Chemistry: The Study of Change

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Chemistry The Study of Change
 

• Chapter 1

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Chapter 1 Chemistry The Study of Change
1.1 Chemistry and Some Fundamental
Definitions 1.2 The Scientific Approach
Developing a Model 1.3 Classification of
Matters 1.4 The Three States of Matter 1.5
Physical and Chemical properties of Matter 1.6
Measurement in Scientific Study 1.7 Handling
Numbers 1.8 The factor-Label Method of Solving
Problems
3
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.

water, ammonia, sucrose, gold, oxygen
1.4
4
Definitions
Matter
anything that has mass and volume -the stuff of
the universe books, planets, trees, professors,
students
Composition
the types and amounts of simpler substances that
make up a sample of matter
Properties
the characteristics that give each substance a
unique identity
Chemical Properties those which the substance
shows as it interacts with, or transforms into,
other substances such as flammability,
corrosiveness
Physical Properties those which the substance
shows by itself without interacting with another
substance such as color, melting point, boiling
point, density
5
Sample Problem 1.1
Distinguishing Between Physical and Chemical
Change
(a) Frost forms as the temperature drops on a
humid winter night.
(b) A cornstalk grows from a seed that is watered
and fertilized.
(c) Dynamite explodes to form a mixture of gases.
(d) Perspiration evaporates when you relax after
jogging.
(e) A silver fork tarnishes in air.
PLAN
Does the substance change composition or just
change form?
SOLUTION
(a) physical change
(b) chemical change
(c) chemical change
(d) physical change
(e) chemical change
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Scientific Approach Developing a Model
Natural phenomena and measured events
universally consistent ones can be stated as a
natural law.
Observations
Hypothesis
Tentative proposal that explains observations.
Procedure to test hypothesis measures one
variable at a time.
Experiment
Set of conceptual assumptions that explains data
from accumulated experiments predicts related
phenomena.
Model (Theory)
Further Experiment
Tests predictions based on model.
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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.

• Heterogeneous mixture composition is not
  uniform throughout.

1.4
8
Physical means can be used to separate a mixture
into its pure components.
1.4
9

• An element is a substance that cannot be
  separated into simpler substances by chemical
  means.
• 115 elements have been identified
• 83 elements occur naturally on Earth
• gold, aluminum, lead, oxygen, carbon
• 32 elements have been created by scientists
• technecium, amerecium, seaborgium

1.4
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A compound is a substance composed of atoms of
two or more elements chemically united in fixed
proportions.
Compounds can only be separated by chemical means
into their pure components (elements).
1.4
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1.4
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Three States of Matter
1.5
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Physical or Chemical?
A physical change does not alter the composition
or identity of a substance.
A chemical change alters the composition or
identity of the substance(s) involved.
1.6
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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
1.7
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1.7
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Common Decimal Prefixes Used with SI Units
Table 1.3
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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
1.7
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Common SI-English Equivalent Quantities
Table 1.4
English to SI Equivalent
English Equivalent
Quantity
SI Unit
SI Equivalent
Length
1 kilometer(km)
1000(103)m
0.62miles(mi)
1 mi 1.61km
1 meter(m)
100(102)m
1.094yards(yd)
1 yd 0.9144m
1000(103)mm
39.37inches(in)
1 foot (ft) 0.3048m
1 centimeter(cm)
0.01(10-2)m
0.3937in
1 in 2.54cm (exactly!)
1 kilometer(km)
1000(103)m
0.62mi
Volume
1,000,000(106) cubic centimeters
35.2cubic feet (ft3)
1 cubic meter(m3)
1 ft3 0.0283m3
1 cubic decimeter (dm3)
1000cm3
0.2642 gallon (gal) 1.057 quarts (qt)
1 gal 3.785 dm3
1 qt 0.9464 dm3
1 cubic centimeter (cm3)
0.001 dm3
0.0338 fluid ounce
1 qt 946.4 cm3
1 fluid ounce 29.6 cm3
Mass
1 kilogram (kg)
1000 grams
2,205 pounds (lb)
1 (lb) 0.4536 kg
1 gram (g)
1000 milligrams
0.03527 ounce(oz)
1 lb 453.6 g
1 ounce 28.35 g
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Some Interesting Quantities
Length Volume Mass
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Density SI derived unit for density is kg/m3
1 g/cm3 1 g/mL 1000 kg/m3
m d x V
21.5 g/cm3 x 4.49 cm3 96.5 g
1.7
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Densities of Some Common Substances
Table 1.5
Substance Physical State
Density (g/cm3)
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Sample Problem 1.5
Calculating Density from Mass and Length
PLAN
Density is expressed in g/cm3 so we have to the
mass in grams and the volume in cm3.
SOLUTION
lengths (mm) of sides
1.49g
1.49x103mg x
mass (mg) of Li
lengths (cm) of sides
20.9mm x
2.09cm
Similarly the other sides will be 1.11cm and
1.20cm, respectively.
mass (g) of Li
volume (cm3)
2.09 x 1.11 x 1.20 2.76cm3
density (g/cm3) of Li
density of Li
0.540g/cm3
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The Freezing and Boiling Points of Water
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Temperature Scales and Interconversions
Kelvin ( K ) - The Absolute temperature scale
begins at absolute zero and only has positive
values.
Celsius ( oC ) - The temperature scale used by
science, formally called centigrade and most
commonly used scale around the world, water
freezes at 0oC, and boils at 100oC.
Fahrenheit ( oF ) - Commonly used scale in the
U.S. for our weather reports water freezes at
32oF,and boils at 212oF.
T (in K) T (in oC) 273.15 T (in oC) T (in
K) - 273.15
T (in oF) 9/5 T (in oC) 32 T (in oC) T
(in oF) - 32 5/9
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Some Interesting Temperatures
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Convert 172.9 0F to degrees Celsius.
1.7
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Scientific Notation
6.022 x 1023
1.99 x 10-23
N x 10n
N is a number between 1 and 10
n is a positive or negative integer
1.8
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Scientific Notation
568.762
0.00000772
n gt 0
n lt 0
568.762 5.68762 x 102
0.00000772 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 104 3.9 x 103
4.31 x 104 0.39 x 104
4.70 x 104
1.8
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Scientific Notation
Multiplication
(4.0 x 10-5) x (7.0 x 103) (4.0 x 7.0) x
(10-53) 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 104-9 1.7
x 10-5

• Divide N1 and N2
• Subtract exponents n1 and n2

1.8
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The Number of Significant Figures in a
Measurement Depends Upon the Measuring Device
Figure 1.14
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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

1.8
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Issues Concerning Significant Figures
be sure to correlate with the problem
FIX function on some calculators
graduated cylinder lt buret pipet
60 min 1 hr
numbers with no uncertainty
1000 mg 1 g
These have as many significant digits as the
calculation requires.
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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
2 significant figures
1.8
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Rules for Rounding Off Numbers
1. If the digit removed is more than 5, the
preceding number increases by 1. 5.379 rounds to
5.38 if three significant figures are retained
and to 5.4 if two significant figures are
retained.
2. If the digit removed is less than 5, the
preceding number is unchanged. 0.2413 rounds to
0.241 if three significant figures are retained
and to 0.24 if two significant figures are
retained.
3.If the digit removed is 5, the preceding number
increases by 1 if it is odd and remains unchanged
if it is even. 17.75 rounds to 17.8, but 17.65
rounds to 17.6. If the 5 is followed only by
zeros, rule 3 is followed if the 5 is followed
by nonzeros, rule 1 is followed 17.6500 rounds
to 17.6, but 17.6513 rounds to 17.7
4. Be sure to carry two or more additional
significant figures through a multistep
calculation and round off only the final answer.
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Significant Figures
Addition or Subtraction
The answer cannot have more digits to the right
of the decimal point than any of the original
numbers.
1.8
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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
6.8 112.04 0.0606926
0.061
1.8
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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?
Because 3 is an exact number
1.8
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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
1.8
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Precision and Accuracy in the Laboratory
Figure 1.16
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Precision and Accuracy in the Laboratory
Figure 1.16 continued
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Factor-Label 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.

How many mL are in 1.63 L?
1 L 1000 mL
1.9
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The speed of sound in air is about 343 m/s. What
is this speed in miles per hour?
meters to miles
seconds to hours
1 mi 1609 m
1 min 60 s
1 hour 60 min
1.9