Chemistry
Fourth Edition
Julia Burdge
Lecture PowerPoints
Chapter 1
Chemistry: The Central
Science
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CHAPTER
1.1
1.2
1.3
1.4
1.5
1.6
1
Chemistry: The Central
Science
The Study of Chemistry
Classification of Matter
Scientific Measurement
The Properties of Matter
Uncertainty in Measurement
Using Units and Solving Problems
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2
1.1
The Study of Chemistry
Topics
Chemistry You May Already Know
The Scientific Method
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3
1.1
The Study of Chemistry
Chemistry You May Already Know
Chemistry is the study of matter and the changes that matter
undergoes.
Matter is what makes up our bodies, our belongings, our
physical environment, and in fact our universe. Matter is
anything that has mass and occupies space.
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4
1.1
The Study of Chemistry
Chemistry You May Already Know
© Steve Allen/Getty
© The McGraw-Hill Companies, Inc./Charles D. Winters, photographer
© Stockbyte/PunchStock
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5
1.1
The Study of Chemistry
The Scientific Method
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6
1.1
The Study of Chemistry
The Scientific Method
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7
1.1
The Study of Chemistry
The Scientific Method
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8
1.2
Classification of Matter
Topics
States of Matter
Elements
Compounds
Mixtures
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9
1.2
Classification of Matter
States of Matter
Chemists classify matter as either a substance or a mixture of
substances.
A substance may be further categorized as either an element
or a compound.
A substance is a form of matter that has a definite (constant)
composition and distinct properties.
Examples are salt (sodium chloride), iron, water, mercury,
carbon dioxide, and oxygen.
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10
1.2
Classification of Matter
States of Matter
Substances can be either elements (such as iron, mercury,
and oxygen) or compounds (such as salt, water, and carbon
dioxide).
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11
1.2
Classification of Matter
States of Matter
All substances can, in principle, exist as a solid, a liquid, and a
gas, the three physical states .
Solids and liquids sometimes are referred to collectively as
the condensed phases.
Liquids and gases sometimes are referred to collectively as
fluids.
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12
1.2
Classification of Matter
States of Matter
© The McGraw-Hill Companies, Inc./Charles
D. Winters, photographer
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13
1.2
Classification of Matter
Elements
An element is a substance that cannot be separated into
simpler substances by chemical means.
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14
1.2
Classification of Matter
Compounds
Most elements can combine with other elements to form
compounds.
A compound is a substance composed
of atoms of two or more elements
chemically united in fixed
proportions.
A compound cannot be separated into
simpler substances by any physical
process.
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15
1.2
Classification of Matter
Mixtures
A mixture is a combination of two or
more substances in which the
substances retain their distinct
identities.
Like pure substances, mixtures can be solids, liquids, or gases.
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16
1.2
Classification of Matter
Mixtures
Mixtures are either homogeneous or heterogeneous.
When we dissolve a teaspoon of sugar in a glass of water, we
get a homogeneous mixture because the composition of the
mixture is uniform throughout.
If we mix sand with iron filings, however, the sand and the
iron filings remain distinct and discernible from each other.
This type of mixture is called a heterogeneous mixture
because the composition is not uniform.
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17
1.2
Classification of Matter
Mixtures
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18
1.3
Scientific Measurement
Topics
SI Base Units
Mass
Temperature
Derived Units: Volume and Density
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19
1.3
Scientific Measurement
SI Base Units
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20
1.3
Scientific Measurement
SI Base Units
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21
1.3
Scientific Measurement
Mass
Mass is a measure of the amount of matter in an object or
sample.
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22
1.3
Scientific Measurement
Temperature
There are two temperature scales used in chemistry.
The Celsius scale was originally defined using the freezing
point (0°C) and the boiling point (100°C) of pure water at sea
level.
The SI base unit of temperature is the kelvin. Kelvin is known
as the absolute temperature scale, meaning that the lowest
temperature possible is 0 K, a temperature referred to as
“absolute zero.”
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23
1.3
Scientific Measurement
Temperature
Units of the Celsius and Kelvin scales are equal in magnitude,
so a degree Celsius is equivalent to a kelvin.
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24
SAMPLE PROBLEM
1.1
Normal human body temperature can range over the course
of the day from about 36°C in the early morning to about
37°C in the afternoon.
Express these two temperatures and the range that they span
using the Kelvin scale.
Strategy
Use K = °C + 273.15 to convert temperatures from the Celsius
scale to the Kelvin scale.
Then convert the range of temperatures from degrees Celsius
to kelvin, keeping in mind that 1°C is equivalent to 1 K.
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25
SAMPLE PROBLEM
1.1
Solution
36°C + 273 = 309 K
37°C + 273 = 310 K
The range of 1°C is equal to a range of 1 K.
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26
1.3
Scientific Measurement
Temperature
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27
SAMPLE PROBLEM
1.2
A body temperature above 39°C constitutes a high fever.
Convert this temperature to the Fahrenheit scale.
Strategy
We are given a temperature in Celsius and are asked to
convert it to Fahrenheit.
Setup
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28
SAMPLE PROBLEM
1.2
Solution
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29
1.3
Scientific Measurement
Derived Units: Volume and Density
There are many quantities, such as volume and density, that
require units not included in the base SI units.
In these cases, we must combine base units to derive
appropriate units for the quantity.
The derived SI unit for volume, the meter cubed (m3), is a
larger volume than is practical in most laboratory settings.
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30
1.3
Scientific Measurement
Derived Units: Volume and Density
The more commonly used
metric unit, the liter (L), is
derived by cubing the
decimeter (one-tenth of a
meter) and is therefore also
referred to as the cubic
decimeter (dm3).
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31
1.3
Scientific Measurement
Derived Units: Volume and Density
Density is the ratio of mass to volume.
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32
SAMPLE PROBLEM
1.3
Ice cubes float in a glass of water because solid water is less
dense than liquid water.
(a) Calculate the density of ice given that, at 0°C, a cube that
is 2.0 cm on each side has a mass of 7.36 g, and
(b) determine the volume occupied by 23 g of ice at 0°C.
Strategy
(a) Determine density by dividing mass by volume.
(b) use the calculated density to determine the volume
occupied by the given mass.
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33
SAMPLE PROBLEM
1.3
Setup
(a) We are given the mass of the ice cube, but we must
calculate its volume from the dimensions given. The
volume of the ice cube is (2.0 cm)3, or 8.0 cm3.
(b) Rearranging the density equation to solve for volume
gives V = m/d.
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34
SAMPLE PROBLEM
1.3
Solution
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35
1.4
The Properties of Matter
Topics
Physical Properties
Chemical Properties
Extensive and Intensive Properties
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36
1.4
The Properties of Matter
Physical Properties
Substances are identified by their properties as well as by
their composition.
Properties of a sub- stance may be quantitative (measured
and expressed with a number) or qualitative (not requiring
explicit measurement).
A physical property is one that can be observed and
measured without changing the identity of a substance.
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37
1.4
The Properties of Matter
Physical Properties
Melting is a physical change; one in which the state of matter
changes, but the identity of the matter does not change.
We can recover the original ice by cooling the water until it
freezes.
Therefore, the melting point of a substance is a physical
property.
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38
1.4
The Properties of Matter
Chemical Properties
The statement “Hydrogen gas burns in oxygen gas to form
water” describes a chemical property of hydrogen, because
to observe this property we must carry out a chemical
change—burning in oxygen (combustion), in this case.
After a chemical change, the original substance (hydrogen gas
in this case) will no longer exist.
What remains is a different substance (water, in this case). We
cannot recover the hydrogen gas from the water by means of
a physical process, such as boiling or freezing.
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39
1.4
The Properties of Matter
Extensive and Intensive Properties
All properties of matter are either extensive or intensive.
The measured value of an extensive property depends on the
amount of matter.
Mass is an extensive property.
The value of an intensive property does not depend on the
amount of matter.
Density and temperature are intensive properties.
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40
SAMPLE PROBLEM
1.4
The diagram in (a) shows a compound made up of atoms of
two elements (represented by the green and red spheres) in
the liquid state.
Which of the diagrams in (b) to (d) represent a physical
change, and which diagrams represent a chemical change?
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41
SAMPLE PROBLEM
1.4
Strategy
A physical change does not change the identity of a
substance, whereas a chemical change does change the
identity of a substance.
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42
SAMPLE PROBLEM
1.4
Solution
Diagrams (b) and (c) represent chemical changes. Diagram (d)
represents a physical change.
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43
1.5
Uncertainty in Measurement
Topics
Significant Figures
Calculations with Measured Numbers
Accuracy and Precision
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44
1.5
Uncertainty in Measurement
Significant Figures
Chemistry makes use of two types of numbers: exact and
inexact.
Exact numbers include numbers with defined values, such as
2.54 in the definition 1 inch (in) = 2.54 cm
1000 in the definition 1 kg = 1000 g
12 in the definition 1 dozen = 12 objects.
Numbers measured by any method other than counting are
inexact.
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45
1.5
Uncertainty in Measurement
Significant Figures
An inexact number must be
reported in such a way as to
indicate the uncertainty in its value.
This is done using significant
figures. Significant figures are the
meaningful digits in a reported
number.
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46
1.5
Uncertainty in Measurement
Significant Figures
The number of significant figures in any number can be
determined using the following guidelines:
1. Any digit that is not zero is significant (112.1 has four
significant figures).
2. Zeros located between nonzero digits are significant (305
has three significant figures, and 50.08 has four significant
figures).
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47
1.5
Uncertainty in Measurement
Significant Figures
The number of significant figures in any number can be
determined using the following guidelines:
3. Zeros to the left of the first nonzero digit are not
significant (0.0023 has two significant figures, and
0.000001 has one significant figure).
4. Zeros to the right of the last nonzero digit are significant if
the number contains a decimal point (1.200 has four
significant figures
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48
1.5
Uncertainty in Measurement
Significant Figures
The number of significant figures in any number can be
determined using the following guidelines:
5. Zeros to the right of the last nonzero digit in a number
that does not contain a decimal point may or may not be
significant (100 may have one, two, or three significant
figures—it is impossible to tell without additional
information.
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49
1.5
Uncertainty in Measurement
Significant Figures
To avoid ambiguity in such cases, it is best to express such
numbers using scientific notation [Appendix 1].
1.3 × 102 two significant figures
1.30 × 102 three significant figures
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50
SAMPLE PROBLEM
1.5
Determine the number of significant figures in the following
measurements:
(a)
(b)
(c)
(d)
(e)
(f)
443 cm
15.03 g
0.0356 kg
3.000 3 10–7 L
50 mL
0.9550 m
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51
SAMPLE PROBLEM
1.5
Strategy
All nonzero digits are significant, so the goal will be to
determine which of the zeros is significant.
Setup
Zeros are significant if they appear between nonzero digits or
if they appear after a nonzero digit in a number that contains
a decimal point.
Zeros may or may not be significant if they appear to the right
of the last nonzero digit in a number that does not contain a
decimal point.
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52
SAMPLE PROBLEM
Solution
(a)
(b)
(c)
(d)
(e)
(f)
1.5
443 cm
15.03 g
0.0356 kg
3.000 3 10–7 L
50 mL
0.9550 m
(a) 3; (b) 4; (c) 3; (d) 4; (e) 1 or 2, an ambiguous case; (f) 4
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53
1.5
Uncertainty in Measurement
Calculations with Measured Numbers
1. In addition and subtraction, the answer cannot have
more digits to the right of the decimal point than the
original number with the smallest number of digits to the
right of the decimal point.
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54
1.5
Uncertainty in Measurement
Calculations with Measured Numbers
If the leftmost digit to be dropped is less than 5, round down.
If the leftmost digit to be dropped is equal to or greater than
5, round up.
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55
1.5
Uncertainty in Measurement
Calculations with Measured Numbers
2. In multiplication and division, the number of significant
figures in the final product or quotient is determined by
the original number that has the smallest number of
significant figures.
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56
1.5
Uncertainty in Measurement
Calculations with Measured Numbers
3. Exact numbers can be considered to have an infinite
number of significant figures and do not limit the
number of significant figures in a calculated result.
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57
1.5
Uncertainty in Measurement
Calculations with Measured Numbers
4. In calculations with multiple steps, rounding the result of
each step can result in “rounding error.”
In general, it is best to retain at least one extra digit until
the end of a multistep calculation to minimize rounding
error.
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58
SAMPLE PROBLEM
1.6
Perform the following arithmetic operations and report the
result to the proper number of significant figures:
(a)
(b)
(c)
(d)
(e)
317.5 mL + 0.675 mL
47.80 L – 2.075 L
13.5 g ÷ 45.18 L
6.25 cm × 1.175 cm
5.46 × 102 g + 4.991 × 103 g
Strategy
Apply the rules for significant figures in calculations, and
round each answer to the appropriate number of digits.
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59
SAMPLE PROBLEM
1.6
Setup
(a)
(b)
(c)
(d)
(e)
317.5 mL + 0.675 mL
47.80 L – 2.075 L
13.5 g ÷ 45.18 L
6.25 cm × 1.175 cm
5.46 × 102 g + 4.991 × 103 g
(a) The answer will contain one digit to the right of the
decimal point to match 317.5, which has the fewest digits
to the right of the decimal point.
(b) The answer will contain two digits to the right of the
decimal point to match 47.80.
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60
SAMPLE PROBLEM
1.6
Setup
(a)
(b)
(c)
(d)
(e)
317.5 mL + 0.675 mL
47.80 L – 2.075 L
13.5 g ÷ 45.18 L
6.25 cm × 1.175 cm
5.46 × 102 g + 4.991 × 103 g
(c) The answer will contain three significant figures to match
13.5, which has the fewest number of significant figures in
the calculation.
(d) The answer will contain three significant figures to match
6.25.
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61
SAMPLE PROBLEM
1.6
Setup
(a)
(b)
(c)
(d)
(e)
317.5 mL + 0.675 mL
47.80 L – 2.075 L
13.5 g ÷ 45.18 L
6.25 cm × 1.175 cm
5.46 × 102 g + 4.991 × 103 g
(e) To add numbers expressed in scientific notation, first write
both numbers to the same power of 10.
That is, 4.991 × 103 = 49.91 × 102, so the answer will
contain two digits to the right of the decimal point (when
multiplied by 102) to match both 5.46 and 49.91.
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62
SAMPLE PROBLEM
1.6
Solution
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63
SAMPLE PROBLEM
1.6
Solution
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64
SAMPLE PROBLEM
1.7
An empty container with a volume of 9.850 × 102 cm3 is
weighed and found to have a mass of 124.6 g.
The container is filled with a gas and reweighed. The mass of
the container and the gas is 126.5 g.
Determine the density of the gas to the appropriate number
of significant figures.
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65
SAMPLE PROBLEM
1.7
Strategy
This problem requires two steps: subtraction to determine
the mass of the gas, and division to determine its density.
Apply the corresponding rule regarding significant figures to
each step.
Setup
126.5 g – 124.6 g = 1.9 g.
Thus, in the division of the mass of the gas by the volume of
the container, the result can have only two significant figures.
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66
SAMPLE PROBLEM
1.7
Solution
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67
1.5
Uncertainty in Measurement
Accuracy and Precision
Accuracy tells us how close a measurement is to the true
value.
Precision tells us how close multiple measurements of the
same thing are to one another.
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68
1.5
Uncertainty in Measurement
Accuracy and Precision
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69
1.5
Uncertainty in Measurement
Accuracy and Precision
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70
1.6
Using Units and Solving Problems
Topics
Conversion Factors
Dimensional Analysis—Tracking Units
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71
1.6
Using Units and Solving Problems
Conversion Factors
A conversion factor is a fraction in which the same quantity is
expressed one way in the numerator and another way in the
denominator.
Because both forms of this conversion factor are equal to 1,
we can multiply a quantity by either form without changing
the value of that quantity.
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72
1.6
Using Units and Solving Problems
Conversion Factors
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73
1.6
Using Units and Solving Problems
Dimensional Analysis—Tracking Units
The use of conversion factors in problem solving is called
dimensional analysis or the factor-label method.
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74
SAMPLE PROBLEM
1.8
The Food and Drug Administration (FDA) recommends that
dietary sodium intake be no more than 2400 mg per day.
What is this mass in pounds (lb), if 1 lb = 453.6 g?
Strategy
This problem requires a two-step dimensional analysis,
because we must convert milligrams to grams and then grams
to pounds.
Assume the number 2400 has four significant figures.
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75
SAMPLE PROBLEM
1.8
Setup
The necessary conversion factors are derived from the
equalities 1 g = 1000 mg and 1 lb = 453.6 g.
Solution
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76
SAMPLE PROBLEM
1.9
An average adult has 5.2 L of blood. What is the volume of
blood in cubic meters?
Strategy
Convert liters to cubic centimeters and then cubic
centimeters to cubic meters.
Setup
1 L = 1000 cm3 and 1 cm = 1 × 10–2 m.
When a unit is raised to a power, the corresponding
conversion factor must also be raised to that power in order
for the units to cancel appropriately.
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77
SAMPLE PROBLEM
1.9
Solution
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78