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Chemistry:
The Study of Change
Chapter 1
The Study of Chemistry
Macroscopic
Microscopic
1.2
• Chemistry is the study of matter and the
changes it undergoes.
•
Matter is anything that occupies space
and has mass.
Water
Sugar
Gold
What about air?
Yes it is matter
1.4
MAJOR AREAS OF CHEMISTRY
• Organic Chemistry - the study of matter which is carbonbased.
• 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.
Determination of 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)
Chemistry Related Other Fields
•
•
•
•
Medical Science
Pharmaceutical Industry
Oil Industry
Paints and Coatings
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
Chemistry: A Science for the 21st Century
• Materials and Technology
• Polymers, ceramics
• Food and Agriculture
• Genetically modified crops
• Specialized fertilizers
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
Latin names: aurum (Au), ferrum (Fe),natrium
(Na)
Classification of Matter
1. Based on Physical State:
a. Gas: does not have a fixed shape or a fixed volume. Both the
volume & the shape of the gas is that of the container.
b. Liquid: Has a fixed volume but variable shape = shape of the
container.
c. Solid: Has fixed shape and volume
The Three States of Matter (water in three forms)
gas
liquid
solid
Pure Substances
Classification of Pure Substances
1. 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
A mixture is a combination of two or more substances
in which the substances retain their distinct identities.
1. Homogenous mixture – composition of the
mixture is the same throughout.
soft drink, milk, sugar solution
2. Heterogeneous mixture – composition is not
uniform throughout.
cement,
iron filings in sand
Physical means can be used to separate a mixture
into its pure components.
distillation
magnet
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
H 2O
Needs a chemical change to bring the original substances, not
physical like boiling or melting.
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 0.99713g/mL at
25 C
• density – M/V
• color
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
Classifications of Matter
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
International System of Units (SI)
Revised Metric System
• Truly systematic
- 1 meter = 10 decimeters = 100
centimeters
Basic Units of the Metric System
Mass
Length
volume
gram
meter
liter
g
m
L
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
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
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
mass
m
density =
d= V
volume
A piece of platinum metal with a density of 21.5
g/cm3 has a volume of 4.49 cm3. What is its mass?
m
d= V
m = d x V = 21.5 g/cm3 x 4.49 cm3 = 96.5 g
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)
8. Pressure: force/Area; unit = kg m/s2 ÷ m2 = kg/ms2 = Pa
(Pascal)
1.
2.
3.
4.
5.
6.
7.
K = 0C + 273.15
273 K = 0 0C
373 K = 100 0C
0F
= 9 x 0C + 32
5
32 0F = 0 0C
212 0F = 100 0C
Conversions between Fahrenheit and Celsius
o
o
o
F - 32
C
1.8
F  1.8 ( C)  32
o
1. Convert 75oC to oF.
2. Convert -10oF to oC.
1. Ans. 167 oF
2. Ans. -23oC
Convert 172.9 0F to degrees Celsius.
9
=
x 0C + 32
5
0F – 32 = 9 x 0C
5
0F
5 x (0F – 32) = 0C
9
0C = 5 x (0F – 32)
9
0C = 5 x (172.9 – 32) = 78.3
9
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.”
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:
0.0000000000000000000000199
1.99 x 10-23
N x 10n
N is a number
between 1 and 10
n is a positive or
negative integer
Scientific Notation
568.762
0.00000772
move decimal left
move decimal right
n>0
n<0
568.762 = 5.68762 x 102
0.00000772 = 7.72 x 10-6
Addition or Subtraction
1. Write each quantity with
the same exponent n
2. Combine N1 and N2
3. The exponent, n, remains
the same
4.31 x 104 + 3.9 x 103 =
4.31 x 104 + 0.39 x 104 =
4.70 x 104
Scientific Notation
Multiplication
1. Multiply N1 and N2
2. Add exponents n1 and n2
Division
1. Divide N1 and N2
2. Subtract exponents n1 and n2
(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
8.5 x 104 ÷ 5.0 x 109 =
(8.5 ÷ 5.0) x 104-9 =
1.7 x 10-5
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.
Significant Figures
Figure TA 1.2
Significant figures - all digits in a number
representing data or results that are known with
certainty plus one uncertain digit.
Copyr ight © 2001 T he McGr aw- Hill Companies, Inc. Permission required for reproduction or display.
Significant Figures
• Any digit that is not zero is significant
1.234 kg
4 significant figures
• Zeros between nonzero digits are significant
606 m
3 significant figures
• Zeros to the left of the first nonzero digit are not significant
0.08 L
1 significant figure
• If a number is greater than 1, then all zeros to the right of the
decimal point are significant
2.0 mg
2 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
0.00420 g
3 significant figures
How many significant figures are in
each of the following measurements?
24 mL
2 significant figures
3001 g
4 significant figures
0.0320 m3
3 significant figures
6.4 x 104 molecules
2 significant figures
560 kg (5.6 x 102 )
2 significant figures
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
3.70
-2.9133
0.7867
one digit after decimal point
round off to 90.4
two digit after decimal point
round off to 0.79
Scientific Notation
568.762
0.00000772
move decimal left
move decimal right
n>0
n<0
568.762 = 5.68762 x 102
0.00000772 = 7.72 x 10-6
Addition or Subtraction
1. Write each quantity with
the same exponent n
2. Combine N1 and N2
3. The exponent, n, remains
the same
4.31 x 104 + 3.9 x 103 =
4.31 x 104 + 0.39 x 104 =
4.70 x 104
1.8
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 3.6666 = 16.536366 = 16.5
3 sig figs
round to
3 sig figs
6.8 ÷ 112.04 = 0.0606926 = 0.061
2 sig figs
round to
2 sig figs
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?
6.64 + 6.68 + 6.70
= 6.67333 = 6.67 = 7
3
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
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 FactorLabel Method or Dimensional Analysis
ALWAYS NEEDED
Dimensional Analysis Method of Solving Problems
1. Determine which unit conversion factor(s) are needed
2. Carry units through calculation
3. If all units cancel except for the desired unit(s), then the
problem was solved correctly.
given quantity x conversion factor = desired quantity
given unit x
desired unit
given unit
= desired unit
• 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.
Start with the given information...
3.5 dozen
12 donuts

1 dozen
= 42 donuts
Then set up your unit factor...
See that the units cancel...
Then multiply and divide all numbers...
Dimensional Analysis Method of Solving Problems
How many mL are in 1.63 L?
Conversion Unit 1 L = 1000 mL
1000 mL
1.63 L x
= 1630 mL
1L
2
1L
L
1.63 L x
= 0.001630
1000 mL
mL
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 mi
60 s
m
x
x
343
s 1609 m
1 min
1 hour = 60 min
60 min
mi
x
= 767
hour
1 hour
Common Relationships Used in the English System
A. Weight
1 pound = 16 ounces
1 ton = 2000 pounds
1 foot = 12 inches
1 yard = 3 feet
1 mile = 5280 feet
1 gallon = 4 quarts
1 quart = 2 pints
1 quart = 32 fluid ounces
B. Length
C. Volume
Commonly Used “Bridging” Units for Intersystem conversion
Quantity
Mass
Length
Volume
English
1 pound
2.2 pounds
1 inch
=
1 yard
1 quart
1 gallon
Metric
=
454 grams
=
1 kilogram
2.54 centimeters
=
0.91 meter
=
0.946 liter
=
3.78 liters
1.3 Measurement in
Chemistry
1. Convert 5.5 inches to millimeters
(139.7 mm)
2. Convert 50.0 milliliters to pints
(0.1057 pint)
3. Convert 1.8 in2 to cm2
(11.613 cm2)
1.7
Derived Units
• The density of an object is its mass per
unit volume,
m
d
V
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.
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?
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
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
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
3
Density =
=
=
7.5609
=
7.56
g/cm
volume
1.64 cm3
cork
water
brass nut
liquid mercury
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 =
0.00129 g/mL) occupied by 1.0 grams.
WORKED
EXAMPLES
Worked Example 1.1
Worked Example 1.2
Worked Example 1.4
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