Title: CHAPTER ONE
1CHAPTER ONE
- The Foundations of Chemistry
2Chapter Outline
- Matter and Energy
- States of Matter
- Chemical and Physical Properties
- Chemical and Physical Changes
- Mixtures, Substances, Compounds, and
Elements - Measurements in Chemistry
- Units of Measurement
3Chapter Outline
- Use of Numbers
- The Unit Factor Method (Dimensional Analysis)
- Percentage
- Density and Specific Gravity
- Heat and Temperature
- Heat Transfer and the Measurement of Heat
4Matter and Energy - Vocabulary
- Chemistry
- Science that describes matter its properties,
the changes it undergoes, and the energy changes
that accompany those processes - Matter
- Anything that has mass and occupies space.
- Energy
- The capacity to do work or transfer heat.
- Scientific (natural) law
- A general statement based the observed behavior
of matter to which no exceptions are known.
5Natural Laws
- Law of Conservation of Mass
- Law of Conservation of Energy
- Law of Conservation of Mass-Energy
- Einsteins Relativity
- Emc2
6Scientific Method
- Observation
- Hypothesis
- Observation or experiment
- Theory
- Observation or experiment
- Law
7States of Matter
8States of Matter
9States of Matter
10States of Matter
- Change States
- heating
- cooling
11States of Matter
- Illustration of changes in state
- requires energy
12Chemical and Physical Properties
- Chemical Properties - chemical changes
- rusting or oxidation
- chemical reactions
- Physical Properties - physical changes
- changes of state
- density, color, solubility
- Extensive Properties - depend on quantity
- Intensive Properties - do not depend on quantity
13Mixtures, Substances, Compounds, and Elements
- Substance
- matter in which all samples have identical
composition and properties - Elements
- substances that cannot be decomposed into simpler
substances via chemical reactions - Elemental symbols
- found on periodic chart
14Mixtures, Substances, Compounds, and Elements
15Mixtures, Substances, Compounds, and Elements
- Compounds
- substances composed of two or more elements in a
definite ratio by mass - can be decomposed into the constituent elements
- Water is a compound that can be decomposed into
simpler substances hydrogen and oxygen
16Mixtures, Substances, Compounds, and Elements
17Mixtures, Substances, Compounds, and Elements
- Mixtures
- composed of two or more substances
- homogeneous mixtures
- heterogeneous mixtures
18Measurements in Chemistry
- Quantity Unit Symbol
- length meter m
- mass kilogram kg
- time second s
- current ampere A
- temperature Kelvin K
- amt. substance mole mol
19Measurements in ChemistryMetric Prefixes
- Name Symbol Multiplier
- mega M 106
- kilo k 103
- deka da 10
- deci d 10-1
- centi c 10-2
20Measurements in ChemistryMetric Prefixes
- Name Symbol Multiplier
- milli m 10-3
- micro ? 10-6
- nano n 10-9
- pico p 10-12
- femto f 10-15
21Units of Measurement
- Definitions
- Mass
- measure of the quantity of matter in a body
- Weight
- measure of the gravitational attraction for a
body
22Units 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
- See Table 1-7 for more conversion factors
23Use of Numbers
- Exact numbers
- 1 dozen 12 things for example
- Accuracy
- how closely measured values agree with the
correct value - Precision
- how closely individual measurements agree with
each other
24Use of Numbers
- Significant figures
- digits believed to be correct by the person
making the measurement - Measure a mile with a 6 inch ruler vs. surveying
equipment - Exact numbers have an infinite number of
significant figures - 12.000000000000000 1 dozen
- because it is an exact number
25Use of Numbers
- Significant Figures - Rules
- Leading zeroes are never significant
- 0.000357 has three significant figures
- Trailing zeroes may be significant
- must specify significance by how the number is
written - 1300 nails - counted or weighed?
- Use scientific notation to remove doubt
- 2.40 x 103 has ? significant figures
26Use 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
27Use of Numbers
- Piece of Black Paper with rulers beside the
edges
28Use of Numbers
- Piece of Paper Side B enlarged
- How long is the paper to the best of your ability
to measure it?
29Use of Numbers
- Piece of Paper Side A enlarged
- How wide is the paper to the best of your ability
to measure it?
30Use of Numbers
- Determine the area of the piece of black paper
using your measured values. - Compare your answer with your classmates.
- Where do your answers differ in the numbers?
- Significant figures rules for multiplication and
division must help us determine where answers
would differ.
31Use of Numbers
- Multiplication Division rule
- Easier of the two rules
- Product has the smallest number of significant
figures of multipliers
32Use of Numbers
- Multiplication Division rule
- Easier of the two rules
- Product has the smallest number of significant
figures of multipliers
33Use of Numbers
- Multiplication Division rule
- Easier of the two rules
- Product has the smallest number of significant
figures of multipliers
34Use of Numbers
- Determine the perimeter of the piece of black
paper using your measured values. - Compare your answer with your classmates.
- Where do your answers differ in the numbers?
- Significant figures rules for addition and
subtraction must help us determine where answers
would differ.
35Use of Numbers
- Addition Subtraction rule
- More subtle than the multiplication rule
- Answer contains smallest decimal place of the
addends.
36Use of Numbers
- Addition Subtraction rule
- More subtle than the multiplication rule
- Answer contains smallest decimal place of the
addends.
37Use of Numbers
- Addition Subtraction rule
- More subtle than the multiplication rule
- Answer contains smallest decimal place of the
addends.
38The Unit Factor Method
- Simple but important method to get correct
answers in word problems. - Method to change from one set of units to
another. - Visual illustration of the idea.
39The Unit Factor Method
- Change from a to a by obeying the
following rules.
40The Unit Factor Method
- Change from a to a by obeying the
following rules. - Must use colored fractions.
41The Unit Factor Method
- Change from a to a by obeying the
following rules. - Must use colored fractions.
- The box on top of the fraction must be the same
color as the next fractions bottom box.
42The Unit Factor Method
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43The Unit Factor Method
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44The Unit Factor Method
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45The Unit Factor Method
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46The Unit Factor Method
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47The Unit Factor Method
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48The Unit Factor Method
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49The Unit Factor Method
- colored fractions represent unit factors
- 1 ft 12 in becomes or
- Example 1-1 Express 9.32 yards in millimeters.
50The Unit Factor Method
51The Unit Factor Method
52The Unit Factor Method
53The Unit Factor Method
54The Unit Factor Method
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55The Unit Factor Method
- Example 1-2 Express 627 milliliters in gallons.
- You do it!
56The Unit Factor Method
- Example 1-2. Express 627 milliliters in gallons.
57The Unit Factor Method
- Area is two dimensional thus units must be in
squared terms. - Example 1-3 Express 2.61 x 104 cm2 in ft2.
58The Unit Factor Method
- Area is two dimensional thus units must be in
squared terms. - Example 1-3 Express 2.61 x 104 cm2 in ft2.
- common mistake
59The Unit Factor Method
- Area is two dimensional thus units must be in
squared terms. - Example 1-3 Express 2.61 x 104 cm2 in ft2.
O
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60The Unit Factor Method
- Area is two dimensional thus units must be in
squared terms. - Example 1-3 Express 2.61 x 104 cm2 in ft2.
O
R
R
61The Unit Factor Method
- Area is two dimensional thus units must be in
squared terms. - Example 1-3 Express 2.61 x 104 cm2 in ft2.
62The Unit Factor Method
- Area is two dimensional thus units must be in
squared terms. - Example 1-3 Express 2.61 x 104 cm2 in ft2.
63The Unit Factor Method
- Volume is three dimensional thus units must be in
cubic terms. - Example 1-4 Express 2.61 ft3 in cm3.
- You do it!
64The Unit Factor Method
- Volume is three dimensional thus units must be in
cubic terms. - Example 1-4 Express 2.61 ft3 in cm3.
65Percentage
- Percentage is the parts per hundred of a sample.
- Example 1-5 A 335 g sample of ore yields 29.5 g
of iron. What is the percent of iron in the ore? - You do it!
66Percentage
- Percentage is the parts per hundred of a sample.
- Example 1-5 A 335 g sample of ore yields 29.5 g
of iron. What is the percent of iron in the ore?
67Density and Specific Gravity
- density mass/volume
- What is density?
- Why does ice float in liquid water?
68Density and Specific Gravity
- density mass/volume
- What is density?
- Why does ice float in liquid water?
69Density and Specific Gravity
- Example 1-6 Calculate the density of a substance
if 742 grams of it occupies 97.3 cm3.
70Density and Specific Gravity
- Example 1-6 Calculate the density of a substance
if 742 grams of it occupies 97.3 cm3.
71Density and Specific Gravity
- Example 1-7 Suppose you need 125 g of a corrosive
liquid for a reaction. What volume do you need? - liquids density 1.32 g/mL
- You do it!
72Density and Specific Gravity
- Example 1-7 Suppose you need 125 g of a corrosive
liquid for a reaction. What volume do you need? - liquids density 1.32 g/mL
73Density and Specific Gravity
- Example 1-7 Suppose you need 125 g of a corrosive
liquid for a reaction. What volume do you need? - liquids density 1.32 g/mL
74Density and Specific Gravity
- Waters density is essentially 1.00 at room T.
- Thus the specific gravity of a substance is very
nearly equal to its density. - Specific gravity has no units.
75Density and Specific Gravity
- Example 1-8 A 31.0 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? - You do it
76Density and Specific Gravity
- Example1-8 A 31.0 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?
77Density and Specific Gravity
- Example1-8 A 31.0 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?
78Density and Specific Gravity
- Example 1-9 A concentrated hydrochloric acid
solution is 36.31 HCl and 63.69 water by mass.
The specific gravity of the solution is 1.185.
What mass of pure HCl is contained in 175 mL of
this solution? - You do it!
79Density and Specific Gravity
80Density and Specific Gravity
81Density and Specific Gravity
82Try This One
- Battery acid is 40.0 sulfuric acid and 60 water
by mass. Its specific gravity is 1.31.
83Solution
- From the given specific value number 1.31
- We may write Density 1.31g/mL
- Next
84- The solution is 40 sulfuric acid and 60 water
by massfrom this information we get
Because 100g of solution contains 40.0 g of
sulfuric acid
sulfuric acid
solution
85We can now solve the problem
86Heat 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
87Heat 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
88Heat and Temperature
- MP water BP water
- Fahrenheit 32 oF 212 oF
- Celsius 0.0 oC 100 cC
- Kelvin 273 K 373 K
89Relationships of the Three Temperature Scales
90Relationships of the Three Temperature Scales
91Relationships of the Three Temperature Scales
92Relationships of the Three Temperature Scales
- Easy method to remember how to convert from
Centigrade to Fahrenheit. - Double the Centigrade temperature.
- Subtract 10 of the doubled number.
- Add 32.
93Heat and Temperature
- Example 1-10 Convert 211oF to degrees Celsius.
94Heat and Temperature
- Example 1-11 Express 548 K in Celsius degrees.
95Heat Transfer and The Measurement of Heat
- SI unit J (Joule)
- calorie
- Amount of heat required to heat 1 g of water 1 oC
- 1 calorie 4.184 J
- Calorie
- Large calorie, kilocalorie, dietetic calories
- Amount of heat required to heat 1 kg of water 1
oC - English unit BTU
- Specific Heat
- amount of heat required to raise the T of 1g of a
substance by 1oC - units J/goC
96Heat Transfer and the Measurement of Heat
- Heat capacity
- amount of heat required to raise the T of 1 mole
of a substance by 1oC - units J/mol oC
- Example 1-12 Calculate the amt. of heat to raise
T of 200.0 g of water from 10.0oC to 55.0oC
97Heat Transfer and the Measurement of Heat
- Heat transfer equation
- necessary to calculate amounts of heat
- amount of heat amount of substance x
specific heat x?DT
98Heat Transfer and the Measurement of Heat
- Heat transfer equation
- necessary to calculate amounts of heat
- amount of heat amount substance x specific
heat x?DT
99Heat Transfer and the Measurement of Heat
- Heat transfer equation
- necessary to calculate amounts of heat
- amount of heat amount substance x specific
heat x??T
100Heat Transfer and the Measurement of Heat
- Example 1-13 Calculate the amount of heat to
raise the temperature of 200.0 grams of mercury
from 10.0oC to 55.0oC. Specific heat for Hg is
0.138 J/g oC. - You do it!
101Heat Transfer and the Measurement of Heat
- Example 1-13 Calculate the amount of heat to
raise T of 200.0 g of Hg from 10.0oC to 55.0oC.
Specific heat for Hg is 0.138 J/g oC. - Requires 30.3 times more heat for water
- 4.184 is 30.3 times greater than 0.138
102Heating Curve for 3 Substances
Which substance has the largest specific heat?
Which substances T will decrease the most after
the heat has been removed?
103Heating Curve for 3 Substances