Transcript Introductory Chemistry, 2nd Edition Nivaldo Tro

Matters and Measurement
Edward Wen, PhD
Chemistry is about Everyday
experience
Photo credit: itsnicethat.com
• Why Cookies tastes different from
Cookie Dough?
• Why Baking Powder or Baking
Soda?
• Why using Aluminum Foil, not Paper
Towel?
• What if the Temperature is set too
high?
2
Chapter Outline
•
•
•
•
•
•
Classification of matters
Measurement, Metric system (SI)
Scientific Notation
Significant figures
Conversion factor
Density
3
In Your Room
• Everything you can see,
touch, smell or taste in
your room is made of
matter.
• Chemists study the
differences in matter and
how that relates to the
structure of matter.
4
What is Matter?
• Matter: anything that
occupies space and has
mass
• Matter is actually
composed of a lot of tiny
little pieces: Atoms and
Molecules
5
Atoms and Molecules
• Atoms: the tiny particles that
make up all matter. Helium
gas (for blimp) is made up of
Helium atoms.
• Molecules: In most
substances, the atoms are
joined together in units.
Liquid water is made up of
water molecules (2 Hydrogen
atoms + 1 Oxygen atoms)
6
Physical States of Matters
• Matter can be classified as solid, liquid or
gas based on what properties it exhibits
State
Shape
Volume
Compress
Flow
Solid
Fixed
Fixed
No
No
Liquid
Indef.
Fixed
No
Yes
Gas
Indef.
Indef.
Yes
Yes
7
Why different States of a Matter?
Structure Determines Properties
• the atoms or molecules have different
structures in solids, liquid and gases
8
Solids
• Particles in a solid: packed
close together and are fixed
in position
though they may vibrate
Incompressible
retaining their shape and
volume
Unable to flow
9
Liquids
• Particles are closely packed, but they
have some ability to move around
Incompressible
Able to flow, yet not to escape and
expand to fill the container (not
“antigravity”)
10
Gases
• The particles have complete freedom from
each other (not sticky to each other)
• The particles are constantly flying around,
bumping into each other and the container
• There is a lot of empty space between the
particles (low density)
 Compressible
 Able to flow and Fill space (“antigravity”)
11
Classifying Matter:
Sugar, Copper, Coke, Gasoline/Water
12
Classification of Matter
•Matter
Pure Substance
•Constant Composition
Mixture
•Variable Composition
•Homogeneous
13
Pure substance
Matter that is composed of only one kind of
piece.
Solid: Salt, Sugar, Dry ice, Copper, Diamond
Liquid: Propane, distilled water (or
Deionized water, DI water)
Gas: Helium gas (GOODYEAR blimp)
14
Classifying Pure Substances:
Elements and Compounds
Elements: Substances which can not be
broken down into simpler substances
by chemical reactions. (A,B)
Compounds: Most substances are
chemical combinations of elements. (C)
• Examples: Pure sugar, pure water
can be broken down into elements
Properties of the compound not
related to the properties of the
elements that compose it
15
Elements
• Example: Diamond (pure carbon), helium gas.
• 116 known, 91 are found in nature
 others are man-made
• Abundance = percentage found in nature
 Hydrogen: most abundant in the universe
 Oxygen: most abundant element (by mass) on earth and
in the human body
 Silicon: abundant on earth surface
• every sample of an element is made up of lots of
identical atoms
16
Compounds
• Composed of elements in fixed percentages
water is 89% O & 11% H
• billions of known compounds
• Organic (sugar, glycerol) or inorganic (table salt)
• same elements can form more than one different
compound
water and hydrogen peroxide contain just
hydrogen and oxygen
carbohydrates all contain just C, H & O (sugar,
starch, glucose)
17
Mixture
Matter that is composed of different
kinds of pieces. Different samples may
have the same pieces in different
percentages. (D)
Examples:
Solid: Flour, Brass (Copper and
Zinc), Rock
Liquid: Salt water, soda, Gasoline
Gas: air
18
Classification of Mixtures
• Homogeneous = composition is uniform
throughout
 appears to be one thing
 every piece of a sample has identical
properties, though another sample with
the same components may have
different properties
 solutions (homogeneous mixtures):
Air; Tap water
• Heterogeneous = matter that is nonuniform throughout
 contains regions with different
properties than other regions: gasoline
mixed with water; Italian salad
dressing
19
What is a Measurement?
• Quantitative observation
• comparison to an agreed upon standard
Every measurement has a number and a unit:
• 77 Fahrenheit: Room temperature
• 7.5 pounds: Average newborn body weight in the US:
• 55 ± 0.5 grams: amount of sugar in one can of Coca
Cola
UNIT: what standard you are comparing your object to
the number tells you
1. what multiple of the standard the object measures
2. the uncertainty in the measurement (±)
20
Some Standard Units in the
Metric System
Quantity
Measured
Name of
Unit
Abbreviation
Mass
gram
g
Length
meter
m
Volume
liter
L
Time
seconds
s
Temperature
Kelvin
K
21
Related Units in the SI System
All units in the SI system are related to the standard
unit by a power of 10 (exactly!)
• 1 kg = 103 g
• 1 km = 103 m
• 1 m = 102 cm
• The power of 10 is indicated by a prefix
• The prefixes are always the same, regardless of
the standard unit
22
Prefixes Used to Modify Standard Unit
• kilo = 1000 times base unit = 103
 1 kg = 1000 g = 103 g
• deci = 0.1 times the base unit = 10-1
 1 dL = 0.1 L = 10-1 L; 1 L = 10 dL
• centi = 0.01 times the base unit = 10-2
 1 cm = 0.01 m = 10-2 m; 1 m = 100 cm
• milli = 0.001 times the base unit = 10-3
 1 mg = 0.001 g = 10-3 g; 1 g = 1000 mg
• micro = 10-6 times the base unit
 1 m = 10-6 m; 106 m = 1 m
• nano = 10-9 times the base unit
 1 nL = 10-9L; 109 nL = 1 L
23
Common Prefixes in the SI System
Prefix
Symbol
Decimal
Equivalent
Power of 10
1,000,000
Base  106
1,000
Base  103
mega-
M
kilo-
k
deci-
d
0.1
Base  10-1
centi-
c
0.01
Base  10-2
milli-
m
0.001
Base  10-3
micro-
 or mc
0.000 001
Base  10-6
nano-
n
0.000 000 001 Base  10-9
24
Standard Unit vs. Prefixes
Using meter as example:
1 km = 1000 m = 103 m
1g
= 10 dm
= 100 cm
= 1000 mm
= 1,000,000 m
= 1,000,000,000 nm
= 102 cm
= 103 mm
= 106 m
= 109 nm
25
Length
• Two-dimensional distance an object covers
• SI unit: METER (abbreviation as m)
 About 3½ inches longer than a yard
1 m = 10-7 the distance from the North Pole to the
Equator
• Commonly use centimeters (cm)
 1 m = 100 cm = 1.094 yard
 1 cm = 0.01 m = 10 mm
 1 inch = 2.54 cm (exactly)
26
Mass
• Amount of matter present in an object
• SI unit: kilogram (kg)
 about 2 lbs. 3 oz.
• Commonly measure mass in grams (g) or
milligrams (mg)
 1 kg = 2.2046 pounds
(1 lbs. = 0.45359)
 1 g = 1000 mg = 103 mg
 1 g = 0.001 kg = 10-3 kg
27
Volume
• Amount of three-dimensional space occupied
• SI unit = cubic meter (m3)
• Commonly measure solid volume in cubic
centimeters (cm3)
 1 m3 = 106 cm3
 1 cm3 = 10-6 m3 = 0.000001 m3
• Commonly measure liquid or gas volume in
milliliters (mL)
 1 gallon (gal) = 3.78 L = 3.78  103 mL
 1 L = 1 dm3 = 1000 mL = 103 mL
 1 mL = 1 cm3 = 1 cc (cubic centimeter)
28
Common Everyday Units and Their
EXACT Conversions
11 cm
1 inch (in)
1 mile
1 foot
1 yard
=
=
=
=
2.54 cm
5280 feet (ft)
12 in
3 ft
29
Common Units and Their Equivalents
Mass
1 kilogram (km) = 2.205 pounds (lb)
1 pound (lb) = 453.59 grams (g)
1 ounce (oz) = 28.35 (g)
Volume
1 liter (L) = 1.057 quarts (qt)
1 U.S. gallon (gal) = 3.785 liters (L)
30
Units
• Always write every number with its
associated unit
• Always include units in your calculations
you can do the same kind of operations on units
as you can with numbers
• cm × cm = cm2
• cm + cm = cm
• cm ÷ cm = 1
using units as a guide to problem solving
31
Conversion Factor
• Relationships to Convert one unit of measurement to
another: US dollar  Canadian dollar, dollar  cent
• Conversion Factors: Relationships between two
units
 Both parts of the conversion factor have the same
number of significant figures
• Conversion factors generated from equivalence
statements
 e.g. 1 inch = 2.54 cm can give
or
2 .54 cm
1in
1in
2.54cm
32
How to Use Conversion Factor
• Arrange conversion factors so starting unit cancels
 Arrange conversion factor so starting unit is on the
bottom of the conversion factor
unit 1 x
unit 2
unit 1
= unit 2
Conversion Factor
33
We have been using the Conversion
Factor ALL THE TIME! 
• How are we converting #cents into #dollars? Why?
 From 1 dollar = 100 cents
45,000 cents x
1 dollar
dollar
= 450 dollars
100 cents
cents
Conversion Factor
34
Convert 325 mg to grams
Given: 325 mg
Find: ? g
Conv. Fact. 1 mg = 10-3 g Soln.
Map:
mg  g
0.325 g
35
Practice: How to set up
Conversion?
• To convert 5.00 inches to cm, from 1 in =
2.54 cm (exact), which one of the two
conversion factors should be used?
1 in
2.54 cm
or
2.54 cm
1 in
36
Practice: Conversion among Units
• 500 mg = ? g
500 mg 
* 8.0 in = ? m

0 .5 g

3780 mL
? mg
• 3.78 L = ? mL
• 1.2 nm = ? m
?g
3 . 78 L 
? mL
?L
1 . 2 nm 
?m

1 . 2  10
? nm
8 . 0 in 
? cm
? in

?m
? cm

9
nm
0 . 203 m
37
Scientific Notation
Very Large vs. Very Small numbers:
•The sun’s diameter is 1,392,000,000 m;
•An atom’s diameter is 0.000 000 000 3
m
Scientific Notation: 1.392 x 109 m & 3 x
10-10 m
the sun’s
diameter is
1,392,000,000 m
an atom’s
average diameter is
0.000 000 000 3 m
Scientific Notation (SN)
Power of 10 (Math language):
• 10 x 10 = 100  100 = 102 (2nd power of 10)
• 10 x 10 x 10 = 1,000  1,000 = 103 (3rd power of 10)
each Decimal Place in our number system represents a
different power of 10
• 24 = 2.4 x 101 = 2.4 x 10
• 1,000,000,000 (1 billion) = 109
• 0.0000000001 (1/10 billionth ) = 10-10
Easily comparable by looking at the power of 10
39
Exponents 10Y
exponent
1.23 x 10-8
decimal part
exponent part
1.23 x 105 > 4.56 x 102
4.56 x 10-2 > 7.89 x 10-5
7.89 x 1010 > 1.23 x 1010
• when the exponent on 10
(Y) is positive, the number
is that many powers of 10
larger
 sun’s diameter = 1.392 x
109 m = 1,392,000,000
m
• when Y is negative, the
number is that many
powers of 10 smaller
 avg. atom’s diameter =
3 x 10-10 m =
0.0000000003 m
40
Writing Numbers in SN
Big numbers:
12,340,000
1.234 x 107
Small numbers:
0.0000234
2.34 x 10-5
41
Writing a Number in Standard Form
1.234 x 10-6
• since exponent is -6, move the decimal point to the left 6
places
if you run out of digits, add zeros
000 001.234
0.000 001 234
If the exponent > 1, add trailing zeros:
1.234 x 1010
1.2340000000
12,340,000,000
42
Scientific calculators
43
Inputting Scientific Notation into a Calculator
-1.23 x 10-3
• input decimal part of the
number
 if negative press +/- key
• (–) on some
• press EXP key
 EE on some (maybe 2nd
function)
• input exponent on 10
 press +/- key to change
exponent to negative
Input 1.23
1.23
Press +/-
-1.23
Press EXP
-1.23 00
Input 3
-1.23 03
Press +/-
-1.23 -03
44
Significant Figures (Sig. Fig.)
Definition: The non-place-holding
digits in a reported measurement
some zero’s in a written number
are only there to help you locate
the decimal point
12.3 cm
has 3 sig. figs.
and its range is
12.2 to 12.4 cm
0.1230 cm
What is Sig. Fig. for?
has 4 sig. figs.
the range of values to expect for
and its range is
repeated measurements
0.1229 to 0.1231 cm
the more significant figures there
are in a measurement, the smaller
the range of values is
45
Counting Significant Figures
1.
All non-zero digits are significant
 1.5 :
2 Sig. Fig.s
2.
Interior zeros are significant
 1.05 :
3 Sig. Fig.s
3.
Trailing zeros after a decimal point are significant
 1.050 :
4 Sig. Fig.s. Leading zeros are NOT
significant
 0.001050 : 4 Sig. Fig.s Place-holding zero’s
= SN : 1.050 x 10-3
46
Counting Significant Figures (Contd)
4. Exact numbers has infinite () number of significant
figures:
example:
 1 pound = 16 ounces
 1 kilogram = 1,000 grams = 1,000,000 milligrams
 1 water molecule contains 2 hydrogen atoms
5. Zeros at the end of a number without a written decimal
point are ambiguous and should be avoided by using
scientific notation.
 Example: 150. has 3 sig. fig
 150 is ambiguous number
 1.50 x 102 has 3 sig. fig.
47
Example–Counting Sig. Fig. in a Number
How many significant figures are in each of the
following numbers?
 2 Sig. Fig. – leading zeros not sig.
 4 Sig. Figs – trailing & interior zeros
sig.
 2 sig. Figs, all digits sig.
27
 3 Sig. Figs – only decimal parts count
2.97 × 105
sig.
1 m = 1000 mm  both 1 and 1000 are exact numbers.
unlimited sig. figs.
0.0035
1.080
48
Practice:
How many Significant figures vs.
Decimal places?
• 2.2 cm
2 sig. Figs;
1 decimal place
3 sig. Figs;
• 2.50 cm 2 decimal places
49
Sig. Fig. in Multiplication/Division;
Rounding vs. Zeroing
• When multiplying or dividing measurements with
Sig. Fig., the result has the same number of
significant figures as the measurement with the
fewest number of significant figures
Rounding
• 5.02 × 89,665 × 0.10 = 45.0118 = 45
3 SF
5 SF
2 SF
2 SF
• 5.892 ÷
4 SF
6.10
3 SF
= 0.96590
= 0.966
3 SF
50
Sig. Fig. in Multiplication/Division:
Scientific notation
• Occasionally, scientific notation is needed to
present results with proper significant figures.
5.89 × 6,103 = 35946.67 = 3.59 × 104
51
Example: Determine the Correct
Number of Sig. Fig.
1. 1.01 × 0.12 × 53.51 ÷ 96 = 0.067556 = 0.068
3 SF
2 SF
4 SF
2 SF
2. 56.55 × 0.920 ÷ 34.684 = 1.5
4 SF.
3 SF.
6 SF.
result should
have 2 Sig. Fig.
= 1.50
result should
have 3 Sig. Fig.
52
Sig. Fig. in Addition/Subtraction
• when adding or subtracting measurements with
significant figures, the result has the same number of
decimal places as the measurement with the fewest
number of decimal places
5.74
+
2 dp
4.865 3 dp
0.823
3 dp
3.965
3 dp
+ 2.651 = 9.214 = 9.21
3 dp
= 0.9
2 dp
= 0.900
3 dp
53
Example: Determine the Correct Number of
Significant Figures
1. 0.987 + 125.1 – 1.22 = 124.867 = 124.9
3 dp
1 dp
2 dp
result should
have 1 dp
2. 0.764 – 3.449 – 5.315 = -8 = -8.000
3 dp
3 dp
3 dp
result should
have 3 dp
54
Sig. Fig. in Combined Calculations
• Do  and/or , then + and/or 3.489 – 5 .67 × 2.3
3 dp
3 Sig. Fig.
2 Sig. Fig.
=
3.489 –
3 dp
13
0 dp
= -9.511
= -10
0 dp (2 sig. fig.)
• Parentheses (): Do calculation in () first, then the rest
3.489 × (5.67 – 2.3)
2 dp
1 dp
= 3.489
×
4 Sig. Fig.
3.4
= 11.8628
2 Sig. Fig. 2 Sig. Fig.
= 12
55
Practice: Calculation with Proper
Significant Figures
a. 12.99 + 2.09 x 1.921 = 12.99 + 4.01 = 17.00
b. 2.00 x 3.5 - 1.000 =
c.
2 . 54  12 . 46

3 . 75
d. ( 0 . 0025  6 . 7 )  8 . 8 
7.0 – 1.000 = 6.0
15.00/3.75 = 4.00
6.7  8.8 = 59
56
How to solve Unit Conversion Problems
1) Write down Given Amount and Unit
2) Write down what you want to Find and Unit
3) Write down needed Conversion Factors or
Equations
4) Design a Solution Map for the Problem
 order Conversions to cancel previous units
or
 arrange Equation so Find amount is
isolated. Example: from Equation A = b  c
to solve for b
57
Solution Map for Unit Conversion
4) Apply the Steps in the Solution Map
 check that units cancel properly
 multiply terms across the top and divide
by each bottom term
Example:
2 . 54 g  12 . 46 g
3 . 1cm  2 . 5 cm

15 . 0 0 g
7 .7 5 cm
2
 1 . 9 g / cm
2
5) Check the Answer to see if its Reasonable
 correct size and unit
58
Example: Unit Conversion
Alternative Route:
Convert 7.8 km to miles
km
m
cm
Given: 7.8 km
Find: ? mi
Conv. Fact. 1 mi = 5280 ft
1 foot = 12 in
1 in = 2.54 cm (exact)
Soln. Map:
km  mi
in
ft
mi
60
Given: 7.8 km
Find: ? mi
Conv. Fact. 1 mi = 5280 ft
Alternative Route:
Convert 7.8 km to miles
1 foot = 12 in
1 in = 2.54 cm (exact)
Soln. Map:
km  mi
• Apply the Solution Map:
7.8 km 
1000 m
1 km

100 cm
1m

1 in
2.54 cm

1 ft
12 in

1 mi
 mi
5280 ft
= 4.84692 mi
• Sig. Figs. & Round:
= 4.8 mi
61
Temperature
• Temperature is a measure of the average
kinetic energy of the molecules in a sample
• Not all molecules have in a sample the
same amount of kinetic energy
• a higher temperature means a larger
average kinetic energy
62
Fahrenheit Temperature Scale
Two reference points:
• Freezing point of concentrated saltwater
(0°F)
• Average body temperature (100°F)
more accurate measure now set average body
temperature at 98.6°F
• Room temperature is about 75°F
63
Celsius Temperature Scale
Two reference points:
• Freezing point of distilled water (0°C)
• Boiling point of distilled water (100°C)
more reproducible standards
most commonly used in science
• Room temperature is about 25°C
64
Fahrenheit vs. Celsius
• a Celsius degree is 1.8 times larger than a
Fahrenheit degree
• the standard used for 0° on the Fahrenheit
scale is a lower temperature than the
standard used for 0° on the Celsius scale

F - 32
C 
1.8
F  1.8  C  32
65
The Kelvin Temperature Scale
• both the Celsius and Fahrenheit scales have “-”
numbers
• Kelvin scale is an absolute scale, meaning it
measures the actual temperature of an object
• 0 K is called Absolute Zero: all molecular motion
would stop, theoretically the lowest temperature in
the universe
 0 K = -273°C = -459°F
 Absolute Zero is a theoretical value
66
Kelvin vs. Celsius
• the size of a “degree” on the Kelvin scale is the
same as on the Celsius scale
though technically, call the divisions on the
Kelvin as kelvins, not degrees
that makes 1 K 1.8 times larger than 1°F
• the 0 standard on the Kelvin scale is a much lower
temperature than on the Celsius scale
K  C  273
67
Ext
remes
of Temperature
On the Earth,
• Lowest temperature recorded: -89.2°C (-128.6 °F,
184 K)
• Highest air temperature recorded: ~60°C (140 F)
In science lab,
• the highest temperature: 4 x 1012 K (?)
• the lowest temperature: ~10-10 K (?)
68
Conversion Between
Fahrenheit and Kelvin
Temperature Scales
Convert 104°F into Celsius
and Kelvin
Information
Given: 104 F
Find: ? °C, ? K
Eq’ns: K  273  C 1.8  C  32  F
• Fahrenheit to Celsius:
= 40 °C (keep 2 significant figures)
Celsius to Kelvin:
= 313 K
70
Mass & Volume
M ass
D en sity 
Mass & Volume:
V olu m e
• two main characteristics of matter
• even though mass and volume are individual
properties - for a given type of matter they
are related to each other!
 Density (ratio of mass vs. volume): for a
certain matter, its density is one of the
characteristic to distinguish from one
another
71
M ass
Unit for density
D en sity 
• Solids = g/cm3
V olu m e
1 cm3 = 1 mL
• Liquids = g/mL: Density of water = 1.00 g/mL
• Gases = g/L: Density of Air ~ 1.3 g/L
Volume of a solid can be determined by water
displacement
• Density : solids > liquids >>> gases
except ice and dry wood are less dense than
liquid water!
72
Density of Common Matters
73
Density
• Temperature affects the density: Heating
objects causes objects to expand, density
The Lava Lamp: heating/cooling
• In a heterogeneous mixture, the denser
object sinks
Why do hot air balloons rise?
The “Gold Rush”: Extracting gold
particle from sand
Density of gasoline changes over the
day!
74
Density and Volume
Styrofoam vs.
Quarter:
Both of these items
have a mass of 23
grams, but they
have very different
volumes; therefore,
their densities are
different as well.
75
Density and Buoyancy
•
•
•
•
Average density of human body = 1.0 g/cm3
Average density of sea water = 1.03 g/mL
Density of mercury, liquid metal, = 13.6 g/mL
Density of copper penny = 8.9 g/cm3
76
Density of Body and
Body Fat
Density of fat tissue < Density of
Muscle/Bones
Estimate the mass percentage of
body fat:

4 . 57
Body fat%  
 4 . 142
3
 Density ( g / cm )

  100 %

Average body fat%:
Female 28%, Male 22%
77
M
Using Density in Calculations
D V
Density 
Mass
Volume
• Both sides multiplied by Volume
Solution Maps:
m, V
D
 M a ss  D e n sity  V o lu m e
m
V, D
• Both sides divided by Density
Mass
Volume 
Density

m, D
V
78
Application of Density
A man gives a woman an engagement ring and tells
her that it is made of platinum (Pt).
Critical thinking : test to determine the ring’s
density before giving him an answer about marriage.
Data: She places the ring on a balance and finds it
has a mass of 5.84 grams. She then finds that the
ring displaces 0.556 cm3 of water.
Density Pt = 21.4 g/cm3
79
Test results
Given: Mass = 5.84 grams
Volume = 0.556 cm3
Density Pt = 21.4 g/cm3
Find: Density in grams/cm3
m
D
V
5.84 g
0.556 cm
3
 10.5
g
cm
3
80
Density as a Conversion Factor
• Between mass and volume!!
Density H2O = 1 g/mL \
• 1 g H2O  1 mL H2O
Density lead = 11.3 g/cm3
• 11.3 g lead  1
cm3
lead
1.00 g
or
1 mL
1 mL
1.00 g
11.3 g
1 mL
or
1 mL
11.3 g
• How much does 4.0 cm3 of Lead weigh?
4.0 cm3 Pb x
11.3 g Pb
1 cm3 Pb
= 45 g Pb
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Measurement and Problem Solving
Density as a Conversion Factor
• The gasoline in an automobile gas tank has a mass of 60.0 kg
and a density of 0.752 g/cm3. What is the volume?
• Given: 60.0 kg
• Find: Volume in L
• Conversion Factors:
 0.752 grams/cm3
 1000 grams = 1 kg
82
Measurement and Problem Solving
Density as a Conversion Factor
• kg  g  cm3
60.0 kg 
1000 g
1 kg

1 cm
3
4
 7.98  10 cm
3
0.752 g
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Example:
A 55.9 kg person displaces
57.2 L of water when
submerged in a water tank.
What is the density of the
person in g/cm3?
Information:
Given: m = 5.59 x 104 g
Find: density, g/cm3
Solution Map: m,VD
Equation: D  m
V
Volume = 57.2 L = 5.72 x 104 cm3
D 
m
V

5 . 59 x 10
5 . 72 x 10
4
4
g
cm
3
= 0.977 g/cm3
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Practice: Calculation involving
Density
1. The density of air at room temperature and sea
level is 1.29 g/L. Calculate the mass of air in a
5.0-gal bottle (1 gal = 3.78 L).
KEY: 24 g (2SF)
2. A driver filled 15.60 kg of gasoline into his car.
If the density of gasoline is 0.788 g/mL, what is
the volume of gasoline in liters?
KEY: 19.8 L (3SF)
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About Challenging Problems
1.99+: Proper dosage of a drug is 3.5 mg/kg of
body weight. Calculate the milligrams of this drug
for a 138-lb individual? (1 lb = 454 g).
KEY: 2.2×102 mg (2SF)
1.103: 100. mg ibuprofen/5 mL Motrin. Calculate
the grams of ibuprofen in 1.5 teaspoons of Motrin.
(1 teaspoon = 5.0 mL)
KEY: 0.15 g (2SF)
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