2. The Discovery Process Chemistry - The study of matter…
Matter - Anything that has mass and occupies space.
This desk
A piece of Aluminum foil
What about air?
Yes it is matter.

3. Chemistry - The study of matter and the changes it undergoes.
Chemical and physical changes
Energy changes

4. MAJOR AREAS OF CHEMISTRY
Biochemistry - the study of life at the molecular level
Organic Chemistry - the study of matter containing carbon and hydrogen.
Inorganic Chemistry - the study of matter containing all other elements.
Analytic Chemistry - analyze matter to determine identity and composition.

5. Physical Chemistry - attempts to explain the way matter behaves.
Not only does chemistry cover all the above fields but it’s fingers reach out into many other areas of study
other sciences
medical practitioners
pharmaceutical industry

6. Chemistry: A Science for the 21st Century
Health and Medicine
Sanitation systems
Surgery with anesthesia
Vaccines and antibiotics
Energy and the Environment
Fossil fuels
Solar energy
Nuclear energy

7. Chemistry: A Science for the 21st Century
Materials and Technology
Polymers, ceramics, liquid crystals
Room-temperature superconductors?
Molecular computing?
Food and Agriculture
Genetically modified crops
“Natural” pesticides
Specialized fertilizers

8. A main challenge of chemistry is to understand the connection between the macroscopic world that we experience and the microscopic world of atoms and molecules.
You must learn to think on the atomic level.

9. The Study of Chemistry
Macroscopic
Microscopic

10. Science Science is a framework for gaining and organizing knowledge.
Science is a plan of action — a procedure for processing and understanding certain types of information.
Scientists are always challenging our current beliefs about science, asking questions, and experimenting to gain new knowledge.
Scientific method is needed.

11. The scientific method is a systematic approach to research
A hypothesis is a tentative explanation for a set of observations
tested modified

12. Fundamental Steps of the Scientific Method
Process that lies at the center of scientific inquiry.

13. Force = mass x acceleration
A law is a concise statement of a relationship between phenomena that is always the same under the same conditions.
Force = mass x acceleration
A theory is a unifying principle that explains a body of facts and/or those laws that are based on them.
Atomic Theory

14. A possible explanation for an observation.
Scientific Models
Law
A summary of repeatable observed (measurable) behavior.
Hypothesis
A possible explanation for an observation.
Theory (Model)
Set of tested hypotheses that gives an overall explanation of some natural phenomenon.

15. Observation of a
phenomenon
A question
A hypothesis
(a potential answer)
Experimentation
Theory
New hypothesis
Further
experimentation
Development of
new
experimentation
and theory

16. Primordial Helium and the Big Bang Theory
Chemistry In Action:
Primordial Helium and the Big Bang Theory
In 1940 George Gamow hypothesized that the universe began with a gigantic explosion or big bang.
Experimental Support
expanding universe
cosmic background radiation
primordial helium

17. Chemistry is the study of matter and the changes it undergoes
Matter is anything that occupies space and has mass.
A substance is a form of matter that has a definite composition and distinct properties.
Sugar
Water
Gold

18. Atoms vs. Molecules Matter is composed of tiny particles called atoms.
Atom: smallest part of an element that is still that element.
Molecule: Two or more atoms joined and acting as a unit.

19. Oxygen and Hydrogen Molecules
Use subscripts when more than one atom is in the molecule.

20. A Chemical Reaction
One substance changes to another by reorganizing the way the atoms are attached to each other.

21. Heterogeneous Mixture
Mixtures
Have variable composition.
Homogeneous Mixture
Having visibly indistinguishable parts; solution.
Heterogeneous Mixture
Having visibly distinguishable parts.

22. A mixture is a combination of two or more substances in which the substances retain their distinct identities.
Homogenous mixture – composition of the mixture is the same throughout.
soft drink, milk, solder
Heterogeneous mixture – composition is not uniform throughout.
cement,
iron filings in sand

23. Which of the following is a homogeneous mixture?
CONCEPT CHECK!
Which of the following is a homogeneous mixture?
Pure water
Gasoline
Jar of jelly beans
Soil
Copper metal
gasoline

24. Physical Change
Change in the form of a substance, not in its chemical composition.
Example: boiling or freezing water
Can be used to separate a mixture into pure compounds, but it will not break compounds into elements.
Distillation
Filtration
Chromatography

25. Physical means can be used to separate a mixture into its pure components.
distillation
magnet

26. Chemical Change
A given substance becomes a new substance or substances with different properties and different composition.
Example: Bunsen burner (methane reacts with oxygen to form carbon dioxide and water)

27. Which of the following are examples of a chemical change?
CONCEPT CHECK!
Which of the following are examples of a chemical change?
Pulverizing (crushing) rock salt
Burning of wood
Dissolving of sugar in water
Melting a popsicle on a warm summer day
1 (burning of wood)

28. 114 elements have been identified
An element is a substance that cannot be separated into simpler substances by chemical means.
114 elements have been identified
82 elements occur naturally on Earth
gold, aluminum, lead, oxygen, carbon
32 elements have been created by scientists
technetium, americium, seaborgium

29. Elements: sulfur, arsenic, iodine, magnesium, bismuth, mercury.

31. A compound is a substance composed of atoms of two or more elements chemically united in fixed proportions.
Compounds can only be separated into their pure components (elements) by chemical means.
Water (H2O)
Glucose (C6H12O6)
Ammonia (NH3)

32. The Organization of Matter

33. Matter exists in five states:
Solid
Liquid
Gas
*Plasma
*Bose-Einstein condensates:

34. *Plasma
Plasma is the superheated phase where electrons get torn from the atom (ionization). It is the rarest phase of matter on Earth although it is the most common in the universe. They can be natural (lightning and the Sun) or man-made (fluorescent light tubes).

35. *Bose-Einstein condensates
- Examples of Bose-Einstein Condensates are superconductors and superfluids.
- These are materials that are cooled till they almost reach absolute zero (0 K, ­°C).
Superconductors are materials (solids) that have no resistance to the flow of electricity, such as superconducting metals, alloys and compounds (alloy of Niobium and Titanium).
Superfluids are fluids, such as liquid helium, that flows with little or no friction at temperatures close to absolute zero. At these temperatures the molecules exhibit strange quantum effects (strange behaviors).

36. The Common Three States of Matter
solid
liquid
gas

37. Matter and Properties Properties - characteristics of matter
chemical vs. physical
Three states of matter
1. gas - particles widely separated, no definite shape or volume solid
2. liquid - particles closer together, definite volume but no definite shape
3. solid - particles are very close together, define shape and definite volume

38. The Three States of Water

39. Physical or Chemical?
A physical change does not alter the composition or identity of a substance.
ice melting
sugar dissolving
in water
A chemical change alters the composition or identity of the substance(s) involved.
hydrogen burns in air to form water

40. hydrogen + oxygen  water
Chemical property - result in a change in composition and can be observed only through a chemical reaction.
Chemical reaction (chemical change) a process of rearranging, replacing, or adding atoms to produce new substances.
hydrogen + oxygen  water
reactants
products

41. Extensive and Intensive Properties
An extensive property of a material depends upon how much matter is is being considered.
mass
length
volume
An intensive property of a material does not depend upon how much matter is is being considered.
density
temperature
color

42. Nature of Measurement Measurement
Quantitative observation consisting of two parts.
number
scale (unit)
Examples
20 grams
6.63 × joule·second

43. Measurement in Chemistry
Data, Results and Units
Data - individual result of a single measurement or observation.
obtain the mass of a sample
record the temperature of
Results - the outcome of the experiment
Units - the basic quantity of mass, volume or whatever being measured.
A measurement is useless without its units.

44. ENGLISH AND METRIC UNITS
English system - a collection of measures accumulated throughout English history.
no systematic correlation between measurements.
1 gal = 4 quarts = 8 pints
Metric System - composed of a set of units that are related to each other decimally.
That is, by powers of tens

45. 1 meter = 10 decimeters = 100 centimeters
Truly systematic
1 meter = 10 decimeters = 100 centimeters
Basic Units of the Metric System
Mass gram g
Length meter m
volume liter L
prefixes are used to indicate the power of ten used

46. Matter - anything that occupies space and has mass.
mass – measure of the quantity of matter
SI unit of mass is the kilogram (kg)
1 kg = 1000 g = 1 x 103 g
weight – force that gravity exerts on an object
weight = c x mass
on earth, c = 1.0
on moon, c ~ 0.1
A 1 kg bar will weigh
1 kg on earth
0.1 kg on moon

47. Mass ≠ Weight
Mass is a measure of the resistance of an object to a change in its state of motion. Mass does not vary.
Weight is the force that gravity exerts on an object. Weight varies with the strength of the gravitational field.

48. International System of Units (SI)

50. Volume – SI derived unit for volume is cubic meter (m3)
1 cm3 = (1 x 10-2 m)3 = 1 x 10-6 m3
1 dm3 = (1 x 10-1 m)3 = 1 x 10-3 m3
1 L = 1000 mL = 1000 cm3 = 1 dm3
1 mL = 1 cm3

51. Density – SI derived unit for density is kg/m3
1 g/cm3 = 1 g/mL = 1000 kg/m3
density =
mass
volume
d = m V
A piece of platinum metal with a density of 21.5 g/cm3 has a volume of 4.49 cm3. What is its mass?
d = m V
m = d x V
= 21.5 g/cm3 x 4.49 cm3 = 96.5 g

52. Example #1
A certain mineral has a mass of 17.8 g and a volume of 2.35 cm3. What is the density of this mineral?

53. Example #2
What is the mass of a 49.6-mL sample of a liquid, which has a density of 0.85 g/mL?

54. Three Systems for Measuring Temperature
Fahrenheit
Celsius
Kelvin

55. K = 0C
273 K = 0 0C
373 K = 100 0C
0F = x 0C + 32 9 5
32 0F = 0 0C
212 0F = 100 0C

56. Conversions between Fahrenheit and Celsius
1. Convert 75oC to oF.
2. Convert -10oF to oC.
1. Ans. 167 oF 2. Ans. -23oC

57. Convert 172.9 0F to degrees Celsius.
0F = x 0C + 32 9 5
0F – 32 = x 0C 9 5
x (0F – 32) = 0C 9 5
0C = x (0F – 32) 9 5
0C = x (172.9 – 32) = 78.3 9 5

58. Chemistry In Action
On 9/23/99, $125,000,000 Mars Climate Orbiter entered Mar’s atmosphere 100 km (62 miles) lower than planned and was destroyed by heat.
1 lb = 1 N
1 lb = 4.45 N
“This is going to be the cautionary tale that will be embedded into introduction to the metric system in elementary school, high school, and college science courses till the end of time.”

59. Scientific Notation The number of atoms in 12 g of carbon:
602,200,000,000,000,000,000,000
6.022 x 1023
The mass of a single carbon atom in grams:
1.99 x 10-23
N x 10n
N is a number
between 1 and 10
n is a positive or
negative integer

60. Scientific Notation Addition or Subtraction 568.762 0.00000772
move decimal left
move decimal right
n > 0
n < 0
= x 102
= 7.72 x 10-6
Addition or Subtraction
Write each quantity with the same exponent n
Combine N1 and N2
The exponent, n, remains the same
4.31 x x 103 =
4.31 x x 104 =
4.70 x 104

61. Scientific Notation Multiplication Division
(4.0 x 10-5) x (7.0 x 103) =
(4.0 x 7.0) x (10-5+3) =
28 x 10-2 =
2.8 x 10-1
Multiply N1 and N2
Add exponents n1 and n2
Division
8.5 x 104 ÷ 5.0 x 109 =
(8.5 ÷ 5.0) x =
1.7 x 10-5
Divide N1 and N2
Subtract exponents n1 and n2

62. Scientific Notation
The measuring devise determines the number of significant figures a measurement has.
In this section you will learn
to determine the correct number of significant figures (sig figs) to record in a measurement
to count the number of sig figs in a recorded value
to determine the number of sig figs that should be retained in a calculation.

63. A digit that must be estimated in a measurement is called uncertain.
A measurement always has some degree of uncertainty. It is dependent on the precision of the measuring device.
Record the certain digits and the first uncertain digit (the estimated number).

64. Significant figures - all digits in a number representing data or results that are known with certainty plus one uncertain digit.

65. Measurement of Volume Using a Buret
The volume is read at the bottom of the liquid curve (meniscus).
Meniscus of the liquid occurs at about mL.
Certain digits: 20.15
Uncertain digit: 20.15

66. Significant Figures Any digit that is not zero is significant
1.234 kg significant figures
Zeros between nonzero digits are significant
606 m significant figures
Zeros to the left of the first nonzero digit are not significant
0.08 L significant figure
If a number is greater than 1, then all zeros to the right of the decimal point are significant
2.0 mg significant figures
If a number is less than 1, then only the zeros that are at the end and in the middle of the number are significant
g 3 significant figures

67. Significant Figures
Trailing zeros are zeros at the right end of the number. They are significant only if the number contains a decimal point.
9.300 has 4 sig figs.
150 has 2 sig figs.

68. How many significant figures are in each of the following measurements?
24 mL
2 significant figures
3001 g
4 significant figures m3
3 significant figures
6.4 x 104 molecules
2 significant figures
560 kg
2 significant figures

69. Significant Figures Addition or Subtraction
The answer cannot have more digits to the right of the decimal
point than any of the original numbers.
89.332
1.1 +
90.432
one significant figure after decimal point
round off to 90.4
3.70
0.7867
two significant figures after decimal point
round off to 0.79

70. Significant Figures Multiplication or Division
The number of significant figures in the result is set by the original number that has the smallest number of significant figures
4.51 x =
= 16.5
3 sig figs
round to
3 sig figs
6.8 ÷ =
= 0.061
2 sig figs
round to
2 sig figs

71. Significant Figures Exact Numbers
Numbers from definitions or numbers of objects are considered
to have an infinite number of significant figures
The average of three measured lengths; 6.64, 6.68 and 6.70? 3
= = 6.67
= 7
Because 3 is an exact number

72. RECOGNITION OF SIGNIFICANT FIGURES
All nonzero digits are significant.
The number of significant digits is independent of the position of the decimal point
Zeros located between nonzero digits are significant
4055 has 4 sig figs

73. Zeros to the left of the first nonzero integer are not significant.
Zeros at the end of a number (trailing zeros) are significant if the number contains a decimal point.
5.700
Trailing zeros are insignificant if the number does not contain a decimal point
versus 2000
Zeros to the left of the first nonzero integer are not significant.

74. How many significant figures are in the following? 3.400 3004 300.
Examples of Significant
Figures
How many significant figures are in the following?
3.400
3004
300.

75. Precision and Accuracy
Agreement of a particular value with the true value.
Precision
Degree of agreement among several measurements of the same quantity.

76. Accuracy – how close a measurement is to the true value
Precision – how close a set of measurements are to each other
accurate
&
precise
precise
but
not accurate
not accurate
&
not precise

77. You need to be able to convert between units
UNIT CONVERSION
You need to be able to convert between units
within the metric system
between the English and metric system
The method used for conversion is called the Factor-Label Method or Dimensional Analysis
!!!!!!!!!!! VERY IMPORTANT !!!!!!!!!!!

78. Use when converting a given result from one system of units to another.
To convert from one unit to another, use the equivalence statement that relates the two units.
Derive the appropriate unit factor by looking at the direction of the required change (to cancel the unwanted units).
Multiply the quantity to be converted by the unit factor to give the quantity with the desired units.

79. Dimensional Analysis Method of Solving Problems
Determine which unit conversion factor(s) are needed
Carry units through calculation
If all units cancel except for the desired unit(s), then the problem was solved correctly.
given quantity x conversion factor = desired quantity
desired unit
given unit
given unit x
= desired unit

80. What does any number divided by itself equal?
Let your units do the work for you by simply memorizing connections between units.
For example: How many donuts are in one dozen?
We say: “Twelve donuts are in a dozen.”
Or: 12 donuts = 1 dozen donuts
What does any number divided by itself equal?
ONE!
or...

81. What does any number times one equal?
This fraction is called a unit factor
What does any number times one equal?
That number.

82. a number divided by itself = 1
We use these two mathematical facts to do the factor label method
a number divided by itself = 1
any number times one gives that number back
Example: How many donuts are in 3.5 dozen?
You can probably do this in your head but let’s see how to do it using the Factor-Label Method.

83. Start with the given information...
3.5 dozen
= 42 donuts
Then set up your unit factor...
See that the units cancel...
Then multiply and divide all numbers...

84. The two unit factors are:
Example #1
A golfer putted a golf ball 6.8 ft across a green. How many inches does this represent?
To convert from one unit to another, use the equivalence statement that relates the two units.
1 ft = 12 in
The two unit factors are:

85. Example #1
A golfer putted a golf ball 6.8 ft across a green. How many inches does this represent?
Derive the appropriate unit factor by looking at the direction of the required change (to cancel the unwanted units).

86. Example #1
A golfer putted a golf ball 6.8 ft across a green. How many inches does this represent?
Multiply the quantity to be converted by the unit factor to give the quantity with the desired units.

87. Example #2
An iron sample has a mass of 4.50 lb. What is the mass of this sample in grams?
(1 kg = lbs; 1 kg = 1000 g)

88. Dimensional Analysis Method of Solving Problems
How many mL are in 1.63 L?
Conversion Unit 1 L = 1000 mL 1L
1000 mL
1.63 L x
= 1630 mL 1L
1000 mL
1.63 L x = L2 mL

89. The speed of sound in air is about 343 m/s
The speed of sound in air is about 343 m/s. What is this speed in miles per hour?
conversion units
meters to miles
seconds to hours
1 mi = 1609 m
1 min = 60 s
1 hour = 60 min
343 m s x
1 mi
1609 m
60 s
1 min x
60 min
1 hour x
= 767 mi
hour

90. C. Volume 1 gallon = 4 quarts 1 quart = 2 pints
Common Relationships Used in the English System
A. Weight 1 pound = 16 ounces
1 ton = 2000 pounds
B. Length foot = 12 inches
1 yard = 3 feet
1 mile = 5280 feet
C. Volume 1 gallon = 4 quarts
1 quart = 2 pints
1 quart = 32 fluid ounces
Commonly Used “Bridging” Units for Intersystem conversion
Quantity English Metric
Mass 1 pound = 454 grams
2.2 pounds = 1 kilogram
Length 1 inch = 2.54 centimeters
1 yard = meter
Volume 1 quart = liter
1 gallon = 3.78 liters

91. 1.3 Measurement in Chemistry
Examples of Unit Conversion
1. Convert 5.5 inches to millimeters
2. Convert 50.0 milliliters to pints
3. Convert 1.8 in2 to cm2
1.3 Measurement in Chemistry

92. Think Metric!!!!!!! 10E-12 boos = 1 picoboo 1boo E2 = 1 boo boo
10 E1 cards = 1decacards
10 E-2 menals = 1 centimental
10 E6 phones = 1 megaphone
10 E-6 phones = 1 microphone
10 E3 joys = 1 kilo-joy

93. Derived Units The density of an object is its mass per unit volume,
where d is the density, m is the mass, and V is the volume.
Generally the unit of mass is the gram.
The unit of volume is the mL for liquids; cm3 for solids; and L for gases. 2

94. A Density Example
A sample of the mineral galena (lead sulfide) weighs 12.4 g and has a volume of 1.64 cm3. What is the density of galena?

95. A Density Example
A sample of the mineral galena (lead sulfide) weighs 12.4 g and has a volume of 1.64 cm3. What is the density of galena?
mass
Density =
volume

96. A Density Example
A sample of the mineral galena (lead sulfide) weighs 12.4 g and has a volume of 1.64 cm3. What is the density of galena?
mass
12.4 g
Density = =
volume
1.64 cm3

97. A Density Example
A sample of the mineral galena (lead sulfide) weighs 12.4 g and has a volume of 1.64 cm3. What is the density of galena?
mass
12.4 g
Density = =
= = 7.56 g/cm3
volume
1.64 cm3

98. cork
water
brass nut
liquid mercury

99. Calculations using density:
What is the density of a sample of bone with mass of 12.0 grams and volume of 5.9 cm3?
A sinker of lead has a volume of 0.25 cm3. Calculate the mass in grams. The density of lead is 11.3 g/cm3.
What is the volume of air in liters (density = g/mL) occupied by 1.0 grams.

100. WORKED EXAMPLES

101. Worked Example 1.1

102. Worked Example 1.2

104. Worked Example 1.4