1. Matters and Measurement
Edward Wen, PhD

2. Chemistry is about Everyday experience
Why Cookies tastes different from Cookie Dough?
Why Baking Powder or Baking Soda?
Why using Aluminum Foil, not Paper Towel?
What if the Temperature is set too high?
Photo credit: itsnicethat.com

3. Chapter Outline Classification of matters
Measurement, Metric system (SI)
Scientific Notation
Significant figures
Conversion factor
Density

4. In Your Room
Everything you can see, touch, smell or taste in your room is made of matter.
Chemists study the differences in matter and how that relates to the structure of matter.

5. What is Matter? Matter: anything that occupies space and has mass
Matter is actually composed of a lot of tiny little pieces: Atoms and Molecules

6. Atoms and Molecules
Atoms: the tiny particles that make up all matter. Helium gas (for blimp) is made up of Helium atoms.
Molecules: In most substances, the atoms are joined together in units. Liquid water is made up of water molecules (2 Hydrogen atoms + 1 Oxygen atoms)

7. Physical States of Matters
Matter can be classified as solid, liquid or gas based on what properties it exhibits

8. Why different States of a Matter? Structure Determines Properties
the atoms or molecules have different structures in solids, liquid and gases

9. Solids
Particles in a solid: packed close together and are fixed in position
though they may vibrate
Incompressible
retaining their shape and volume
Unable to flow

10. Liquids
Particles are closely packed, but they have some ability to move around
Incompressible
Able to flow, yet not to escape and expand to fill the container (not “antigravity”)

11. Gases
The particles have complete freedom from each other (not sticky to each other)
The particles are constantly flying around, bumping into each other and the container
There is a lot of empty space between the particles (low density)
 Compressible
 Able to flow and Fill space (“antigravity”)

12. Classifying Matter: Sugar, Copper, Coke, Gasoline/Water

13. Classification of Matter
Pure Substance
Constant Composition
Homogeneous
Mixture
Variable Composition
Matter

14. Pure substance Matter that is composed of only one kind of piece.
Solid: Salt, Sugar, Dry ice, Copper, Diamond
Liquid: Propane, distilled water (or Deionized water, DI water)
Gas: Helium gas (GOODYEAR blimp)

15. Classifying Pure Substances: Elements and Compounds
Elements: Substances which can not be broken down into simpler substances by chemical reactions. (A,B)
Compounds: Most substances are chemical combinations of elements. (C)
Examples: Pure sugar, pure water
can be broken down into elements
Properties of the compound not related to the properties of the elements that compose it

16. Elements Example: Diamond (pure carbon), helium gas.
116 known, 91 are found in nature
others are man-made
Abundance = percentage found in nature
Hydrogen: most abundant in the universe
Oxygen: most abundant element (by mass) on earth and in the human body
Silicon: abundant on earth surface
every sample of an element is made up of lots of identical atoms

17. Compounds Composed of elements in fixed percentages
water is 89% O & 11% H
billions of known compounds
Organic (sugar, glycerol) or inorganic (table salt)
same elements can form more than one different compound
water and hydrogen peroxide contain just hydrogen and oxygen
carbohydrates all contain just C, H & O (sugar, starch, glucose)

18. Mixture
Matter that is composed of different kinds of pieces. Different samples may have the same pieces in different percentages. (D)
Examples:
Solid: Flour, Brass (Copper and Zinc), Rock
Liquid: Salt water, soda, Gasoline
Gas: air

19. Classification of Mixtures
Homogeneous = composition is uniform throughout
appears to be one thing
every piece of a sample has identical properties, though another sample with the same components may have different properties
solutions (homogeneous mixtures): Air; Tap water
Heterogeneous = matter that is non-uniform throughout
contains regions with different properties than other regions: gasoline mixed with water; Italian salad dressing

20. What is a Measurement? Quantitative observation
comparison to an agreed upon standard
Every measurement has a number and a unit:
77 Fahrenheit: Room temperature
7.5 pounds: Average newborn body weight in the US:
55 ± 0.5 grams: amount of sugar in one can of Coca Cola
UNIT: what standard you are comparing your object to
the number tells you
what multiple of the standard the object measures
the uncertainty in the measurement (±)

21. Some Standard Units in the Metric System
Quantity Measured
Name of Unit
Abbreviation
Mass
gram g
Length
meter m
Volume
liter L
Time
seconds s
Temperature
Kelvin K

22. Related Units in the SI System
All units in the SI system are related to the standard unit by a power of 10 (exactly!)
1 kg = 103 g
1 km = 103 m
1 m = 102 cm
The power of 10 is indicated by a prefix
The prefixes are always the same, regardless of the standard unit

23. Prefixes Used to Modify Standard Unit
kilo = 1000 times base unit = 103
1 kg = 1000 g = 103 g
deci = 0.1 times the base unit = 10-1
1 dL = 0.1 L = 10-1 L; 1 L = 10 dL
centi = 0.01 times the base unit = 10-2
1 cm = 0.01 m = 10-2 m; 1 m = 100 cm
milli = times the base unit = 10-3
1 mg = g = 10-3 g; 1 g = 1000 mg
micro = 10-6 times the base unit
1 m = 10-6 m; 106 m = 1 m
nano = 10-9 times the base unit
1 nL = 10-9L; 109 nL = 1 L

24. Common Prefixes in the SI System
Symbol
Decimal
Equivalent
Power of 10
mega- M
1,000,000
Base  106
kilo- k
1,000
Base  103
deci- d
0.1
Base  10-1
centi- c
0.01
Base  10-2
milli- m
0.001
Base  10-3
micro-
m or mc
Base  10-6
nano- n
Base  10-9

25. Standard Unit vs. Prefixes
Using meter as example:
1 km = 1000 m = 103 m
1 g = 10 dm
= 100 cm = 102 cm
= 1000 mm = 103 mm
= 1,000,000 m = 106 m
= 1,000,000,000 nm = 109 nm

26. Length Two-dimensional distance an object covers
SI unit: METER (abbreviation as m)
About 3½ inches longer than a yard
1 m = 10-7 the distance from the North Pole to the Equator
Commonly use centimeters (cm)
1 m = 100 cm = yard
1 cm = 0.01 m = 10 mm
1 inch = 2.54 cm (exactly)

27. Mass Amount of matter present in an object SI unit: kilogram (kg)
about 2 lbs. 3 oz.
Commonly measure mass in grams (g) or milligrams (mg)
1 kg = pounds
(1 lbs. = )
1 g = 1000 mg = 103 mg
1 g = kg = 10-3 kg

28. Volume Amount of three-dimensional space occupied
SI unit = cubic meter (m3)
Commonly measure solid volume in cubic centimeters (cm3)
1 m3 = 106 cm3
1 cm3 = 10-6 m3 = m3
Commonly measure liquid or gas volume in milliliters (mL)
1 gallon (gal) = 3.78 L = 3.78  103 mL
1 L = 1 dm3 = 1000 mL = 103 mL
1 mL = 1 cm3 = 1 cc (cubic centimeter)

29. Common Everyday Units and Their EXACT Conversions
11 cm
1 inch (in) =
2.54 cm
1 mile
5280 feet (ft)
1 foot
12 in
1 yard
3 ft

30. Common Units and Their Equivalents
Mass
1 kilogram (km) =
2.205 pounds (lb)
1 pound (lb)
grams (g)
1 ounce (oz)
28.35 (g)
Volume
1 liter (L) =
1.057 quarts (qt)
1 U.S. gallon (gal)
3.785 liters (L)

31. Units Always write every number with its associated unit
Always include units in your calculations
you can do the same kind of operations on units as you can with numbers
cm × cm = cm2
cm + cm = cm
cm ÷ cm = 1
using units as a guide to problem solving

32. Conversion Factor
Relationships to Convert one unit of measurement to another: US dollar  Canadian dollar, dollar  cent
Conversion Factors: Relationships between two units
Both parts of the conversion factor have the same number of significant figures
Conversion factors generated from equivalence statements
e.g. 1 inch = 2.54 cm can give or 32

33. How to Use Conversion Factor
Arrange conversion factors so starting unit cancels
Arrange conversion factor so starting unit is on the bottom of the conversion factor
unit 2
unit 1
unit 1 x =
unit 2
Conversion Factor 33

34. We have been using the Conversion Factor ALL THE TIME! 
How are we converting #cents into #dollars? Why?
From 1 dollar = 100 cents
dollar
cents
1 dollar
100 cents
45,000 cents x =
450 dollars
Conversion Factor 34

35. Convert 325 mg to grams 0.325 g Given: 325 mg Find: ? g
Conv. Fact. 1 mg = 10-3 g Soln. Map: mg  g
0.325 g 35 35

36. Practice: How to set up Conversion?
To convert 5.00 inches to cm, from 1 in = 2.54 cm (exact), which one of the two conversion factors should be used? or

37. Practice: Conversion among Units
500 mg = ? g
3.78 L = ? mL
1.2 nm = ? m
* 8.0 in = ? m

38. Scientific Notation Very Large vs. Very Small numbers:
The sun’s diameter is 1,392,000,000 m;
An atom’s diameter is m
Scientific Notation: x 109 m & 3 x m
the sun’s
diameter is
1,392,000,000 m
an atom’s
average diameter is m

39. Scientific Notation (SN)
Power of 10 (Math language):
10 x 10 = 100  100 = 102 (2nd power of 10)
10 x 10 x 10 = 1,000  1,000 = 103 (3rd power of 10)
each Decimal Place in our number system represents a different power of 10
24 = 2.4 x 101 = 2.4 x 10
1,000,000,000 (1 billion) = 109
(1/10 billionth ) = 10-10
Easily comparable by looking at the power of 10

40. Exponents 10Y
when the exponent on 10 (Y) is positive, the number is that many powers of 10 larger
sun’s diameter = x 109 m = 1,392,000,000 m
when Y is negative, the number is that many powers of 10 smaller
avg. atom’s diameter = 3 x m = m
1.23 x
10-8
decimal part
exponent part
exponent
1.23 x 105 > 4.56 x 102
4.56 x 10-2 > 7.89 x 10-5
7.89 x 1010 > 1.23 x 1010

41. Writing Numbers in SN 1.234 x 107 2.34 x 10-5 Big numbers: 12,340,000
Small numbers:
1.234 x 107
2.34 x 10-5

42. Writing a Number in Standard Form
1.234 x 10-6
since exponent is -6, move the decimal point to the left 6 places
if you run out of digits, add zeros
If the exponent > 1, add trailing zeros:
1.234 x 1010
12,340,000,000

43. Scientific calculators

44. Inputting Scientific Notation into a Calculator
-1.23 x 10-3
input decimal part of the number
if negative press +/- key
(–) on some
press EXP key
EE on some (maybe 2nd function)
input exponent on 10
press +/- key to change exponent to negative
Press
+/-
Input 1.23
1.23
EXP
Input 3
-1.23

45. Significant Figures (Sig. Fig.)
Definition: The non-place-holding digits in a reported measurement
some zero’s in a written number are only there to help you locate the decimal point
What is Sig. Fig. for?
the range of values to expect for repeated measurements
the more significant figures there are in a measurement, the smaller the range of values is
12.3 cm
has 3 sig. figs.
and its range is
12.2 to 12.4 cm cm
has 4 sig. figs.
and its range is
to cm

46. Counting Significant Figures
All non-zero digits are significant
1.5 : 2 Sig. Fig.s
Interior zeros are significant
1.05 : 3 Sig. Fig.s
Trailing zeros after a decimal point are significant
1.050 : 4 Sig. Fig.s. Leading zeros are NOT significant
: 4 Sig. Fig.s Place-holding zero’s
= SN : x 10-3

47. Counting Significant Figures (Contd)
4. Exact numbers has infinite () number of significant figures:
example:
1 pound = 16 ounces
1 kilogram = 1,000 grams = 1,000,000 milligrams
1 water molecule contains 2 hydrogen atoms
5. Zeros at the end of a number without a written decimal point are ambiguous and should be avoided by using scientific notation.
Example: 150. has 3 sig. fig
150 is ambiguous number
1.50 x 102 has 3 sig. fig.

48. Example–Counting Sig. Fig. in a Number
How many significant figures are in each of the following numbers?
0.0035
1.080 27
2.97 × 105
1 m = 1000 mm
2 Sig. Fig. – leading zeros not sig.
4 Sig. Figs – trailing & interior zeros sig.
2 sig. Figs, all digits sig.
3 Sig. Figs – only decimal parts count sig.
both 1 and 1000 are exact numbers. unlimited sig. figs.

49. Practice: How many Significant figures vs. Decimal places?
2.2 cm
2.50 cm
2 sig. Figs;
1 decimal place
3 sig. Figs;
2 decimal places

50. Sig. Fig. in Multiplication/Division; Rounding vs. Zeroing
When multiplying or dividing measurements with Sig. Fig., the result has the same number of significant figures as the measurement with the fewest number of significant figures
Rounding
× 89, × =
÷ 6.10 =
= 45
3 SF
5 SF
2 SF
2 SF
= 0.966
4 SF
3 SF
3 SF

51. Sig. Fig. in Multiplication/Division: Scientific notation
Occasionally, scientific notation is needed to present results with proper significant figures.
5.89 × 6,103 = = 3.59 × 104

52. Example: Determine the Correct Number of Sig. Fig.
1.01 × 0.12 × ÷ 96 =
56.55 × ÷ = 1.5
= 0.068
3 SF
2 SF
4 SF
2 SF
result should
have 2 Sig. Fig.
= 1.50
4 SF.
3 SF.
6 SF.
result should
have 3 Sig. Fig.

53. Sig. Fig. in Addition/Subtraction
when adding or subtracting measurements with significant figures, the result has the same number of decimal places as the measurement with the fewest number of decimal places
= 9.214
= 0.9
= 9.21
2 dp
3 dp
3 dp
2 dp
= 0.900
3 dp
3 dp
3 dp

54. Example: Determine the Correct Number of Significant Figures
– 1.22 =
0.764 – – = -8
= 124.9
3 dp
1 dp
2 dp
result should
have 1 dp =
3 dp
3 dp
3 dp
result should
have 3 dp

55. Sig. Fig. in Combined Calculations
Do  and/or , then + and/or -
– × 2.3
3 dp Sig. Fig Sig. Fig.
= – = = -10
3 dp dp dp (2 sig. fig.)
Parentheses (): Do calculation in () first, then the rest
× ( – )
2 dp dp
= × = = 12
4 Sig. Fig Sig. Fig. 2 Sig. Fig.

56. Practice: Calculation with Proper Significant Figures
a x =
b x =
= 17.00
7.0 – = 6.0
15.00/3.75 = 4.00
6.7  8.8 = 59

57. How to solve Unit Conversion Problems
Write down Given Amount and Unit
Write down what you want to Find and Unit
Write down needed Conversion Factors or Equations
Design a Solution Map for the Problem
order Conversions to cancel previous units or
arrange Equation so Find amount is isolated. Example: from Equation A = b  c to solve for b

58. Solution Map for Unit Conversion
Apply the Steps in the Solution Map
check that units cancel properly
multiply terms across the top and divide by each bottom term
Example:
5) Check the Answer to see if its Reasonable
correct size and unit 58

59. Example: Unit Conversion

60. Alternative Route: Convert 7.8 km to miles
Given: 7.8 km
Find: ? mi
Conv. Fact. 1 mi = 5280 ft
1 foot = 12 in 1 in = 2.54 cm (exact)
Soln. Map: km  mi km m cm in ft mi 60 60

61. Alternative Route: Convert 7.8 km to miles
Given: 7.8 km
Find: ? mi
Conv. Fact. 1 mi = 5280 ft
1 foot = 12 in 1 in = 2.54 cm (exact)
Soln. Map: km  mi
Apply the Solution Map:
= mi
Sig. Figs. & Round:
= 4.8 mi 61 61

62. Temperature
Temperature is a measure of the average kinetic energy of the molecules in a sample
Not all molecules have in a sample the same amount of kinetic energy
a higher temperature means a larger average kinetic energy

63. Fahrenheit Temperature Scale
Two reference points:
Freezing point of concentrated saltwater (0°F)
Average body temperature (100°F)
more accurate measure now set average body temperature at 98.6°F
Room temperature is about 75°F

64. Celsius Temperature Scale
Two reference points:
Freezing point of distilled water (0°C)
Boiling point of distilled water (100°C)
more reproducible standards
most commonly used in science
Room temperature is about 25°C

65. Fahrenheit vs. Celsius
a Celsius degree is 1.8 times larger than a Fahrenheit degree
the standard used for 0° on the Fahrenheit scale is a lower temperature than the standard used for 0° on the Celsius scale

66. The Kelvin Temperature Scale
both the Celsius and Fahrenheit scales have “-” numbers
Kelvin scale is an absolute scale, meaning it measures the actual temperature of an object
0 K is called Absolute Zero: all molecular motion would stop, theoretically the lowest temperature in the universe
0 K = -273°C = -459°F
Absolute Zero is a theoretical value

67. Kelvin vs. Celsius
the size of a “degree” on the Kelvin scale is the same as on the Celsius scale
though technically, call the divisions on the Kelvin as kelvins, not degrees
that makes 1 K 1.8 times larger than 1°F
the 0 standard on the Kelvin scale is a much lower temperature than on the Celsius scale

68. Extremes of Temperature
On the Earth,
Lowest temperature recorded: -89.2°C ( °F, 184 K)
Highest air temperature recorded: ~60°C (140 F)
In science lab,
the highest temperature: 4 x 1012 K (?)
the lowest temperature: ~10-10 K (?)

69. Conversion Between Fahrenheit and Kelvin Temperature Scales

70. Convert 104°F into Celsius and Kelvin
Information
Given: 104 F
Find: ? °C, ? K
Eq’ns:
Fahrenheit to Celsius:
= 40 °C (keep 2 significant figures)
Celsius to Kelvin:
= 313 K

71. Mass & Volume Mass & Volume: two main characteristics of matter
even though mass and volume are individual properties - for a given type of matter they are related to each other!
 Density (ratio of mass vs. volume): for a certain matter, its density is one of the characteristic to distinguish from one another

72. Unit for density
Solids = g/cm3
1 cm3 = 1 mL
Liquids = g/mL: Density of water = 1.00 g/mL
Gases = g/L: Density of Air ~ 1.3 g/L
Volume of a solid can be determined by water displacement
Density : solids > liquids >>> gases
except ice and dry wood are less dense than liquid water!

73. Density of Common Matters

74. Density
Temperature affects the density: Heating objects causes objects to expand, density
The Lava Lamp: heating/cooling
In a heterogeneous mixture, the denser object sinks
Why do hot air balloons rise?
The “Gold Rush”: Extracting gold particle from sand
Density of gasoline changes over the day!

75. Density and Volume Styrofoam vs. Quarter:
Both of these items have a mass of 23 grams, but they have very different volumes; therefore, their densities are different as well.

76. Density and Buoyancy Average density of human body = 1.0 g/cm3
Average density of sea water = 1.03 g/mL
Density of mercury, liquid metal, = 13.6 g/mL
Density of copper penny = 8.9 g/cm3

77. Density of Body and Body Fat
Density of fat tissue < Density of Muscle/Bones
Estimate the mass percentage of body fat:
Average body fat%:
Female 28%, Male 22%

78. Using Density in Calculations
Solution Maps:
Both sides multiplied by Volume 
Both sides divided by Density
m, V D
V, D m
m, D V

79. Application of Density
A man gives a woman an engagement ring and tells her that it is made of platinum (Pt). Critical thinking : test to determine the ring’s density before giving him an answer about marriage. Data: She places the ring on a balance and finds it has a mass of 5.84 grams. She then finds that the ring displaces cm3 of water. Density Pt = 21.4 g/cm3

80. Test results Given: Mass = 5.84 grams Volume = 0.556 cm3
Density Pt = 21.4 g/cm3
Find: Density in grams/cm3

81. Density as a Conversion Factor
Between mass and volume!!
Density H2O = 1 g/mL \
1 g H2O  1 mL H2O
Density lead = 11.3 g/cm3
11.3 g lead  1 cm3 lead
How much does 4.0 cm3 of Lead weigh? =
4.0 cm3 Pb
11.3 g Pb
1 cm3 Pb
45 g Pb x

82. Measurement and Problem Solving Density as a Conversion Factor
The gasoline in an automobile gas tank has a mass of 60.0 kg and a density of g/cm3. What is the volume?
Given: kg
Find: Volume in L
Conversion Factors:
0.752 grams/cm3
1000 grams = 1 kg

83. Measurement and Problem Solving Density as a Conversion Factor
kg  g  cm3

84. Example: A 55. 9 kg person displaces 57
Example: A 55.9 kg person displaces 57.2 L of water when submerged in a water tank. What is the density of the person in g/cm3?
Information:
Given: m = 5.59 x 104 g
Find: density, g/cm3
Solution Map: m,VD
Equation:
Volume = 57.2 L = 5.72 x 104 cm3
= g/cm3

85. Practice: Calculation involving Density
The density of air at room temperature and sea level is 1.29 g/L. Calculate the mass of air in a 5.0-gal bottle (1 gal = 3.78 L).
A driver filled kg of gasoline into his car. If the density of gasoline is g/mL, what is the volume of gasoline in liters?
KEY: 24 g (2SF)
KEY: 19.8 L (3SF)

86. About Challenging Problems
1.99+: Proper dosage of a drug is 3.5 mg/kg of body weight. Calculate the milligrams of this drug for a 138-lb individual? (1 lb = 454 g).
1.103: 100. mg ibuprofen/5 mL Motrin. Calculate the grams of ibuprofen in 1.5 teaspoons of Motrin. (1 teaspoon = 5.0 mL)
KEY: 2.2×102 mg (2SF)
KEY: 0.15 g (2SF)