Chapter 2
Measurement and
Problem Solving
Homework
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Exercises (optional, in Tro textbook)
 1 through 27 (odd)
Problems (in Tro textbook)
 29-65 (odd)
 67-91 (odd)
 93-99 (odd)
Cumulative Problems (in Tro textbook)
 101-117 (odd)
Highlight Problems (optional, in Tro textbook)
 119, 121
Scientific Notation:
Writing Large and Small Numbers
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Scientific Notation
 A system in which an ordinary decimal
number (m) is expressed as a product of a
number between 1 and 10, multiplied by 10
raised to a power (n)
 Used to write very large or very small
numbers
 Based on powers of 10
m  10
n
Scientific Notation:
Writing Large and Small Numbers

Consists of a number (coefficient) followed
by a power of 10 (x 10n)
7.03  10
Coefficient or
decimal part
2
exponent
exponential term
or part

Negative exponent: Number is less than 1
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Positive exponent: Number is greater than 1
Scientific Notation:
Writing Large and Small Numbers
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In an ordinary cup of water there are:
7,910,000,000,000,000,000,000,000 molecules
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Each molecule has a mass of:
0.0000000000000000000000299 gram
In scientific notation:
7.91 х 1024 molecules
2.99 х 10-23 gram
To Express a Number in Scientific Notation:
For small numbers (<1):
1) Locate the decimal point
2) Move the decimal point to the right to give a number
(coefficient) between 1 and 10
3) Write the new number multiplied by 10 raised to the
“nth power”
 where “n” is the number of places you moved the
decimal point. It has a negative sign
 If the decimal point is moved to the right, then the
exponent is a negative number (× 10-n)
To Express a Number in Scientific
For large numbers (>1):
1) Locate the decimal point
2) Move the decimal point to the left to give a number
(coefficient) between 1 and 10
3) Write the new number multiplied by 10 raised to the
“nth power”
 where “n” is the number of places you moved the
decimal point. It has a positive sign.
 If the decimal point is moved to the left, the
exponent is a positive number (× 10n)
Examples
 Write
each of the following in
scientific notation
 12,500
 0.0202
 37,400,000
 0.0000104
Examples
12,500
 Decimal place is at the far right
 Move the decimal place to a
position between the 1 and 2
 Coefficient (1.25)
 The decimal place was moved 4
places to the left (large number) so
exponent is positive
 1.25x104
Examples
 0.0202
 Move
the decimal place to a
position between the 2 and 0
 Coefficient (2.02)
 The decimal place was moved 2
places to the right (small number)
so exponent is negative
 2.02x10-2
Examples
 37,400,000
 Decimal
place is at the far right
 Move the decimal place to a
position between the 3 and 7
 Coefficient (3.74)
 The decimal place was moved 7
places to the left (large number) so
exponent is positive
 3.74x107
Examples
 0.0000104
 Move
the decimal place to a
position between the 1 and 0
 Coefficient (1.04)
 The decimal place was moved 5
places to the right (small number)
so exponent is negative
 1.04x10-5
Scientific Notation and
Calculators
1) Enter the coefficient (number)
2) Push the key: EE or EXP
Then enter only the power of 10
3) If the exponent is negative, use the
key: (+/-)
4) DO NOT use the multiplication key:
X
to express a number in sci. notation
Converting Back to a Standard Number
1) Determine the sign of the exponent, n
If n is + the decimal point will move to the
right (gives a number greater than one)
 If n is – the decimal point will move to the
left (gives a number less than one)
2) Determine the value of the exponent of 10
 The “power of ten” determines the
number of places to move the decimal
point
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Using Scientific Notation
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To compare numbers written in scientific
notation
First compare the exponents of 10
 The larger the exponent, the larger the
number
If the exponents are the same, then compare
coefficients directly
 Which number is larger?
21.8 х 103 or 2.05 х 104
2.18 х 104 > 2.05 х 104
Significant Figures:
Writing Numbers to Reflect Precision
 Two
kinds of numbers
 Counted (exact)
 Measured
Measured Numbers
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Scientific numbers are reported so that all
digits are certain except the last digit which
is estimated
A measurement:
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involves reading a measuring device
always has some amount of uncertainty
uncertainty comes from the tool used for
comparison
e.g. some rulers show smaller divisions (more
precise) than others
Measured Numbers
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Always have to
estimate the value
between the two
smallest divisions on
a measuring device
Every person will
estimate it slightly
differently, so there
is some uncertainty
present as to the
true value
2.8 cm
2.8 to 2.9 cm
2.9 cm
Significant Figures: Writing Numbers to
Reflect Precision
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Scientific numbers are reported so every
digit is certain except the last which is
estimated
To indicate the uncertainty of a single
measurement scientists use a system
called significant figures
Significant figures: All digits known with
certainty plus one digit that is uncertain
Counting Significant Figures
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The last digit written in a measurement is the
number that is considered to be uncertain
(estimated)
 Unless stated otherwise, the uncertainty in
the last digit is ±1
The precision of a measured quantity is
determined by number of sig. figures
A zero in a measurement may or may not be
significant
significant zeros
 place-holding zeros (not significant)
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Counting Significant Figures
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Nonzero integers are always significant
Zeros (may or may not be significant)
It is determined by its position in a
sequence of digits in a measurement
 Leading zeros never count as significant
figures
 Captive (interior) zeros are always
significant
 Trailing zeros are significant if the
number has a decimal point
Exact Numbers
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Exact numbers occur in definitions or in
counting
Numbers known with no uncertainty
Unlimited number of significant figures (never
limit the no. of sig. figures in a calculation)
They are either
 Counting
numbers
 7 pennies, 6 pills, 4 chairs
 Defined numbers
 12 in = 1 ft
 1 gal = 4 quarts
 1 minute = 60 seconds
Significant Figures in Calculations
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Calculations cannot improve the precision of
experimental measurements
The number of significant figures in any
mathematical calculation is limited by the least
precise measurement used in the calculation
Two operational rules to ensure no increase in
measurement precision
 addition and subtraction
 multiplication and division
Significant Figures in Calculations:
Multiplication and Division
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Product or quotient has the same number of
significant figures as the factor with the fewest
significant figures
Count the number of significant figures in each
number. The least precise factor (number) has
the fewest significant figures
Rounding
 Round the result so it has the same number
of significant figures as the number with the
fewest significant figures
Rounding
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To round the result to the correct number of
significant figures
If the last (leftmost) digit to be removed:
• is less than 5, the preceding digit stays the
same (rounding down)
• is equal to or greater than 5, the preceding
digit is rounded up
• In multiple step calculations, carry the extra
digits to the final result and then round off
Multiplication/Division Example:
5 SF
4 SF
3 SF
0.1021  0.082103  273 
1.1
2.1
 2.080438
2 SF
2 SF
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The number with the fewest
significant figures is 1.1 so the
answer has 2 significant figures
Multiplication/Division Example:
4 SF
5 SF
3 SF
3 SF
0.1021 × 0.082103 × 273 = 2.288481
2.29
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The number with the fewest
significant figures is 273 so the
answer has 3 significant figures
Significant Figures in Calculations:
Addition and Subtraction
 Sum
or difference is limited by the
quantity with the smallest number
of decimal places
 Find quantity with the fewest
decimal places
 Round answer to the same
decimal place
Addition/Subtraction Example:
1 d.p.
3 d.p.
2 d.p.
171 .5  72 .915  8 .23  236
.185
236.2
1 d.p.
 The
number with the fewest
decimal places is 171.5 so the
answer should have 1 decimal
place
Measurement
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The most useful tool of the chemist
Most of the basic concepts of chemistry
were obtained through data compiled by
taking measurements
How much…?
How long…?
How many...?
These questions cannot be answered
without taking measurements
The concepts of chemistry were
discovered as data was collected and
subjected to the scientific method
Measurement
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The estimation of the magnitude of an
object relative to a unit of measurement
Involves a measuring device
 e.g. meter stick, scale, thermometer
The device is calibrated to compare the
object to some standard (inch/centimeter,
pound/kilogram)
Quantitative observation with two parts:
A number and a unit
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Number tells the total of the quantity measured
Unit tells the scale (dimensions)
Measurement
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A unit is a standard (accepted) quantity
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Describes what is being added up
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Units are essential to a measurement
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For example, you need “six of sugar”
 teaspoons?
 ounces?
 cups?
 pounds?
The Basic Units of Measurement
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Units tells the magnitude of the standard
Two most commonly used systems of units of
measurement
 U.S. (English) system: Used in everyday
commerce (USA and Britain*)
 Metric system: Used in everyday commerce
and science (The rest of the world)
SI Units (1960): A modern, revised form of the
metric system set up to create uniformity of
units used worldwide (world’s most widely
used)
The Standard Units:
The Metric/SI System
A decimal system of measurement based
on the meter and the gram
 It has a single base unit per physical
quantity
 All other units are multiples of 10 of the
base unit
 The power (multiple) of 10 is indicated by
a prefix
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The Standard Units:
The Metric System
In the metric system there is one base unit for
each type of measurement
length
 volume
 mass
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The base units multiplied by the appropriate
power of 10 form smaller or larger units
The prefixes are always the same, regardless
of the base unit
 milligrams and milliliters both mean 1/1000 of
the base unit
The Standard Units: Length
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Meter
Base unit of length in metric and SI system
About 3 ½ inches longer than a yard
 1 m = 1.094 yd
The Standard Units: Length
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Other units of length
are derived from the
meter
Commonly use
centimeters (cm)
 1 m = 100 cm
 1 inch = 2.54 cm
(exactly)
The Standard Units: Volume
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Measure of the amount
of three-dimensional
space occupied by a
object
Derived from length
SI unit = cubic meter
(m3)
Metric unit = liter (L) or
10 cm3
Commonly measure
smaller volumes in
cubic centimeters (cm3)
Volume = side × side × side
Volume = side × side × side
The Standard Units: Volume
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Since it is a threedimensional measure,
its units have been
cubed
SI base unit = cubic
meter (m3)
This unit is too large
for practical use in
chemistry
Take a volume 1000
times smaller than the
cubic meter, 1dm3
The Standard Units: Volume
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Metric base unit =
1dm3 = liter (L)
1L = 1.057 qt
Commonly measure
smaller volumes in
cubic centimeters
(cm3)
Take a volume 1000
times smaller than the
cubic decimeter, 1cm3
The Standard Units: Volume
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Metric base unit =
1dm3 = liter (L)
1L = 1.057 qt
Commonly measure
smaller volumes in
cubic centimeters
(cm3)
Take a volume 1000
times smaller than the
cubic decimeter, 1cm3
The Standard Units: Volume
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The most commonly
used unit of volume in
the laboratory:
milliliter (mL)
1 mL = 1 cm3
1 L= 1 dm3 = 1000 mL
1 m3 = 1000 dm3 =
1,000,000 cm3
Use a graduated
cylinder or a pipette to
measure liquids in the
lab
The Standard Units: Mass
Measure of the total quantity of matter
present in an object
 SI unit (base) = kilogram (kg)
 Metric unit (base) = gram (g)
 Commonly measure mass in grams (g) or
milligrams (mg)
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kg = 1000 g
 1 g = 1000 mg
 1 kg = 2.205 pounds
 1 lb = 453.6 g
Prefixes Multipliers
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One base unit for each type of measurement
Length (meter), volume (liter), and mass (gram*)
The base units are then multiplied by the appropriate
power of 10 to form larger or smaller units
base unit = meter, liter, or gram
Prefixes Multipliers (memorize)
× base unit
Mega
 Kilo
 Base
 Deci
 Centi
 Milli
 Micro
 Nano
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(M) 1,000,000
(k)
1,000
1 meter liter gram
(d)
0.1
(c)
0.01
(m)
0.001
(µ)
0.000001
(n)
0.000000001
106
103
100
10-1
10-2
10-3
10-6
10-9
Prefix Multipliers
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For a particular measurement:
 Choose
the prefix which is similar in size
to the quantity being measured
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Keep in mind which unit is larger
 e.g.
A kilogram is larger than a gram, so
there must be a certain number of
grams in one kilogram
 Choose the prefix most convenient for a
particular measurement
n < µ < m < c < base < k < M
Converting from One Unit to Another:
Equalities
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A fixed relationship between two quantities
Shows the relationship between two units
that measure the same quantity
The relationships are exact, not measured
 1 min = 60 s
 12 inches = 1 ft
 1 dozen = 12 items (units)
 1L = 1000 mL
 16 oz = 1 lb
 4 quarts = 1 gallon
Converting from One Unit to Another:
Conversion Factors
Many problems in chemistry involve a
conversion of units
 Conversion factor: An equality expressed
as a fraction
 Used as a multiplier to convert a quantity
in one unit to its equivalent in another unit
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 May
be exact or measured
 Both parts of the conversion factor should
have the same number of significant figures
Solving Multistep Conversion Problems:
Dimensional Analysis Example
(Conversion Factors Stated within a Problem)
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The average person in the U.S.
consumes one-half pound of sugar per
day. How many pounds of sugar
would be consumed in one year?
1) State the initial quantity given (+unit):
One year
State the final quantity to find (+unit):
Pounds
2) Write a sequence of units (map) which
begins with the initial unit and ends with the
desired unit: year
day
pounds
1 cal  4.184 J
Solving Multistep Conversion Problems:
Dimensional Analysis Example
3) For each unit change,
State the equalities:
 Every equality will have two conversion
factors
365 days = 1 year
0.5 lb sugar =1day
year
day
pounds
Solving Multistep Conversion Problems:
Dimensional Analysis Example
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State the conversion factors:
0.5 lb. sugar
1day
and
1 day
0.5 lb. sugar
4) Set Up the problem:
1 year
365 day(s)
1 year
0.5 lb sugar
1 day
 183 lbs. sugar
Guide to Problem Solving when
Working Dimensional Analysis Problems
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Identify the known or given quantity and the
units of the new quantity to be determined
Write out a sequence of units which starts with
your initial units and ends with the desired units
(“solution map”)
Write out the necessary equalities and
conversion factors
Perform the mathematical operations that
connect the units
Check that the units cancel properly to obtain
the desired unit
Does the answer make sense?
Density
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The ratio of the mass of an object to the volume
occupied by that object
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Tells how tightly the matter within an object is packed
together
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Units for solids and liquids = g/cm3
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1 cm3 = 1 mL so can also use g/mL
Unit for gases = g/L
Density: solids > liquids >>> gases
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Density 
mass
volume
m
d 
v
Density
Can use density as a conversion factor
between mass and volume
 Density of some common substances given
in Table 2.4, page 33
 You will be given any densities on tests
EXCEPT water
 Density of water is 1.0 g/cm3 at room
temperature
 1.0 mL of water weighs how much?
 How many mL of water weigh 15 g?
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Density
 To
determine the density of an object
 Use a scale to determine the mass
 Determine the volume of the object
 Calculate
it if possible (cube shaped)
 Can also calculate volume by
determining what volume of water is
displaced by an object
Volume of Water Displaced = Volume of Object
Density Problem
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Iron has a density of 7.87 g/cm3. If 52.4 g
of iron is added to 75.0 mL of water in a
graduated cylinder, to what volume
reading will the water level in the cylinder
rise?
Vf  ?
m  52.4 g
d  7.87 g cm
Vi  7 5 .0 mL
3
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Density Problem
Solve for volume of iron
density 
mass
volume
52.4 g iron
volume 
mass
density
1 mL iron
7.87 g iron
1 cm
3
= 1 mL
 6 .658 mL iron

6 .658 mL iron
+ 75.0 mL water
= 8 1.7 mL total
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End