Chapter One

Matter and Measurement

http://www.chemistry.armstrong.edu/

nivens/course_list.htm

OWL HOMEWORK REQUIRED!!!

2

Matter and Energy -• Chemistry

– the study of matter and its changes

– the "central science"

Matter– Anything that has mass and occupies space. (mass is the amount of matter, weight is force of gravitational attraction on the mass)

• Energy– The ability to do work or transfer heat.

• Scientific (natural) law– A general statement based the observed behavior of matter to which no exceptions are known.

3

Natural Laws

• Law of Conservation of Mass- amount of matter (mass) is a constant under a chemical or physical change

• Law of Conservation of Energy – total energy in a process is fixed, only converted from one form to another.

• Law of Conservation of Mass-Energy- Universal sum of mass + energy is a constant

• Einstein’s Relativity

• E=mc2

4

States of Matter

• Solid – definite shape and volume

• Liquid – definite volume (assumes shape of container)

• Gas – no shape or volume (expand to fill volume of container

•Change of States

–Heating or cooling

5

Chemical and Physical Properties

• Chemical Properties - observed during a chemical change (change in composition)– rusting or oxidation

– chemical reactions

• Physical Properties – observed as a physical state a compound resides in– changes of state

– density, color, solubility

• Extensive Properties - depend on quantity

• Intensive Properties - do not depend on quantity

6

Mixtures, Substances, Compounds, and Elements

• Substance

– identical composition and properties

• Elements

– cannot be decomposed into simpler substances via chemical reactions

• Elemental symbols

– found on periodic chart

7

Mixtures, Substances, Compounds, and Elements

• Compounds

– two or more elements in a definite ratio by mass - can be decomposed into the constituent elements

• Water– hydrogen and oxygen

• Carbon dioxide – carbon and oxygen

8

Mixtures, Substances, Compounds, and Elements

• Mixtures

– two or more substances

– homogeneous mixtures –same phase

• Also termed “solutions”

– heterogeneous mixtures

9

Math Review and Measurements

• Make measurements to understand the environment– Humans—sight, taste, smell, sound…

• Limited and biased

– Use instruments—meter sticks, thermometers, balances• More accurate and precise

– All measurements have units—METRIC SYSTEM vs. British System

10

Units of Measurement

Definitions

• Time

– Interval or duration of forward events

• Mass

– measure of the quantity of

matter in a body

• Weight

– measure of the gravitational attraction for a body (w=mg)

11

Units of Measurement

Definitions

• Length

– Measure of space in any direction

12

Units of Measurement

• Volume

– Amount of space

occupied by a system

– Derived unit, mL or cc, cm3

13

Measurements in Chemistry

QuantityQuantity UnitUnit SymbolSymbol

� length meter m

� mass kilogram kg

� time second s

� current ampere A

� temperature Kelvin K

� amt. substance mole mol

14

Measurements in ChemistryUse Metric System

NameName SymbolSymbol MultiplierMultiplier

• mega M 106 (1,000,000)

• kilo k 103 (1,000)

• deka da 10

• deci d 10-1 (0.1)

• centi c 10-2 (0.01)

15

Measurements in ChemistryMetric Prefixes

NameName SymbolSymbol MultiplierMultiplier

• milli m 10-3(0.001)

• micro µ 10-6(0.000001)

• nano n 10-9

• pico p 10-12

• femto f 10-15

16

Units of Measurement

Definitions

• Mass

–mass is the amount of matter

Weight

– force of gravitational attraction on mass

17

Units of Measurement

Common Conversion Factors

• Length

– 1 m = 39.37 inches

– 2.54 cm = 1 inch

• Volume

– 1 liter = 1.06 qt

– 1 qt = 0.946 liter

18

Use of Numbers

• Exact numbers– 1 dozen = 12 things for example

• Measured Numbers—– use rules for significant figures

– Use scientific notation were possible

• Accuracy – how closely measured values agree with the correct value

• Precision– how closely individual measurements agree with each other

19

Use of Numbers

• Significant figures

– digits believed to be correct by the person making the measurement

• Exact numbers have an infinite number of significant figures

12.000000000000000 = 1 dozen

because it is an exact number

20

Significant figures

• Calculators give 8+ numbers

• People estimate numbers differently

• Scientists develop rules to help you determine which digits are “significant”

• Dictated by the precision (graduation) on your measuring device

• In the lab, the last significant digit is the digit you have to estimate

21

Use of Numbers

Significant Figures - Rules

• Non zero numbers are significant

• Leading zeroes are never significant0.000357 has three significant figures

• Imbedded zeroes are always significant3.0604 has five significant figures

• Trailing zeroes may be significantmust specify significance by how the number is written

1300 nails - counted or weighed?

1.30000 –what was the precision of the measurement?

22

Use of Numbers

• Use scientific notation to remove doubtExpress answers as powers of 10 by moving the decimal place right (-) or left (+)

2000 � 2 x 103

15000 � 1.5 x 10?

0.004 � 4 x 10-3

0.000053 � __.__ x 10?

In scientific notation, zeros are given if they are significant

1.000 x 103 has 4 significant figures

2.40 x 103 has ? significant figures

23

Use of Numbers

• -Scientific notation for logarithms

take the log of 2.40 x 103

log(2.40 x 103) = 3.380

How many significant figures?

• -Imbedded zeroes are always significant

3.0604 has five significant figures

24

Use of Numbers

• Multiplication & Division rule

Easier of the two rules

Product has the smallest number of significant figures of multipliers

5.22 tooff round

21766.5

31.2x

224.4

25

Use of Numbers

• Multiplication & Division rule

Easier of the two rules

Product has the smallest number of significant figures of multipliers

5.22 tooff round

21766.5

31.2x

224.4

3.9 tooff round

89648.3

41.x

2783.2

26

Use of Numbers

• Addition & Subtraction rule

More subtle than the multiplication rule

Answer contains smallest decimal place of the addends.

6.95 tooff round

9463.6

20.2

423.1

3692.3

+

+

27

Use of Numbers

• Addition & Subtraction ruleMore subtle than the multiplication rule

Answer contains smallest decimal place of the addends.

6.95 tooff round

9463.6

20.2

423.1

3692.3

+

+

6.671 tooff round

6707.6

312.2

7793.8

−

28

Units of Measurement

Common Conversion Factors (Equalities)

• Length – 1 m = 39.37 inches

– 2.54 cm = 1 inch

• Volume– 1 liter = 1.06 qt

– 1 qt = 0.946 liter

• See Text and Handout for more conversion factors

29

Using Conversion factors

• Scientists often must convert between units.

• Conversion factors can be made for any relationship.

• Use known equivalence to make a fraction that can be used to “convert” from one unit to the other.

– Fraction equals 1, so you are multiplying by 1

30

Factor Label Method

• 1 inch = 2.54 cm

– Ratio

– Use ratio to perform a calculation Units will divide out.

• Convert 60 in to centimeters

in

cmor

cm

in

1

54.211

54.2

1==

cmin

cmxin 4.152

1

54.260 =

31

1 ft = 12 in becomes or

• Example - Express 9.32 yards in millimeters.

in 12

ft 1

ft 1

in 12

Factor Label Method

32

More practice

Express 9.32 yards in millimeters.

in 12

ft 1

ft 1

in 12

33

More Practice

)yd1

ft3( yd 9.32

mm ?yd 9.32 =

34

More practice

)ft1

in12( )

yd1

ft3( yd 9.32

mm ?yd 9.32 =

35

More Practice

)in1

cm2.54( )

ft1

in12( )

yd1

ft3( yd 9.32

mm ?yd 9.32 =

36

More Practice

mm 108.52)cm1

mm10( )

in1

cm2.54( )

ft1

in12( )

yd1

ft3( yd 9.32

mm ?yd 9.32

3×=

=

37

More Practice

mm 108.52)cm1

mm10( )

in1

cm2.54( )

ft1

in12( )

yd1

ft3( yd 9.32

mm ?yd 9.32

3×=

=

38

Factor Label Method

• Example - Express 627 milliliters in gallons.

39

The Unit Factor Method

• Example - Express 627 milliliters in gallons.

gal 0.166gal 0.166155gal ?

)qt4

gal1( )

L1

qt1.06( )

mL1000

L1( mL 627gal ?

mL 627gal ?

≈=

=

=

40

Practice

• Using your conversion handout

– Convert 25 g to lbs

– Convert 1 mL to Liters

– Convert 20 meters to cm

41

Area – length x width

� Area is two dimensional thus units must be in squared terms.

• Example 1-4: Express 2.61 x 104 cm2 in ft2.

• common mistake

)cm 2.54

in 1(cm102.61ft ? 242

×=

42

Area – length x width

� Area is two dimensional thus units must be in squared terms.

• Example 1-4: Express 2.61 x 104 cm2 in ft2.

2242 )cm 2.54

in 1(cm102.61ft ? ×=

43

Area – length x width

� Area is two dimensional thus units must be in squared terms.

• Example 1-4: Express 2.61 x 104 cm2 in ft2.

22242 )in 12

ft 1()

cm 2.54

in 1(cm102.61ft ? ×=

44

Volume – length x width x height

• Volume is three dimensional thus units must be in cubic terms—this volume is used in medical measurements--cc

• Example 1-5: Express 2.61 ft3 in cm3.

343

3333

cm 107.39cm 73906.9696

)in 1

cm 2.54()

ft 1

in 12(ft 2.61cm ?

×≈=

=

45

Percentage

• Percentage is the parts per hundred of a sample. = (g / g total ) X 100%

• Example - A 335 g sample of ore yields 29.5 g of iron. What is the percent of iron in the ore?

46

Percentage

• Percentage is the parts per hundred of a sample.

• Example - A 335 g sample of ore yields 29.5 g of iron. What is the percent of iron in the ore?

8.81%

100%x ore g 335

Feg 29.5

100%x ore of grams

iron of grams iron % ?

=

=

=

47

Density and Specific Gravity

• What is density?

• Density (how compact something is, mass per unit volume)

• density = mass/volume

48

Density and Specific Gravity

• Example - Calculate the density of a substance if 742 grams of it occupies 97.3 cm3.

Vm density

mL 3.97cm 97.3 mL 1 cm 1 33

=

=∴=

49

Density and Specific Gravity

• Example - Calculate the density of a substance if 742 grams of it occupies 97.3 cm3.

g/mL 7.63 density

mL 97.3g 742

density

Vm density

mL 3.97cm 97.3 mL 1 cm 1 33

=

=

=

=∴=

50

Density and Specific Gravity

• Example - You need 125 g of a corrosive liquid for a reaction. What volume do you need?

– liquid’s density = 1.32 g/mL

51

Density and Specific Gravity

• Example -

mL 94.7 1.32

g 125V

density

mV

V

mdensity

mLg

==

=∴=

52

Density and Specific Gravity

• Water’s density is essentially 1.00 g/mL at room T.

• Thus the specific gravity of a substance is very nearly equal to its density.

• Specific gravity has no units.

)water(density

)substance(densityGravity Specific =

53

Density and Specific Gravity

• Example1-9: A 31.10 gram piece of chromium is dipped into a graduated cylinder that contains 5.00 mL of water. The water level rises to 9.32 mL. What is the specific gravity of chromium?

mL 4.32

g 31.10 Cr ofdensity

mL 4.32

mL 5.00 - mL 9.32 Cr of Volume

=

=

=

54

Density and Specific Gravity

• Example1-9: A 31.10 gram piece of chromium is dipped into a graduated cylinder that contains 5.00 mL of water. The water level rises to 9.32 mL. What is the specific gravity of chromium?

20.7 1.00

7.20Cr ofGravity Specific

mLg 7.20

mLg 7.19907

mL 4.32

g 31.10 Cr ofdensity

mLg

mLg

==

≈=

=

55

Heat and Temperature

• Heat and Temperature are not the same thing

T is a measure of the intensity of heat in a body

• 3 common temperature scales - all use water as a reference

56

Heat and Temperature

MP water BP water

• Fahrenheit 32 oF 212 oF

• Celsius 0.0 oC 100 cC

• Kelvin 273 K 373 K

• Body temperature 37.0 oC, 98.6 oF– 37.2 oC and greater—sick

– 41 oC and greater, convulsions

– <28.5 oC hypothermia

57

Relationships of the Three Temperature Scales

273KC

or

273 C K

ipsRelationsh Centigrade andKelvin

o

o

−=

+=

58

1.8

32FC

or

32C 1.8F

1.85

9

10

18

100

180

ipsRelationsh Centigrade and Fahrenheit

oo

oo

−=

+×=

===

Example - Express 548 K in Celsius degrees.

•Example - Convert 211 oF to degrees Celsius.

59

Heat and Temperature

• Example - Convert 211oF to degrees Celsius.

C4.991.8

32112C

1.8

32FC

oo

oo

=−

=

−=

60

Heat and Temperature

• Example - Express 548 K in Celsius degrees.

275C

273485C

273KC

o

o

o

=

−=

−=