Download AP Chemistry Chapter 1: Chemical Foundations

AP Chemistry
Chapter 1: Chemical Foundations
The only thing that matters is Matter!
The Scientific Method
1. Observations
(collecting data)
-quantitative or qualitative
2. Formulating hypothesis
- possible explanation for the
observation
3. Performing experiments
- gathering new information to
decide whether the hypothesis
is valid
Outcomes Over the Long-Term
Theory (Model)
- A set of tested hypotheses that give an overall explanation
of some natural phenomenon.
Natural Law
- The same observation applies to many different systems
- ex: Law of conservation of mass
Law vs. Theory
A natural law summarizes what happens
Law of gravity
A theory (model) is an attempt to explain why it happens.
Einstein's theory of gravity
describes gravitational forces in
terms of the curvature of spacetime
caused by the presence of mass
Nature of Measurement
• A quantitative observation is a
measurement.
• A measurement consists of both a number
and a unit.
Examples:
20 grams
6.63 x 10-34 J·s
International System of Units
 The International System of Units, symbolized SI, is
the modern version of the metric system
 Seven SI Units:
Quantity
Unit
Symbol
length
meter
m
mass
kilogram
kg
time
second
s
electric current
ampere
A
temperature
kelvin
K
amount of substance
mole
mol
luminous intensity
candela
cd
SI Units
SI Prefixes Common in Chemistry
1 Liter = 1dm3 = (10 cm)3 = 1000cm3
Since 1 cm3 = 1 mL,
1L= 1000ml
Precision and Accuracy
Accuracy refers to the agreement of a particular value
with the true value.
Precision refers to the degree of agreement among
several measurements made in the same manner.
Neither
accurate nor
precise
Precise but not
accurate
Precise AND
accurate
Types of Error
Random Error (Indeterminate Error)
• measurement has an equal probability of being
high or low.
Systematic Error (Determinate Error)
• Occurs in the same direction each time (high or
low), often resulting from poor technique or
incorrect calibration. This can result in
measurements that are precise, but not accurate.
Uncertainty in Measurement
• The volume of a buret is read
at the bottom of the liquid
curve (meniscus).
• Meniscus of the liquid occurs
at about 20.15 mL.
 Certain digits: 20.15
 Uncertain digit: 20.15
Uncertainty in Measurement
A digit that must be estimated is called
uncertain. A measurement always has
some degree of uncertainty.
 Measurements are performed with
instruments
 No instrument can read to an infinite
number of decimal places
Significant Figures
Significant Figures: all the digits that can
be known precisely in a measurement,
plus one last estimated (uncertain) digit
To determine if a figure is significant, you
need to follow the rules!
4
5
9
2
7
9
8
4
0
3
1. All non-zero integers are always significant
341 = 3 Sig Figs
2. All “trapped” zeros are always significant
7003 = 4 Sig Figs
3. “Leading” zeros are NEVER significant.
0.0071 = 2 Sig Figs
4. “Trailing” zeros are ONLY significant when
there is a DECIMAL in the number.
43.00 = 4 Sig Figs
5. Zeros that are “placeholders” at the end of
a number are NOT significant.
300 = 1 Sig. Figs
6. Unlimited number of sig. figs:
1. Counted objects
Ex: 36 students in the class.
2. Defined or Exact quantities
Ex: 60 minutes = 1 hour
Sig Fig Practice #1
How many significant figures in each of the following?
1.0070 m 
5 sig figs
17.10 kg 
4 sig figs
100,890 L 
5 sig figs
3.29 x 103 s 
3 sig figs
0.0054 cm 
2 sig figs
3,200,000 mol2 sig figs
Significant Figures in Mathematical Operations
Multiplication and Division:
The answer is rounded to the same number of
significant figures as the measurement with the least
significant figures.
6.38 x 2.0 = 12.76  13
(2.0 ONLY has 2 sig figs; therefore, round
final answer to 2 sig figs)
Sig Fig Practice
Calculation
Calculator says:
Answer
3.24 m x 7.0 m
22.68 m2
100.0 g ÷ 23.7 cm3
4.219409283 g/cm3 4.22 g/cm3
23 m2
0.02 cm x 2.371 cm 0.04742 cm2
0.05 cm2
710 m ÷ 3.0 s
236.6666667 m/s
240 m/s
1818.2 lb x 3.23 ft
5872.786 lb·ft
5870 lb·ft
1.030 g ÷ 2.87 mL
2.9561 g/mL
2.96 g/mL
Rules for Significant Figures in
Mathematical Operations
Addition and Subtraction:
Display the final answer with the same number of
decimal places as the least precise measurement
used in the calculation.
6.8 + 11.934 = 18.734  18.7
(6.8 goes to the tenth; therefore, round to the
tenth: 3 sig figs)
Calculation
Sig Fig Practice #3
Calculator says:
Answer
3.24 m + 7.0 m
10.24 m
10.2 m
100.0 g - 23.73 g
76.27 g
76.3 g
0.02 cm + 2.371 cm
2.391 cm
2.39 cm
713.1 L - 3.872 L
709.228 L
709.2 L
1818.2 lb + 3.37 lb
1821.57 lb
1821.6 lb
2.030 mL - 1.870 mL
0.16 mL
0.160 mL
Concept Check
You have water in each graduated
cylinder shown. You then add both
samples to a beaker (assume that
all of the liquid is transferred).
How would you write the number
describing the total volume?
2.8 + 0.28 = 3.08 = 3.1 mL
What limits the precision of the
total volume?
Units of Temperature:
Kelvin, Celsius, Fahrenheit
The lowest
theoretical
temperature
possible where all
motion of
particles stop is
O K or absolute
zero.
K = oC + 273.15
Density
• Density: The ratio of the mass of an object to
its volume.
• Common units are g/cm3 or g/mL
• Density =
Mass
Volume
Example: A student determines that a piece of
metal has a volume of 285 mL and a mass of
612 g. Is the shiny piece of metal Aluminum,
which has a density of 2.70 g/mL?
NO
D=
m
v
=
612 g
285 mL
= 2.147
= 2.15 g/mL
Water has a density of 1.00 g/cm3 at 40C
• Materials that have a density lower than 1 g/cm3
will float in water.
• Materials that have a density greater than 1 g/cm3
will sink in water
Density generally decreases as its
temperature increases.
• Water is an exception
• Ice has a density of .917
g/cm3 at 00C. Water has
a density of 1.00 g/cm3.
That’s why ice floats!
DA
• Covered while reviewing summer worksheets.
Square and Cubic units
• Use the conversion factors you already know,
but when you square or cube the unit, don’t
forget to cube the number also!
• Best way: Square or cube the ENITRE
conversion factor
• Example: Convert 4.3 cm3 to mm3
3
4.3 cm
(
10 mm
1cm
3
)
3
=
4.3 cm
3
10 mm
3
3
= 4300 mm
3
3
1 cm
•
•
Classification of Matter
Anything occupying space and having mass.
Matter exists in three main states.
 Solid – definite volume & shape
 Liquid – definite volume, indefinite shape
 Gas – indefinite volume & shape
Properties of Matter
Extensive properties depend on the amount
of matter that is present.
Volume
Mass
Energy Content (think Calories!)
Intensive properties do not depend on the
amount of matter present.
Melting point
Boiling point
Density
Kinetic Nature of Matter
Matter
consists of
atoms and
molecules
in motion.
Structure of a Solid
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Structure of a Liquid
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Structure of a Gas
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OTHER STATES OF MATTER
• PLASMA — an electrically charged gas;
Example: the sun or any other star
• BOSE-EINSTEIN CONDENSATE — a
condensate that forms near absolute zero
that has superconductive properties;
Example: supercooled Rb gas
Mixtures
• Have variable composition.
Homogeneous Mixture
 Having visibly indistinguishable
parts; solution.
Heterogeneous Mixture
 Having visibly distinguishable
parts.
Homogeneous Mixtures
Homogeneous vs. Heterogeneous
Mixtures
Compound vs. Mixture
Concept Check
Which of the following is a homogeneous mixture?





Pure water
Gasoline
Jar of jelly beans
Soil
Copper metal
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Physical Change
•
Change in the form of a substance, not in its
chemical composition.
• A physical change will not break up compounds
Example:
 boiling or freezing water
 Distillation
 Filtration
 Chromatography
Separation of Mixtures
Physical means can be used
to separate a mixture into its
pure components.
Ex: dyes such as ink may
be separated by paper
chromatography.
magnet
distillation
1.4
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)
Separation of a Compound
The Electrolysis of water
Compounds must be
separated by chemical
means.
With the application of
electricity, water can be
separated into its elements
Reactant

Water

2 H2O

Products
Hydrogen + Oxygen
2 H2
+
O2
Physical vs. Chemical Change
• Physical changes do not
result in new substances.
• Chemical changes result in
NEW substances
Mrs. Kalmer is the one on
the left . . .
The Organization of Matter
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