution.

Physical properties and chemical properties (see previous slides)

Extensive vs. Intensive properties

Separation processes

1- 11

Sample Problem 1.2 Distinguishing Between Physical and Chemical Change

PROBLEM: Decide whether each of the following process is primarily a physical or a chemical change, and explain briefly:

SOLUTION:

(a) Frost forms as the temperature drops on a humid winter night.

(b) A cornstalk grows from a seed that is watered and fertilized.

(c) Dynamite explodes to form a mixture of gases.

(d) Perspiration evaporates when you relax after jogging.

(e) A silver fork tarnishes slowly in air.

(a) physical change (b) chemical change (c) chemical change

(d) physical change (e) chemical change

1- 12

Energy due to the position of the object or energy from a potential chemical reactionPotential Energy

Nonpotential Energy

Energy due to the motion of the object E=1/2 mv2

Heat transfer q=mcpTLight energy e=hRadiant energy, blackbody radiationEtc.

EnergyEnergy is the capacity to do work.

Potential and nonpotential energy can be interconverted. (See figure 1.3 in your textbook)

1- 13

Scientific Approach: Developing a Model

Observations : Natural phenomena and measured events; universally consistent ones can be stated as a natural law.

Hypothesis: Tentative proposal that explains observations.

Experiment: Procedure to test hypothesis; measures one variable at a time.

Model (Theory):Set of conceptual assumptions that explains data from accumulated experiments; predicts related phenomena.

Further Experiment: Tests predictions based on model.

revised if experiments do not support it

altered if predictions do not support it

Fig 1.4

1- 14

A Systematic Approach to Solving Chemistry Problems

1. Problem statement

2. Plan

Clarify the known and unknown.

Suggest steps from known to unknown.

Prepare a visual summary of steps.

3. Solution

4. Check

5. Comment and 6. Follow-up Problem

1- 15

Table 1. 1 SI Base Units

Physical Quantity (Dimension) Unit Name

Unit Abbreviation

mass

meter

kg

length

kilogram

m

time second s

temperature kelvin K

electric current ampere A

amount of substance mole mol

luminous intensity candela cd

Know the ones outlined in red.

volume liter L

1- 16

Common Decimal Prefixes Used with SI Units Table 1.2

Prefix PrefixSymbol

Word ConventionalNotation

ExponentialNotation

tera T trillion 1,000,000,000,000 1x1012

giga G billion 1,000,000,000 1x109

mega M million 1,000,000 1x106

kilo k thousand 1,000 1x103

hecto h hundred 100 1x102

deka da ten 10 1x101

----- ---- one 1 1x100

deci d tenth 0.1 1x10-1

centi c hundredth 0.01 1x10-2

milli m thousandth 0.001 1x10-3

micro millionth 0.000001 1x10-6

nano n billionth 0.000000001 1x10-9

pico p trillionth 0.000000000001 1x10-12

femto f quadrillionth 0.000000000000001 1x10-15

Memorize these prefixes!

Just for fun!

Try to find out the distance from the earth to the sun in meters:_____________

Also the diameter of the helium atom:_________________

1- 17

1- 18

English EquivalentLength

1 kilometer(km)1000(103)m

0.62mi

Common SI-English Equivalent QuantitiesTable 1.3

Quantity SI UnitSI Equivalent

English to SI Equivalent

1 kilometer(km)1000(103)m

0.62miles(mi)1 mi = 1.61km

1 meter(m)100(102)cm

1.094yards(yd)1000(103)mm

39.37inches(in)1 yd = 0.9144m

1 foot (ft) = 0.3048m

1 centimeter(cm)0.01(10-2)m

0.3937in1 in = 2.54cm

(exactly!)

1 mi = 5280 ft

1 yd = 3 ft, etc.

1- 19

Volume

1,000,000(106) cubic centimeters

35.2cubic feet (ft3)

1 cubic meter(m3)

1 ft3 = 0.0283m3

1 cubic decimeter(dm3)1000cm3

0.2642 gallon (gal) 1.057 quarts (qt) 1 gal = 3.785 dm3

1 qt = 0.9464 dm3

1 cubic centimeter (cm3)0.001 dm3

0.0338 fluid ounce 1 qt = 946.4 cm3

1 fluid ounce = 29.6 cm3

English Equivalent

Quantity SI UnitSI Equivalent

English to SI Equivalent

Common SI-English Equivalent QuantitiesTable 1.3

1 L = 1.0567 qt

1 gal = 4 qt

1- 20

English Equivalent

Quantity SI UnitSI Equivalent

English to SI Equivalent

Mass

1 kilogram (kg)

1000 grams 2,205 pounds (lb)

1 (lb) = 0.4536 kg

1 gram (g)

1000 milligrams

0.03527 ounce(oz)

1 lb = 453.6 g

1 ounce = 28.35 g

Common SI-English Equivalent QuantitiesTable 1.3

1 lb = 453.6 g

1 lb = 16 oz

1- 21

English-Metric Conversions to Memorize

1" = 2.54 cm exactly

1 lb = 453.6 g

1 cal = 4.184 J

1.0 L = 1.0567 qt

1.000 atm = 14.70 psi

1- 22

Practice:

The O-H bond length is 95.7 pm. Convert it to nm.

Convert 750 mL to L:

Weight = mass x accelerationMetric:

Newton = mass in grams x accel of gravityEnglish:

Pound = mass in slugs x accel of gravityOn earth we use 1 lb = 453.6 g, even though we

should use slugs!

Practice:Convert 2.65 g to kg: Convert 2.65 g to mg:

1- 23

The three dimensions

Length is one dimensional.

Area is two-dimensional: area = length x width

Units will be ft2, m2, cm2, etc.

Volume is three-dimensionalV = length x width x height

Units will be L, mL, cm3,

MEMORIZE: 1 L = 103 mL, 1 mL = 1 cm3

1- 24

Equipment measuring liquid volume:From previous lab experience, describe a

graduated cylinder, a buret, a volumetric flask, and a pipette.

Decide which of them would provide the most accurate volume reading.

1- 25

1- 26

DENSITY

Density is a physical property that represents the amount of volume a certain mass displaces

D = m/V

Find the mass of 250.0mL of ethanol if D = 0.789g/mL at 20C.

1- 27

Substance Physical State Density (g/cm3)

Densities of Some Common Substances*Table 1.5 (4th ed.)

Hydrogen Gas 0.0000899

Oxygen Gas 0.00133

Grain alcohol Liquid 0. 789

Water Liquid 0.998

Table salt Solid 2.16

Aluminum Solid 2.70

Lead Solid 11.3

Gold Solid 19.3

*At room temperature(200C) and normal atmospheric pressure(1atm).

1- 28

Figure 1.6 The freezing and boiling points of water.

1- 29

Temperature Scales and Interconversions

Kelvin ( K ) - The “Absolute temperature scale” begins at absolute zero and only has positive values.

Celsius ( oC ) - The temperature scale used by science, formally called centigrade, most commonly used scale around the world; water freezes at 0oC, and boils at 100oC.

Fahrenheit ( oF ) - Commonly used scale in the U.S. for our weather reports; water freezes at 32oF and boils at 212oF.

K = oC + 273.15 oC = K - 273.15

oF = 1.8 oC + 32oC = [oF - 32 ]/ 1.8

Memorize these!!!

1- 30

TEMPERATURE CALCULATIONSConvert room temperature of 72.0oF to oC and

to Kelvin. Apply sig fig rules of math. (Do you remember them?)

In the northeast, we say “below zero” for very cold temperatures, but this is referring to Fahrenheit. What is “five below zero” in Celsius?

1- 31

PERCENT CALCULATIONS

Percent: can be by mass or by some other measurement.

Percent = part/whole * 100

If a 17.6 g silver bracelet is only 14.1 g Ag, because it is in an alloy with copper for strengthening. What is the percent-mass Ag? And the percent-mass Cu?

1- 32

The number of significant figures in a measurement depends upon the measuring device.

Figure 1.7

32.30C32.330C

The last digit, shaded gray, is the estimated digit that you would read. It is the digit with some uncertainty.

1- 33

Rules for Determining Which Digits are Significant:*

All digits are significant

•Make sure that the measured quantity has a decimal point.•Start at the left of the number and move right until you reach the first nonzero digit.•Count that digit and every digit to its right as significant.

Numbers such as 5300 L are assumed to only have 2 significant figures. A terminal decimal point is often used to clarify the situation, but scientific notation is the best!

except zeros that are used only to position the decimal point.

Zeros that end a number and lie either after or before the decimal point are significant; thus 1.030 ml has four significant figures, and 5300. L has four significant figures also.

*See informational handout “The Rules of Significant Figures”

1- 34

Rules for Significant Figures in Calculations

1. For addition and subtraction. The answer has the same number of decimal places as there are in the measurement with the fewest decimal places.

106.78 mL = 106.8 mL

Example: subtracting two volumes

863.0879 mL = 863.1 mL

865.9 mL - 2.8121 mL

Example: adding two volumes 83.5 mL

+ 23.28 mL

1- 35

= 23.4225 cm3 = 23 cm39.2 cm x 6.8 cm x 0.3744 cm

2. For multiplication and division. The number with the least

certainty limits the certainty of the result. Therefore, the answer

contains the same number of significant figures as there are in the

measurement with the fewest significant figures.

Rules for Significant Figures in Answers

Multiply the following numbers:

1- 36

Rules for Rounding Off Numbers

1. If the digit removed is greater than or equal to 5, the preceding number increases by 1. 5.379 rounds to 5.38 if three significant figures are retained and to 5.4 if two significant figures are retained.

2. If the digit removed is less than 5, the preceding number is unchanged. 0.2413 rounds to 0.241 if three significant figures are retained and to 0.24 if two significant figures are retained.

Be sure to carry two or more additional significant figures through a multi-step calculation and round off only the final answer.

Practice

Pair up with your lab partner or someone in your study group. Look at sample problem 1.9, sig figs and rounding. Then do its follow-up problem. Be prepared to show me your work.

1- 37

1- 38

Precision and Accuracy Errors in Scientific Measurements

Random Error - In the absence of systematic error, some values that are higher and some that are lower than the actual value.

Precision -Refers to reproducibility or how close the measurements are to each other.

Accuracy -Refers to how close a measurement is to the real value.

Systematic error - Values that are either all higher or all lower than the actual value.

1- 39

Figure 1.8

precise and accurate

precise but not accurate

Precision and accuracy in the laboratory.

1- 40

systematic error

random error

Precision and accuracy in the laboratory.Figure 1.8continued