1. Chemistry: The Study of Change
Chapter 1

2. The Study of Chemistry
Macroscopic
Microscopic
1.2

3. Chemistry is the study of matter and the changes it undergoes.
Matter is anything that occupies space and has mass.
Sugar
Water
Gold
What about air?
Yes it is matter
1.4

4. MAJOR AREAS OF CHEMISTRY
Organic Chemistry - the study of matter which is carbon-based.
Inorganic Chemistry - the study of matter containing all other elements (Inorganic is everything else).
Analytical Chemistry - analyze matter to determine identity and composition (involves qualitative and quantitative) (eg. Alcohol in blood)
Biochemistry - the study of life at the molecular level (DNA sequencing)
Physical Chemistry - attempts to explain the way matter behaves (eg. Kinetics)

5. Chemistry Related Other Fields
Medical Science
Pharmaceutical Industry
Oil Industry
Paints and Coatings

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. 114 elements have been identified
An element is a substance that cannot be separated into simpler substances by chemical means (Listed in the Periodic Table)
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
Metals, nonmetals, metalloids

10. Classification of Matter
Based on Physical State:
Gas: does not have a fixed shape or a fixed volume. Both the volume & the shape of the gas is that of the container.
Liquid: Has a fixed volume but variable shape = shape of the container.
Solid: Has fixed shape and volume

11. The Three States of Matter (water in three forms)
solid
liquid
gas

12. Pure Substances Classification of Pure Substances
Elements: These only have 1 type of atoms.
2. Compounds: composed of atoms of two or more elements. eg. H2O chemically united in a fixed proportion – 2 H atoms and 1 O atom. It can be separated only chemically.
Ammonia (NH3) – 1N, 3H (combined at high pressure)
Elements exist in monoatomic or polyatomic form. For eg., Hydrogen and oxygen exist as H2 and O2 , that is diatomic molecules where as Ca exist as monoatomic.
Chemical Symbols: Cobalt – Co, Sodium – Na (natium-Latin)
CO is not cobalt, carbon monoxide

13. soft drink, milk, sugar solution
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, sugar solution
Heterogeneous mixture – composition is not uniform throughout.
cement,
iron filings in sand

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

15. Physical and Chemical Properties
Physical Properties: will keep the identity of the substance. Eg. boiling point (B.P), melting point (M.P), density, mass, volume, area.
(water changes the physical state by heating and cooling, not the composition, so the M.P and B.P are physical properties)
Chemical Properties: involves chemical change.
eg. rusting, burning of H in O H2O
Needs a chemical change to bring the original substances, not physical like boiling or melting.

16. Extensive and Intensive Properties
An extensive property of a material depends upon how much matter is being considered - additive
mass (M)
length
volume (V)
An intensive property of a material does not depend upon how much matter is is being considered. e.g., density of water g/mL at 25 C
density – M/V
color

17. Eg., Newton's Law of Gravity
Hypothesis, Theory, Law
A hypothesis is an educated guess, based on observation. Needs to prove by experiments to be true.
A scientific theory summarizes a hypothesis or group of hypotheses that have been supported with repeated testing.
A law generalizes a body of observations.
Eg., Newton's Law of Gravity
F = Gm1m2/r2

18. Classifications of Matter

19. Measurement in Chemistry
In SI units, not in English units.
Meaurement is always necessary to collect data
Mass, volume etc. with proper equipments or glassware
Always use unit. Mass = 5 for a salt is useless. Needs unit like g or kg.
A measurement is useless without its units.
weight – force that gravity exerts on an object

20. International System of Units (SI)
Revised Metric System

21. 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 often used

23. 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
In Chemistry, the smaller g is more convenient.
weight – force that gravity exerts on an object

24. 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

25. Density – SI derived unit for density is kg/m3
1 g/cm3 = 1 g/mL = 1000 kg/m3
Chemical Application: g/cm3 or g/mL
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

26. Derived Units in SI Area: (length x width); unit = m x m = m2
Volume: length x width x height; unit = m x m x m = m3
Density: mass/volume; unit = kg/m3
Speed: distance/time; unit = m/s
Acceleration: speed/time; unit = m/s/s = m/s2
Force: mass x acceleration; unit = kg m/s2 = N (newton)
Energy: force x distance; unit = kg m/s2 x m = kg m2/s2 = J (joule)
Pressure: force/Area; unit = kg m/s2 ÷ m2 = kg/ms2 = Pa (Pascal)

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

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

29. 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

30. 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 (read p.17, ch.1)
1 lb = 1 N
1 lb = 4.45 N
F=ma=4.45Newton
“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.”

31. Scientific Notation N x 10n 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

32. 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

33. 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

34. Scientific Notation
The measuring device 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.

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

36. 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

37. 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
5.6 x 102
2 significant figures

38. 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 digit after decimal point
round off to 90.4
3.70
0.7867
two digit after decimal point
round off to 0.79

39. 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

40. Significant Figures
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

41. 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

42. 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
ALWAYS NEEDED

43. 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

44. a number divided by itself = 1 Example: 2/2 = 1
We use these two mathematical facts to do the factor label method
a number divided by itself = 1
Example: 2/2 = 1
any number times one gives that number back
Example: 2 x 1 = 2
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.

45. 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...

46. 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

47. 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

48. 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

49. 1.3 Measurement in Chemistry
Examples of Unit Conversion
1. Convert 5.5 inches to millimeters
(139.7 mm)
2. Convert 50.0 milliliters to pints
( pint)
Convert 1.8 in2 to cm2
( cm2)
1.3 Measurement in Chemistry

50. 1.7

51. 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

52. 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?

53. 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

54. 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

55. 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

56. cork
water
brass nut
liquid mercury

57. 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.

58. WORKED EXAMPLES

59. Worked Example 1.1

60. Worked Example 1.2

62. Worked Example 1.4

67. Homework Problems
1.9,1.11,1.12,1.14,1.15,1.16,1.18,1.21,1.22,1.23,1.29,1.33,1.38,1.39,1.40,1.51,1.52,1.64,1.74, 1.80