Chapter One
Matter and Measurement
http://www.chemistry.armstrong.edu/
nivens/course_list.htm
OWL HOMEWORK REQUIRED!!!
Matter and Energy • Chemistry
– the study of matter and its changes
– the "central science"
Matter
– Anything that has mass and occupies space. (mass is
the amount of matter, weight is force of gravitational
attraction on the mass)
• Energy
– The ability to do work or transfer heat.
• Scientific (natural) law
– A general statement based the observed behavior of
matter to which no exceptions are known.
2
Natural Laws
• Law of Conservation of Mass-
amount of matter
(mass) is a constant under a chemical or physical change
• Law of Conservation of Energy –
total energy in a
process is fixed, only converted from one form to another.
• Law of Conservation of Mass-Energy- Universal
sum of mass + energy is a constant
• Einstein’s Relativity
• E=mc2
3
States of Matter
• Solid – definite shape and volume
• Liquid – definite volume (assumes
shape of container)
• Gas – no shape or volume (expand to
fill volume of container
•Change of States
–Heating or cooling
4
Chemical and Physical Properties
• Chemical Properties - observed during a
chemical change (change in composition)
– rusting or oxidation
– chemical reactions
• Physical Properties – observed as a physical
state a compound resides in
– changes of state
– density, color, solubility
• Extensive Properties - depend on quantity
• Intensive Properties - do not depend on
quantity
5
Mixtures, Substances, Compounds,
and Elements
• Substance
– identical composition and properties
• Elements
– cannot be decomposed into simpler
substances via chemical reactions
• Elemental symbols
– found on periodic chart
6
Mixtures, Substances, Compounds,
and Elements
• Compounds
– two or more elements in a definite ratio by
mass - can be decomposed into the
constituent elements
• Water– hydrogen and oxygen
• Carbon dioxide – carbon and oxygen
7
Mixtures, Substances, Compounds,
and Elements
• Mixtures
– two or more substances
– homogeneous mixtures –same phase
• Also termed “solutions”
– heterogeneous mixtures
8
Math Review and Measurements
• Make measurements to understand the
environment
– Humans—sight, taste, smell, sound…
• Limited and biased
– Use instruments—meter sticks,
thermometers, balances
• More accurate and precise
– All measurements have units—METRIC
SYSTEM vs. British System
9
Units of Measurement
Definitions
• Time
– Interval or duration of forward events
• Mass
– measure of the quantity of
matter in a body
• Weight
– measure of the gravitational attraction for a body
(w=mg)
10
Units of Measurement
Definitions
• Length
– Measure of space in any direction
11
Units of Measurement
• Volume
– Amount of space
occupied by a system
– Derived unit, mL or cc, cm3
12
Measurements in Chemistry
Quantity
length
mass
time
current
temperature
amt. substance
Unit
meter
kilogram
second
ampere
Kelvin
mole
Symbol
m
kg
s
A
K
mol
13
Measurements in Chemistry
Use Metric System
Name
• mega
• kilo
• deka
• deci
• centi
Symbol
M
k
da
d
c
Multiplier
106 (1,000,000)
103 (1,000)
10
10-1 (0.1)
10-2 (0.01)
14
Measurements in Chemistry
Metric Prefixes
Name
• milli
• micro
• nano
• pico
• femto
Symbol
m
µ
n
p
f
Multiplier
10-3(0.001)
10-6(0.000001)
10-9
10-12
10-15
15
Units of Measurement
Definitions
• Mass
– mass is the amount of matter
Weight
– force of gravitational attraction on mass
16
Units of Measurement
Common Conversion Factors
• Length
– 1 m = 39.37 inches
– 2.54 cm = 1 inch
• Volume
– 1 liter = 1.06 qt
– 1 qt = 0.946 liter
17
Use of Numbers
• Exact numbers
– 1 dozen = 12 things for example
• Measured Numbers—
– use rules for significant figures
– Use scientific notation were possible
• Accuracy
– how closely measured values agree with the
correct value
• Precision
– how closely individual measurements agree with
each other
18
Use of Numbers
• Significant figures
– digits believed to be correct by the person
making the measurement
• Exact numbers have an infinite number
of significant figures
12.000000000000000 = 1 dozen
because it is an exact number
19
Significant figures
• Calculators give 8+ numbers
• People estimate numbers differently
• Scientists develop rules to help you determine
which digits are “significant”
• Dictated by the precision (graduation) on
your measuring device
• In the lab, the last significant digit is the digit
you have to estimate
20
Use of Numbers
Significant Figures - Rules
• Non zero numbers are significant
• Leading zeroes are never significant
0.000357 has three significant figures
• Imbedded zeroes are always significant
3.0604 has five significant figures
• Trailing zeroes may be significant
must specify significance by how the number is
written
1300 nails - counted or weighed?
1.30000 –what was the precision of the
measurement?
21
Use of Numbers
• Use scientific notation to remove doubt
Express answers as powers of 10 by moving the
decimal place right (-) or left (+)
2000 2 x 103
15000 1.5 x 10?
0.004 4 x 10-3
0.000053 __.__ x 10?
In scientific notation, zeros are given if they are
significant
1.000 x 103 has 4 significant figures
2.40 x 103 has ? significant figures
22
Use of Numbers
• -Scientific notation for logarithms
take the log of 2.40 x 103
log(2.40 x 103) = 3.380
How many significant figures?
• -Imbedded zeroes are always significant
3.0604 has five significant figures
23
Use of Numbers
• Multiplication & Division rule
Easier of the two rules
Product has the smallest number of
significant figures of multipliers
4.242
x 1.23
5.21766
round off to 5.22
24
Use of Numbers
• Multiplication & Division rule
Easier of the two rules
Product has the smallest number of
significant figures of multipliers
2.7832
4.242
x 1.23
x 1.4
5.21766
3.89648
round off to 3.9
round off to 5.22
25
Use of Numbers
• Addition & Subtraction rule
More subtle than the multiplication rule
Answer contains smallest decimal place of the
addends.
3.6923
+ 1.234
+ 2.02
6.9463
round off to 6.95
26
Use of Numbers
• Addition & Subtraction rule
More subtle than the multiplication rule
Answer contains smallest decimal place of
the addends.
3.6923
+ 1.234
+ 2.02
6.9463
round off to 6.95
8.7937
− 2.123
6.6707
round off to 6.671
27
Units of Measurement
Common Conversion Factors (Equalities)
• Length
– 1 m = 39.37 inches
– 2.54 cm = 1 inch
• Volume
– 1 liter = 1.06 qt
– 1 qt = 0.946 liter
• See Text and Handout for more
conversion factors
28
Using Conversion factors
• Scientists often must convert between units.
• Conversion factors can be made for any
relationship.
• Use known equivalence to make a fraction
that can be used to “convert” from one unit
to the other.
– Fraction equals 1, so you are multiplying by 1
29
Factor Label Method
•
1 inch = 2.54 cm
– Ratio
1in
=1
2.54 cm
or 1 =
2.54 cm
1in
– Use ratio to perform a calculation Units will
divide out.
•
Convert 60 in to centimeters
60 in x
2.54 cm
= 152.4 cm
1in
30
Factor Label Method
1 ft = 12 in becomes
1 ft
12 in
or
12 in
1 ft
• Example - Express 9.32 yards in
millimeters.
31
More practice
Express 9.32 yards in millimeters.
1 ft
12 in
12 in
1 ft
32
More Practice
9.32 yd = ? mm
9.32 yd (
3ft
)
1yd
33
More practice
9.32 yd = ? mm
3ft 12in
9.32 yd ( ) (
)
1yd 1ft
34
More Practice
9.32 yd = ? mm
3ft 12in 2.54cm
9.32 yd (
)(
)(
)
1yd 1ft
1in
35
More Practice
9.32 yd = ? mm
3ft 12in 2.54 cm 10mm
9.32 yd (
)(
)(
)(
) = 8.52 × 10 3 mm
1yd 1ft
1in
1cm
36
More Practice
9.32 yd = ? mm
9.32 yd (
3ft 12in 2.54 cm 10mm
)(
)(
)(
) = 8.52 × 10 3 mm
1yd 1ft
1in
1cm
37
Factor Label Method
• Example - Express 627 milliliters in
gallons.
38
The Unit Factor Method
• Example - Express 627 milliliters in
gallons.
? gal =627 mL
? gal = 627 mL (
1L
1.06 qt 1gal
)(
)(
)
1000 mL
1L
4qt
? gal = 0.166155 gal ≈ 0.166 gal
39
Practice
•
Using your conversion handout
– Convert 25 g to lbs
– Convert 1 mL to Liters
– Convert 20 meters to cm
40
Area – length x width
Area is two dimensional thus units must
be in squared terms.
• Example 1-4: Express 2.61 x 104 cm2 in
ft2.
? ft 2 = 2.61× 104 cm2 (
1 in
)
2.54 cm
• common mistake
41
Area – length x width
Area is two dimensional thus units must
be in squared terms.
• Example 1-4: Express 2.61 x 104 cm2 in
ft2.
1 in 2
? ft 2 = 2.61×10 4 cm 2 (
)
2.54 cm
42
Area – length x width
Area is two dimensional thus units must
be in squared terms.
• Example 1-4: Express 2.61 x 104 cm2 in
ft2.
? ft 2 = 2.61 × 104 cm 2 (
1 in 2 1 ft 2
) (
)
2.54 cm 12 in
43
Volume – length x width x height
• Volume is three dimensional thus units
must be in cubic terms—this volume is
used in medical measurements--cc
• Example 1-5: Express 2.61 ft3 in cm3.
? cm 3 = 2.61 ft 3 (
12 in 3 2.54 cm 3
) (
)
1 ft
1 in
= 73906.9696 cm 3 ≈ 7.39 × 10 4 cm 3
44
Percentage
• Percentage is the parts per hundred of
a sample. = (g / g total ) X 100%
• Example - A 335 g sample of ore yields
29.5 g of iron. What is the percent of
iron in the ore?
45
Percentage
• Percentage is the parts per hundred of a
sample.
• Example - A 335 g sample of ore yields 29.5 g
of iron. What is the percent of iron in the ore?
grams of iron
x 100%
grams of ore
29.5 g Fe
=
x 100%
335 g ore
= 8.81%
? % iron =
46
Density and Specific Gravity
• What is density?
• Density (how compact something is, mass per
unit volume)
• density = mass/volume
47
Density and Specific Gravity
• Example - Calculate the density of a
substance if 742 grams of it occupies
97.3 cm3.
1 cm 3 = 1 mL ∴ 97.3 cm 3 = 97.3 mL
density = m
V
48
Density and Specific Gravity
• Example - Calculate the density of a
substance if 742 grams of it occupies
97.3 cm3.
1 cm 3 = 1 mL ∴ 97.3 cm 3 = 97.3 mL
density = m
V
742
g
density =
97.3 mL
density = 7.63 g/mL
49
Density and Specific Gravity
• Example - You need 125 g of a
corrosive liquid for a reaction. What
volume do you need?
– liquid’s density = 1.32 g/mL
50
Density and Specific Gravity
• Example -
m
m
∴V =
V
density
125 g
V=
= 94.7 mL
1.32 g mL
density =
51
Density and Specific Gravity
Specific Gravity =
density(substance)
density(water)
• Water’s density is essentially 1.00 g/mL at
room T.
• Thus the specific gravity of a substance is
very nearly equal to its density.
• Specific gravity has no units.
52
Density and Specific Gravity
• Example1-9: A 31.10 gram piece of chromium is
dipped into a graduated cylinder that contains
5.00 mL of water. The water level rises to 9.32
mL. What is the specific gravity of chromium?
Volume of Cr = 9.32 mL - 5.00 mL
= 4.32 mL
31.10 g
density of Cr =
4.32 mL
53
Density and Specific Gravity
• Example1-9: A 31.10 gram piece of chromium is
dipped into a graduated cylinder that contains
5.00 mL of water. The water level rises to 9.32
mL. What is the specific gravity of chromium?
density of Cr =
31.10 g
4.32 mL
= 7.19907 g
Specific Gravity of Cr =
mL
≈ 7.20 g
7.20 g mL
= 7.20
1.00 g mL
mL
54
Heat and Temperature
• Heat and Temperature
are not the same thing
T is a measure of the
intensity of heat in a
body
• 3 common temperature
scales - all use water as
a reference
55
Heat and Temperature
MP water
32 oF
0.0 oC
273 K
• Fahrenheit
• Celsius
• Kelvin
BP water
212 oF
100 cC
373 K
• Body temperature 37.0 oC, 98.6 oF
– 37.2 oC and greater—sick
– 41 oC and greater, convulsions
– <28.5 oC hypothermia
56
Relationships of the Three
Temperature Scales
Kelvin and Centigrade Relationships
K = o C + 273
or
o
C = K − 273
57
Fahrenheit and Centigrade Relationsh ips
180 18 9
=
= = 1.8
100 10 5
o
F = 1.8 × o C + 32
or
o
o
C=
F − 32
1.8
•Example - Convert 211 oF to degrees Celsius.
Example - Express 548 K in Celsius degrees. 58
Heat and Temperature
• Example - Convert 211oF to degrees
Celsius.
o
F − 32
1.8
2 11 − 32
o
C =
= 99 . 4 o C
1.8
o
C =
59
Heat and Temperature
• Example - Express 548 K in Celsius
degrees.
o
C = K − 273
o
C = 5 48 − 273
o
C = 275
60