Chapter 1

Report
Chemistry: Atoms First
Julia Burdge & Jason Overby
Chapter 1
Chemistry:
The Science of Change
Homework: 5, 9, 15, 17, 23,25, 27, 31,
37, 29, 43, 45, 47, 49, 51, 53,
55, 57, 59, 61, 63, 65, 67, 69,
71, 73, 75, 77 and 79
Kent L. McCorkle
Cosumnes River College
Sacramento, CA
Copyright (c) The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
1
1
Chemistry: The Science of Change
1.1 The Study of Chemistry
1.5 Uncertainty in Measurement
Chemistry You May Already Know
Significant Figures
The Scientific Method
Calculations with Measured Numbers
1.2 Classification of Matter
Accuracy and Precision
States of Matter
1.6 Using Units and Solving Problems
Mixtures
Conversion Factors
1.3 The Properties of Matter
Dimensional Analysis
Physical Properties
Tracking Units
Chemical Properties
Extensive and Intensive Properties
1.4 Scientific Measurement
SI Base Units
Mass
Temperature
Derived Units: Volume and Density
1.1 The Study of Chemistry
Chemistry is the study of matter and the changes that matter undergoes.
Matter is anything that has mass and occupies space.
The Study of Chemistry
Scientists follow a set of guidelines known as the scientific method:
•
gather data via observations and experiments
•
identify patterns or trends in the collected data
•
summarize their findings with a law
•
formulate a hypothesis
•
with time a hypothesis may evolve into a theory
1.2 Classification of Matter
Chemists classify matter as either a substance or a mixture of substances.
A substance is a form of matter that has definite composition and distinct properties.
 Examples: salt (sodium chloride), iron, water, mercury, carbon dioxide, and
oxygen
Substances differ from one another in composition and may be identified by
appearance, smell, taste, and other properties.
A mixture is a physical combination of two or more substances.
A homogeneous mixture is uniform throughout.
 Also called a solution.
 Examples: seawater, apple juice
A heterogeneous mixture is not uniform throughout.
 Examples: trail mix, chicken noodle soup
Classification of Matter
All substances can, in principle,
exist as a solid, liquid or gas.
We can convert a substance from
one state to another without
changing the identity of the
substance.
Classification of Matter
Solidsparticles
Solid
do not conform
are held
closely
to
the shape
together
of their
in an
ordered fashion.
container.
Liquidsparticles
Liquid
do conform
are close
to
together
the
shapebut
of are
theirnot held
rigidly in position.
container.
Gasesparticles
Gas
assume have
both the
significant
shape
and volume
separation
of
from container.
their
each other and
move freely.
Classification of Matter
A mixture can be separated by physical means into its components
without changing the identities of the components.
1.3 The Properties of Matter
There are two general types of properties of matter:
1) Quantitative properties are measured and expressed with a
number.
2) Qualitative properties do not require measurement and are
usually based on observation.
The Properties of Matter
A physical property is one that can be observed and measured
without changing the identity of the substance.
 Examples: color, melting point, boiling point
A physical change is one in which the state of matter changes, but
the identity of the matter does not change.
 Examples: changes of state (melting, freezing, condensation)
The Properties of Matter
A chemical property is one a substance exhibits as it interacts with
another substance.
 Examples: flammability, corrosiveness
A chemical change is one that results in a change of composition;
the original substances no longer exist.
 Examples: digestion, combustion, oxidation
The Properties of Matter
An extensive property depends on the amount of matter.
 Examples: mass, volume
An intensive property does not depend on the amount of matter.
 Examples: temperature, density
1.4 Scientific Measurement
Properties that can be measured
are called quantitative
properties.
A measured quantity must
always include a unit.
The English system has units
such as the foot, gallon, pound,
etc.
The metric system includes
units such as the meter, liter,
kilogram, etc.
SI Base Units
The revised metric system is called the International System of
Units (abbreviated SI Units) and was designed for universal use by
scientists.
There are seven SI base units
Units in Measurements
Factor
Prefix
Symbol
1 x 106
Mega
M
1 x 103
Kilo
k
the base units
grams (g) meters (m)
Moles (mol) volume (L)
1 x 10-2
centi
c
1 x 10-3
milli
m
1 x 10-6
micro
m
1 x 10-9
nano
n
1 x 10-12
pico
p
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Mass
Mass is a measure of the amount of matter in an object or sample.
Because gravity varies from location to location, the weight of an
object varies depending on where it is measured. But mass doesn’t
change.
The SI base unit of mass is the kilogram (kg), but in chemistry the
smaller gram (g) is often used.
1 kg = 1000 g = 1×103 g
Atomic mass unit (amu) is used to express the masses of atoms and
other similar sized objects.
1 amu = 1.6605378×10-24 g
Temperature
There are two temperature scales used in chemistry:
The Celsius scale (°C)
 Freezing point (pure water): 0°C
 Boiling point (pure water): 100°C
The Kelvin scale (K)
 The “absolute” scale
 Lowest possible temperature: 0 K (absolute zero)
K = °C + 273.15
Worked Example 1.1
Normal human body temperature can range over the course of a 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 Celsius to Kelvin.
Solution 36°C + 273 = 309 K
37°C + 273 = 310 K
What range do they span?
Depending on the precision
required, the conversion from
°C to K is often simply done by
adding 273, rather than 273.15.
310 K - 309 K = 1 K
Think About It Remember that converting a temperature from °C to K is
different from converting a range or difference in temperature from °C to K.
Temperature
The Fahrenheit scale is common in the United States.
 Freezing point (pure water): 32°C
 Boiling point (pure water): 212°C
There are 180 degrees between freezing and boiling in Fahrenheit
(212°F-32°F) but only 100 degrees in Celsius (100°C-0°C).
 The size of a degree on the Fahrenheit scale is only 9 of a
5
degree on the Celsius scale.
Temp in °F = (
9
5
×temp in °C ) + 32°F
Worked Example 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 and asked to convert it to degrees
Fahrenheit. We will use the equation below:
Temp in °F = (
9
5
× temp in °C ) + 32°F
9
Solution Temp in °F = ( × 39°C ) + 32°F
5
Temp in °F = 102°F
Think About It Knowing that normal body temperature on the Fahrenheit
scale is approximately 98.6°F, 102°F seems like a reasonable answer.
Taking Measurements
While Erlenmeyer flasks and beakers have volume markings, those
Markings are only approximations and should NEVER be used for
taking measurement readings
Taylor 2010
Taking Measurements
Graduated cylinders are moderately accurate and are primarily what
we use in the laboratory for measuring volumes of liquids
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Taking Measurements
Burettes, Pipettes and volumetric flasks are the most accurate way to make
solutions and to measure volumes. However, they tend to be more
difficult to use than graduated cylinders.
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Taking Measurements
You want to read the volume of solution at the meniscus and
you may need to get down to eye level with the graduated cylinder
to read the meniscus properly.
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Uncertainty in Measurement
An inexact number must be reported so as to indicate its uncertainty.
Significant figures are the
meaningful digits in a reported
number.
The last digit in a measured number
is referred to as the uncertain digit.
When using the top ruler to measure
the memory card, we could estimate
2.5 cm. We are certain about the 2,
but we are not certain about the 5.
The uncertainty is generally
considered to be + 1 in the last digit.
2.5 + 0.1 cm
Uncertainty in Measurement
When using the bottom ruler to
measure the memory card, we might
record 2.45 cm.
Again, we estimate one more digit
than we are certain of.
2.45 + 0.01 cm
Taking Measurements
Which piece of glassware would you use to get the
most accurate measured value?
or
Beaker
Graduated cylinder
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Taking Measurements
Let’s consider the significant figure
value we would obtain?
What is our certain measured value?
What is our estimated value ?
13
10
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Taking Measurements
Let’s consider the significant figure
value we would obtain?
What is our certain measured value?
What is our estimated value ?
13.51
13.5
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Taking Measurements
Which piece of glassware would you use to get the most
accurate measured value?
or
Beaker
Graduated cylinder
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Significant Figures
• Why?
– To show the certainty in a measured value
– To indicate the margin of error when
measuring
• How?
– Report all known values
– Estimate one value past what’s given
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Significant Figures Rules
•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
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Significant Figures Rules
•Any number that has zeros after the digits and no
decimal point according to Tro may have an infinite
number of significant figures
•However, there is considerable number of scientists who
believe that zeros that come after the digits when a
decimal point is not present are considered to be NOT
significant
•For example the number 54, 000 would have infinite or 2
significant figures where as 54, 000. would have 5
Taylor 2010
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
2 significant figures
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Significant Figures
How many significant figures are there in 1.3070 g?
A. 6
B. 5
C. 4
Ans: B
D. 3
E. 2
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Worked Example 1.4
Determine the number of significant figures in the following measurements: (a)
443 cm, (b) 15.03 g, (c) 0.0356 kg, (d) 3.000×10-7 L, (e) 50 mL, (f) 0.9550 m.
Strategy Zeros are significant between nonzero digits or after a nonzero digit
with a decimal. Zeros may or may not be significant if they appear to the right of
a nonzero digit without a decimal.
Solution (a) 443 cm
3 S.F.
(c) 0.0356 kg
3 S.F.
(e) 50 mL
1 or 2, ambiguous
(b) 15.03 g
4 S.F.
(d) 3.000 x 10-7 L
4 S.F.
(f) 0.9550 m
4 S.F.
Think About It Be sure that you have identified zeros correctly as either
significant or not significant. They are significant in (b) and (d); they are not
significant in (c); it is not possible to tell in (e); and the number in (f)
contains one zero that is significant, and one that is not.
Significant Figures
Addition and Subtraction of Significant Figures
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 significant figure after decimal point
round off to 90.4
two significant figures after decimal point
round off to 0.79
Taylor 2010
Significant Figures
Multiplication and Division of Significant Figures
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
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Significant Figures
What are 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?
6.64 + 6.68 + 6.70
= 6.67333 = 6.673 = 7
3
Because 3 is an exact number
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Significant Figures
After carrying out the following operations, how many
significant figures are appropriate to show in the
result?
(13.7 + 0.027)  8.221
A. 1
B. 2
Ans: C
C. 3
D. 4
E. 5
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Calculations with Measured Numbers
In addition and subtraction, the answer cannot have more digits to
the right of the decimal point than any of the original numbers.
102.50 ← two digits after the decimal point
+ 0.231 ← three digits after the decimal point
102.731 ← round to two digits after the decimal point, 102.73
143.29
- 20.1
123.19
← two digits after the decimal point
← one digit after the decimal point
← round to one digit after the decimal point, 123.2
Worked Example 1.5
Perform the following arithmetic operations and report the result to the proper
number of significant figures: (a) 317.5 mL + 0.675 mL, (b) 47.80 L – 2.075 L,
(c) 13.5 g ÷ 45.18 L, (d) 6.25 cm x 1.175 cm, (e) 5.46x102 g + 4.991x103 g
Strategy Apply the rules for significant figures in calculations, and round each
answer to the appropriate number of digits.
Solution (a) 317.5 mL
+ 0.675 mL
318.175 mL
(b) 47.80 L
- 2.075 L
45.725 L
← round to 318.2 mL
← round to 45.73 L
Worked Example 1.5 (cont.)
Perform the following arithmetic operations and report the result to the proper
number of significant figures: (a) 317.5 mL + 0.675 mL, (b) 47.80 L – 2.075 L,
(c) 13.5 g ÷ 45.18 L, (d) 6.25 cm x 1.175 cm, (e) 5.46x102 g + 4.991x103 g
Strategy Apply the rules for significant figures in calculations, and round each
answer to the appropriate number of digits.
Solution (e) 5.46 x 102 g
+ 49.91 x 102 g
55.37 x 102 g
= 5.537 x 103 g
Think About It Changing the answer to correct scientific notation doesn’t
change the number of significant figures, but in this case it changes the number of
places past the decimal place.
Worked Example 1.5 (cont.)
Perform the following arithmetic operations and report the result to the proper
number of significant figures: (a) 317.5 mL + 0.675 mL, (b) 47.80 L – 2.075 L,
(c) 13.5 g ÷ 45.18 L, (d) 6.25 cm x 1.175 cm, (e) 5.46x102 g + 4.991x103 g
Strategy Apply the rules for significant figures in calculations, and round each
answer to the appropriate number of digits.
Solution
3 S.F.
(c) 13.5 g
= 0.298804781 g/L
45.18 L
← round to 0.299 g/L
4 S.F.
(d) 6.25 cm×1.175 cm = 7.34375 cm2
3 S.F.
4 S.F.
← round to 7.34 cm2
Worked Example 1.6
An empty container with a volume of 9.850 x 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.
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.
Solution
126.5 g
– 124.6 g
mass of gas = 1.9 g
density =
← one place past the decimal point (two sig figs)
1.9 g
= 0.00193 g/cm3
2
3
9.850 x 10 cm
← round to 0.0019 g/cm3
Think About It In this case, although each of the three numbers we started
with has four significant figures, the solution only has two significant figures.
Calculations with Measured Numbers
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.
1.4×8.011 = 11.2154 ← fewest significant figures is 2, so
round to 11
2 S.F. 4 S.F.
11.57/305.88 = 0.0378252
4 S.F.
5 S.F.
← fewest significant figures is 4, so
round to 0.03783
Calculations with Measured Numbers
Exact numbers can be considered to have an infinite number of
significant figures and do not limit the number of significant figures
in a result.
Example: Three pennies each have a mass of 2.5 g. What is the
total mass?
3×2.5 = 7.5 g
Exact
(counting number)
Inexact
(measurement)
Scientific Notation
Converting Numbers to Scientific Notation
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
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Calculations with Measured Numbers
In calculations with multiple steps, round at the end of the
calculation to reduce any rounding errors.
Do not round after each step.
Compare the following:
Rounding after each step
Rounding at end
1) 3.66×8.45 = 30.9
2) 30.9×2.11 = 65.2
1) 3.66×8.45 = 30.93
2) 30.93×2.11 = 65.3
In general, keep at least one extra digit until the end of a multistep
calculation.
Scientific Notation
• Why?
– To express large or small numbers in a
simpler format
– All numbers provided in scientific notation are
significant
• How?
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
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Scientific Notation
Addition and Subtraction Of numbers in Scientific notation
4.31 x 104 + 3.9 x 103 =
1. Write each quantity with
the same exponent n
4.31 x 104 + 0.39 x 104 =
2. Combine N1 and N2
3. The exponent, n, remains
the same
4.70 x 104
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Scientific Notation
Multiplication and Division In Scientific Notation
Multiplication Rules
1. Multiply N1 and N2
(4.0 x 10-5) x (7.0 x 103) =
(4.0 x 7.0) x (10-5+3) =
2. Add exponents n1 and n2
28 x 10-2 =
3. Check to make sure N3
is between 1-10
2.8 x 10-1
Division Rules
1. Divide N1 and N2
8.5 x 104 ÷ 5.0 x 109 =
(8.5 ÷ 5.0) x 104-9 =
2. Subtract exponents n1 and n2
3. Check to make sure N3 is between 1-10
1.7 x 10-5
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Scientific Notation
Put the following values into Scientific Notation
24 mL
2.4 x 101 mL
3001 g
3.001 x 103 g
0.0320 m3
3.20 x 10-2 m3
640,000,000 molecules
0.000000000091 kg
6.4 x 108 molecules
9.1 x 10-12 kg
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Scientific Notation
Do the following Mathematical Calculations
using Scientific Notation
2.4 x 104 +
4.4 x 10-5
÷
7.19 x 1015 X
5.6 x 10-4 -
3.72 x 103 =
2.8 x 104
5.92 x 102 =
7.4 x 10-8
8.345 x 10-5 =
6.00 x 1011
3.0 x 10-3 =
5.9 x 10-4
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Units in Measurements
Hints for converting between units
•Always write fractions on 2 lines
•Always use scientific notation for your factor
•Always go through the “base”
•Always pair the base with the factor
Remember these rules regarding working with numbers in Scientific
Notation
•When you multiply values with exponents, ADD the exponents
•When you divide values with exponents, SUBTRACT the exponent of the
denominator from the numerator
•When you raise an exponent to a power, multiply the exponent by the
power
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Scientific Notation and your Calculator
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Scientific Notation and your Calculator
The scientific notation key
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Derived Units: Volume and Density
There are many units (such as
volume) that require units not
included in the base SI units.
The derived SI unit for volume is
the meter cubed (m3).
A more practical unit for volume is
the liter (L).
 1 dm3 = 1 L
 1 cm3 = 1 mL
Derived Units: Volume and Density
The density of a substance is the ratio of mass to volume.
d = density
m = mass
V = volume
SI-derived unit:
d =
m
V
kilogram per cubic meter (kg/m3)
Other common units: g/cm3 (solids)
g/mL (liquids)
g/L
(gases)
Worked Example 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, and (b) use the
calculated density to determine the volume occupied by the given mass.
Solution (a) A cube has three equal sides so the volume is (2.0 cm)3, or 8.0 cm3
7.36 g
d=
= 0.92 g/cm3
3
8.0 cm
(b) Rearranging d = m/V to solve for volume gives V = m/d
23 g
V=
= 25 cm3
3
0.92 g/cm
Think About It For a sample with a density less than 1 g/cm3, the number
of cubic centimeters should be greater than the number of grams. In this
case, 25 cm3 > 23 g.
Units in Measurements
Conversions to try:
1) 3.72 mm2 to Km2
6) 5.34 g/ L to pg/pL
2) 3.72 pmol to Mmol
7) 5.34 KL to nL
3) 3.72 mol/ L to mmol/ mL
8) 5.34 ng to Mg
4) 3.72 KL to mL
9) 5.34 pm to cm
5) 3.72 ng/ mol to g/ mmol
10) 5.34 g/cm3 to Kg/m3
Answers:
1) 3.72 x 10-12 Km2
4) 3.72 x 106 mL
7) 5.34 x 1012 nL
10) 5.34 x 103 Kg/m3
2) 3.72 x 10-18 Mmol
5) 3.72 x 10-12 g/mmol
8) 5.34 x 10-15 Mg
3) 3.72 mmol/ mL
6) 5.34 pg/pL
9) 5.34 x 10-10 cm
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Problem Solving
Dimensional Analysis
• What is it?
– a problem-solving method that uses the fact that any
number or expression can be multiplied by one
without changing its value
– also called Factor-Label Method or the Unit Factor
Method
• Why use it?
– To describe the same or equivalent "amounts" of what
we are interested in. For example, we know that
1 inch = 2.54 centimeters
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Problem Solving
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
How many mL are in 1.63 L?
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
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Problem solving
A rock climber estimates that the rock face is 155 ft high.
The rope he brought is 65 m long. Is the rope long enough
to reach the top? (1 ft = 0.3048m)
To answer we need the unit conversion of
1 foot = 0.3048 meters
X
65 m
1 foot
0.3048m
=
210 feet
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Problem Solving
The speed limit is 55 miles/hr. What is the
speed limit in standard SI units?
To answer we need the unit conversion of
1 mi = 1.6093 Km
1Km = 1000 m
55 miles
1 hr
1.6093
Km X 1000 m=
X
1 mile
1 Km
8.9x104 m
1 hr
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Density
Typicall the units are g/ mL or g/ cm3
mass
density =
volume
d=
m
V
A piece of platinum metal with a density of 21.5 g/cm3 has a volume of
4.49 mL. What is its mass?
d=
m
V
m=dxV
= 21.5 g/cm3 x 4.49 cm3 = 96.5 g
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Density
A piece of metal with a mass of 114 g was placed into
a graduated cylinder that contained 25.00 mL of
water, raising the water level to 42.50 mL. What is
the density of the metal?
A.
B.
C.
D.
E.
0.154 g/mL
0.592 g/mL
2.68 g/mL
6.51 g/mL
7.25 g/mL
Ans: D
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Density
using density as a conversion factor
what is the mass of a piece of metal with a volume of 15.8 grams if
the density of the metal is 2.7 g/mL?
1 The units of volume are either mL, cm3, or L. Given that the density
is given in g/mL, the volume of the piece of metal is most likely 15.8
mL (and not 15.8 g).
Assuming that is the case, you are trying to convert from mL to
grams, so the labels of the conversion factor must be grams on top
and mL on the bottom so that the mL labels will cancel. The density
tells us the numbers that go into the conversion factor: 2.7 grams of
Al has a volume of 1 mL.
So
15.8 mL x ( 2.7 g / 1 mL ) = _________ grams
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1.6
Using Units and Solving Problems
A conversion factor is a fraction in which the same quantity is
expressed one way in the numerator and another way in the
denominator.
For example, 1 in = 2.54 cm, may be written:
1 in
2.54 cm
or
2.54 cm
1 in
Dimensional Analysis – Tracking Units
The use of conversion factors in problem solving is called
dimensional analysis or the factor-label method.
Example: Convert 12.00 inches to meters.
12.00 in ×
= 30.48 cm
The result contains 4 sig
figs because the
conversion, a definition,
is exact.
Which conversion factor will cancel inches and give us centimeters?
1 in
2.54 cm
or
2.54 cm
1 in
Worked Example 1.7
The Food and Drug Administration (FDA) recommends that dietary sodium
intake be no more than 2400 mg per day.
Strategy The necessary conversion factors are derived from the equalities
1 g = 1000 mg and 1 lb = 453.6 g.
1g
1000 mg
or
Solution
2400 mg ×
1000 mg
1g
1g
1000 mg
1 lb
453.6 g
×
1 lb
453.6 g
or
453.6 g
1 lb
= 0.005291 lb
Think About It Make sure that the magnitude of the result is reasonable and
that the units have canceled properly. If we had mistakenly multiplied by 1000
and 453.6 instead of dividing by them, the result
(2400 mg×1000 mg/g×453.6 g/lb = 1.089×109 mg2/lb) would be
unreasonably large and the units would not have canceled properly.
Worked Example 1.8
An average adult has 5.2 L of blood. What is the volume of blood in cubic
meters?
Strategy 1 L = 1000 cm3 and 1 cm = 1x10-2 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.
Solution
5.2 L ×
cm3
1000
1L
×
10-2
1x
m
1 cm
3
= 5.2 x 10-3 m3
Think About It Based on the preceding conversion factors, 1 L = 1×10-3 m3.
Therefore, 5 L of blood would be equal to 5×10-3 m3, which is close to the
calculated answer.
Accuracy and Precision
Accuracy tells us how close a
measurement is to the true value.
Good accuracy and good precision
Precision tells us how close a series of
replicate measurements are to one another.
Poor accuracy but good precision
Poor accuracy and poor precision
Accuracy and Precision
Three students were asked to find the mass of an aspirin tablet. The
true mass of the tablet is 0.370 g.
Student A: Results are precise but not accurate
Student B: Results are neither precise nor accurate
Student C: Results are both precise and accurate
1
Chapter Summary: Key Points
The Scientific Method
States of Matter
Substances
Mixtures
Physical Properties
Chemical Properties
Extensive and Intensive Properties
SI Base Units
Mass
Temperature
Volume and Density
Significant Figures

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