chapter one - georgia southern university-armstrong...
TRANSCRIPT
Chapter One
Matter and Measurement
http://www.chemistry.armstrong.edu/
nivens/course_list.htm
OWL HOMEWORK REQUIRED!!!
2
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.
3
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
4
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
5
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
6
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
7
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
8
Mixtures, Substances, Compounds, and Elements
• Mixtures
– two or more substances
– homogeneous mixtures –same phase
• Also termed “solutions”
– heterogeneous mixtures
9
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
10
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)
11
Units of Measurement
Definitions
• Length
– Measure of space in any direction
12
Units of Measurement
• Volume
– Amount of space
occupied by a system
– Derived unit, mL or cc, cm3
13
Measurements in Chemistry
QuantityQuantity UnitUnit SymbolSymbol
� length meter m
� mass kilogram kg
� time second s
� current ampere A
� temperature Kelvin K
� amt. substance mole mol
14
Measurements in ChemistryUse Metric System
NameName SymbolSymbol MultiplierMultiplier
• mega M 106 (1,000,000)
• kilo k 103 (1,000)
• deka da 10
• deci d 10-1 (0.1)
• centi c 10-2 (0.01)
15
Measurements in ChemistryMetric Prefixes
NameName SymbolSymbol MultiplierMultiplier
• milli m 10-3(0.001)
• micro µ 10-6(0.000001)
• nano n 10-9
• pico p 10-12
• femto f 10-15
16
Units of Measurement
Definitions
• Mass
–mass is the amount of matter
Weight
– force of gravitational attraction on mass
17
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
18
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
19
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
20
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
21
Use of Numbers
Significant Figures - Rules
• Non zero numbers are significant
• Leading zeroes are never significant0.000357 has three significant figures
• Imbedded zeroes are always significant3.0604 has five significant figures
• Trailing zeroes may be significantmust specify significance by how the number is written
1300 nails - counted or weighed?
1.30000 –what was the precision of the measurement?
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Use of Numbers
• Use scientific notation to remove doubtExpress 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
23
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
24
Use of Numbers
• Multiplication & Division rule
Easier of the two rules
Product has the smallest number of significant figures of multipliers
5.22 tooff round
21766.5
31.2x
224.4
25
Use of Numbers
• Multiplication & Division rule
Easier of the two rules
Product has the smallest number of significant figures of multipliers
5.22 tooff round
21766.5
31.2x
224.4
3.9 tooff round
89648.3
41.x
2783.2
26
Use of Numbers
• Addition & Subtraction rule
More subtle than the multiplication rule
Answer contains smallest decimal place of the addends.
6.95 tooff round
9463.6
20.2
423.1
3692.3
+
+
27
Use of Numbers
• Addition & Subtraction ruleMore subtle than the multiplication rule
Answer contains smallest decimal place of the addends.
6.95 tooff round
9463.6
20.2
423.1
3692.3
+
+
6.671 tooff round
6707.6
312.2
7793.8
−
28
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
29
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
30
Factor Label Method
• 1 inch = 2.54 cm
– Ratio
– Use ratio to perform a calculation Units will divide out.
• Convert 60 in to centimeters
in
cmor
cm
in
1
54.211
54.2
1==
cmin
cmxin 4.152
1
54.260 =
31
1 ft = 12 in becomes or
• Example - Express 9.32 yards in millimeters.
in 12
ft 1
ft 1
in 12
Factor Label Method
32
More practice
Express 9.32 yards in millimeters.
in 12
ft 1
ft 1
in 12
33
More Practice
)yd1
ft3( yd 9.32
mm ?yd 9.32 =
34
More practice
)ft1
in12( )
yd1
ft3( yd 9.32
mm ?yd 9.32 =
35
More Practice
)in1
cm2.54( )
ft1
in12( )
yd1
ft3( yd 9.32
mm ?yd 9.32 =
36
More Practice
mm 108.52)cm1
mm10( )
in1
cm2.54( )
ft1
in12( )
yd1
ft3( yd 9.32
mm ?yd 9.32
3×=
=
37
More Practice
mm 108.52)cm1
mm10( )
in1
cm2.54( )
ft1
in12( )
yd1
ft3( yd 9.32
mm ?yd 9.32
3×=
=
38
Factor Label Method
• Example - Express 627 milliliters in gallons.
39
The Unit Factor Method
• Example - Express 627 milliliters in gallons.
gal 0.166gal 0.166155gal ?
)qt4
gal1( )
L1
qt1.06( )
mL1000
L1( mL 627gal ?
mL 627gal ?
≈=
=
=
40
Practice
• Using your conversion handout
– Convert 25 g to lbs
– Convert 1 mL to Liters
– Convert 20 meters to cm
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.
• common mistake
)cm 2.54
in 1(cm102.61ft ? 242
×=
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.
2242 )cm 2.54
in 1(cm102.61ft ? ×=
43
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.
22242 )in 12
ft 1()
cm 2.54
in 1(cm102.61ft ? ×=
44
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.
343
3333
cm 107.39cm 73906.9696
)in 1
cm 2.54()
ft 1
in 12(ft 2.61cm ?
×≈=
=
45
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?
46
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?
8.81%
100%x ore g 335
Feg 29.5
100%x ore of grams
iron of grams iron % ?
=
=
=
47
Density and Specific Gravity
• What is density?
• Density (how compact something is, mass per unit volume)
• density = mass/volume
48
Density and Specific Gravity
• Example - Calculate the density of a substance if 742 grams of it occupies 97.3 cm3.
Vm density
mL 3.97cm 97.3 mL 1 cm 1 33
=
=∴=
49
Density and Specific Gravity
• Example - Calculate the density of a substance if 742 grams of it occupies 97.3 cm3.
g/mL 7.63 density
mL 97.3g 742
density
Vm density
mL 3.97cm 97.3 mL 1 cm 1 33
=
=
=
=∴=
50
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
51
Density and Specific Gravity
• Example -
mL 94.7 1.32
g 125V
density
mV
V
mdensity
mLg
==
=∴=
52
Density and Specific Gravity
• 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.
)water(density
)substance(densityGravity Specific =
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?
mL 4.32
g 31.10 Cr ofdensity
mL 4.32
mL 5.00 - mL 9.32 Cr of Volume
=
=
=
54
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?
20.7 1.00
7.20Cr ofGravity Specific
mLg 7.20
mLg 7.19907
mL 4.32
g 31.10 Cr ofdensity
mLg
mLg
==
≈=
=
55
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
56
Heat and Temperature
MP water BP water
• Fahrenheit 32 oF 212 oF
• Celsius 0.0 oC 100 cC
• Kelvin 273 K 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
57
Relationships of the Three Temperature Scales
273KC
or
273 C K
ipsRelationsh Centigrade andKelvin
o
o
−=
+=
58
1.8
32FC
or
32C 1.8F
1.85
9
10
18
100
180
ipsRelationsh Centigrade and Fahrenheit
oo
oo
−=
+×=
===
Example - Express 548 K in Celsius degrees.
•Example - Convert 211 oF to degrees Celsius.
59
Heat and Temperature
• Example - Convert 211oF to degrees Celsius.
C4.991.8
32112C
1.8
32FC
oo
oo
=−
=
−=
60
Heat and Temperature
• Example - Express 548 K in Celsius degrees.
275C
273485C
273KC
o
o
o
=
−=
−=