[Chapter One: The Science of Chemistry ] [2011]
Chapter One
The Science of Chemistry
PREREADING ACTIVITY
1. Prior to reading take a look at the picture on page 2. What do you think this is a picture of?
Consider the title of the chapter. Do you think the picture if a good match for the title?
Please explain your answer.
2. Now, read the first paragraph in bold print at the top of page 3. Based on this paragraph,
tell me in your own words, what you expect to learn in this chapter.
3. What are properties? Can you give me some general properties that would describe you.
4. What does it mean to classify something? If I were going to classify the shoes of the
students in this classroom, what are some ways I might go about doing this?
SECTION 1: WHAT IS CHEMISTRY? pg. 4-9
In chemistry, we study changes. These changes can be either physical or chemical. These
changes are covered in Section 1-1 of your book. As you read, complete the maps on the
following pages to help your understanding of the material.
Page 1 of 67
[Chapter One: The Science of Chemistry ] [2011]
Changes of Matter
Type
Type
Book Definition
Book Definition
How to Identify
Examples
1.
1.
2.
2.
3.
3.
How to Identify
Examples
1.
1.
2.
2.
3.
3.
4.
4.
Page 2 of 67
[Chapter One: The Science of Chemistry ] [2011]
A physical property is observed with the senses and can be determined without destroying the
object. For example, color, shape, mass, length and odor are all examples of physical properties.
A chemical property indicates how a substance reacts with something else. The original substance
is fundamentally changed in observing a chemical property. For example, the ability of iron to rust
is a chemical property. The iron has reacted with oxygen, and the original iron metal is changed. It
now exists as iron oxide, a different substance.
Classify the following properties as either chemical or physical by putting a check in the appropriate
column.
Physical Property
Chemical Property
Blue color
Density
Flammability
Solubility (Dissolve)
Reacts with acid to form
hydrogen
Color changes from blue
to red
Supports combustion
Melting point
Reacts with water to
form gas
Reacts with a base to
form water
Smells bad as it sours
Hardness (does it break
easily?)
Boiling point
Can neutralize a base
Luster (shiny?)
Smells like lemons
Page 3 of 67
[Chapter One: The Science of Chemistry ] [2011]
PHYSICAL VS. CHEMICAL CHANGES
In a physical change, the original substance still exists. It has only changes in
form. In a chemical change, a new substance is produced. Energy changes
always accompany chemical changes.
Classify the following as being a physical or chemical change and circle the
clue word(s) .
1. Sodium hydroxide dissolves in water.
2. Hydrochloric acid reacts with potassium hydroxide to produce a salt, water
and heat.
3. A pellet of sodium is sliced in two.
4. Water is heated and changed to steam
5. Potassium chlorate decomposes to potassium chloride and oxygen gas.
6. Iron rusts to form iron oxide.
7. When placed in H2O, a sodium pellet catches on fire as hydrogen gas is
liberated and sodium hydroxide is formed.
8. Evaporation
9. Ice melting
10. Milk sours
11. Sugar dissolves in water.
12. Wood rotting
13. Pancakes cooking on a griddle
14. Grass growing in a lawn
15. A tire is inflated with air.
16. Food is digested in the stomach.
Page 4 of 67
[Chapter One: The Science of Chemistry ] [2011]
CLASSIFICATION ACTIVITY I
Consider the list of substances below.
Step 1: Your job is to group these substances into four categories based on similarities and
differences. You must use five categories. No more and no less.
Step 2: Write a brief description for each category, explaining how you decided what belongs in
that category.
Step 3: Come up with an appropriate title or name for each category.
Paint
Mercury
Water
Air
Italian Dressing
Baking Soda
Iron
Sand in Water
Milk
Jello
Tap Water
Rust (Fe2O3)
Sugar Water
Page 5 of 67
Types of
Matter
Semantic
Map
Can be separated by density
Can be separated through boiling
Created through physical means
Differences can be seen with the
naked eye
Created through a chemical
reaction
Can be separated through filtration
Can be separated through physical
means
Can be separated through chemical
means
Mixture
Pure
Contains 2 or more elements in a
definite ratio
Contains 2 or more elements in
which the amount can vary
Contains only one type of atom
[Chapter One: The Science of Chemistry ] [2011]
TYPES OF MATTER
Element
Compound
Heterogeneous
Mixture
Homogeneous
Mixture
Solution
Suspension
Colloid
Page 6 of 67
[Chapter One: The Science of Chemistry ] [2011]
SECTION 3 HOW IS MATTER CLASSIFIED? Pg. 21 - 28
Construct the following flow chart as work through the reading. Each unattached box is provided so that you can fill in a brief definition
Types of Matter
Matter
Page 7 of 67
[Chapter One: The Science of Chemistry ] [2011]
SECTION 2: MEASUREMENT (PG. 10-19)
Complete the following word sort as you read pages 10-19 in your textbook:
Main/Broad Concepts:
(may be used more than once)
Base Measurements
Chemical Properties
Derived Units
Physical Properties
Qualitative Measurements
Quantitative Measurements
Smaller/Specific Concepts:
Colorless
Combines
Decomposes
Density
Ductile
Height
Length
Mass
Oxidizes
Reacts
Volume
Page 8 of 67
[Chapter One: The Science of Chemistry ] [2011]
PRACTICING THE METRIC SYSTEM
K
H
D B D
C
M
Convert the following from the given
unit to the requested unit:
1)
2)
3)
4)
5)
6)
7)
8)
9)
360 g to
0.00238 cg to
13.52 g to
0.014 kg to
43.25 cg to
641.5 mg to
281 ml to
4.305 l to
28.5 ml to
mg
10)
1.832 L to
ml
11)
6.58 cm to
mm
12)
18.05 m to
cm
13)
3.80 km to
m
14)
14.28 m to
km
15)
35.85 cm to
km
16)
40.6 dm to
mm
17)
1.05 mm to
cm
18)
8.75 mm to
g
kg
cg
mg
g
l
ml
l
_________m
Page 9 of 67
[Chapter One: The Science of Chemistry ] [2011]
DENSITY PROBLEMS
Be sure to include units in all answers.
Reminders:
1cm3 =1mL
Density = Mass  Volume
1. What is the density of a material if its mass is 2.02 g and its volume is .500cm3? (4.04)
2. What is the density a substance that has a mass of 87.6g and a volume of 8.09cm3?
3. A substance has a mass of 61.9g and its volume is 5.46cm3, what is this density? (11.3)
4. What's the density of a sample that has a mass of 75.4g and a volume of 5.24cm3?
5. What is the density of a material when the mass is 20.4g and the volume is 18.6cm3?
(1.10)
6. A sample has a volume of 10.2cm3 and a mass of 15.6g. What is the density of the gold?
7. A sample has a mass of 1.02g and a volume of 1.35cm3, what is the density of the nickel?
(0.75)
8. What is the density of a substance that has a mass of 54.2g and a volume of 3.06cm3?
9. What is the density of a substance when its volume is 51.6cm3 and its mass is 134.5g?
(2.6)
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[Chapter One: The Science of Chemistry ] [2011]
MORE DENSITY PROBLEMS
1. What is the volume of a 898 kilogram piece of iron metal (density is 7.75 g/mL)?(115,871)
2. Calculate the mass of a liquid with a density of 3.2 g/mL and a volume of 25 mL.
3. An irregular object with a mass of 18 kg displaces 2.5 L of water when placed in a large
overflow container. Calculate the density of the object. (7.2)
4. A graduated cylinder has a mass of 80 g when empty. When 20 mL of water is added, the
graduated cylinder has a mass of 100 g. If a stone is added to the graduated cylinder, the
water level rises to 45 mL and the total mass is now 156 g. What is the density of the
stone?
5. An object has a mass of 49 kilograms and a volume of 93.2 ml. What is the density? (0.53)
6. If object has a density of 1.12 g/ml and a mass of 96 g. What volume does it occupy?
7. The volume of an object is 492 ml. Its mass is 675 grams. What is the density of this object?
(1.37)
8. If an object has a volume of 100 cm3 and a density of 2.92 g/cm3, what is its mass?
9. An object is placed in a graduated cylinder which previously held 20.0 ml of water. The water level
rose to 39.5 ml. After removing the object and drying it, the student placed the object on the
balance, which read 203 grams. What is the density of this object? (10.4)
10. Calculate the mass of an object, which has a volume of 125 ml and a density of 9.5 g/ml.
11. An empty beaker weighs 243 grams. A liquid is added to the beaker and the mass increases to
493 grams. Calculate the volume of the liquid if the density is 0.70 g/ml. (357)
Page 11 of 67
[Chapter One: The Science of Chemistry ] [2011]
Density Worksheet
1. Define mass?
2. Define volume?
3. Define density and show the formula for calculating density.
4. Why does changing the shape of an object have no effect on the density of that
object?
5. Aluminum is used to make airplanes. Cast iron is used to make weightlifting
equipment. Explain why the densities of these metals make them useful for these
purposes?
6. Calculate the densities of the following objects. Remember to place units after
each number.
Object A length = 6cm width = 3cm height = 1cm mass = 36g
volume = ___________ density = ___________
Object B length = 10cm width = 5cm height = 2cm mass = 300g
volume = ___________ density = _____________
Object C Use the water displacement method to determine the density of object C (silly
putty).
initial water level in graduated cylinder = 25ml
final water level after placing silly putty into graduated cylinder = 29ml
mass of silly putty=8g
volume = ____________ density = ___________
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[Chapter One: The Science of Chemistry ] [2011]
States of Matter (pgs. 4-9)
1. Particles in the solid ice cube appeared to move while staying within their rigidly
defined positions. These particles have_____________________ motion only.
2. From what you observed, applying pressure to a liquid causes its molecules to move
____________________ . If enough pressure is applied, the liquid becomes a(n)
__________________.
3. ______ Solids have the following properties with oneexception. Choose the property that
does not apply to a solid.
a. definite shape
c. orderly arrangement of particles
b. easy to compress
d. definite volume
4. Explain why ice is an example of a solid.
5. Hexane is a liquid. Explain why hexane is also a fluid.
6. How do the shapes and sizes of the graduated cylinder, Erlenmeyer flask, and
measuring cup affect the volume of the 50 mL hexane sample? How does the size of a container
affect the volume of a gas?
Worksheet 1.1
7. Compare and contrast particle motion for the three states of matter using the particle
views in Sections a through c as examples.
8. What happens to the motion of gas particles when they are compressed by a piston?
9. List three of the properties of a liquid.
Page 13 of 67
[Chapter One: The Science of Chemistry ] [2011]
METRIC MATCH GAME
Purpose: In this activity you will use your knowledge of the metric system to construct a 4 X 4
rectangle (4 cards down and 4 cards across) by correctly matching up the measurement
conversions on the given Metric Match cards.
Background: Prior to the time of Napoleon, every country in the world had its own measuring
system, which was often based a body part of the ruling monarch (hence the term ruler). The units
of length might have been based on the length of a king’s foot (thus the measurement of foot).
Weight units might have been called stones fro the number of equal size stones it took to counterbalance a king on a huge balance. Not only did every county have its own measuring system, but
the system might have changed from monarch to monarch.
Napoleon used cannons and cannon shells to conquer most of Europe. He didn’t want to run his
supply too thin and therefore used locally produced cannon balls and artillery supplies. Since
every country had different measuring systems, he found it impossible to have the conquered
countries manufacture his needed supplies. He therefore ordered his scientists back in Paris to
come up with a measuring system that would be easy to understand and could be used by all
people in any country he occupied. He could then have all his supplies made locally to his
specifications. The system the French scientists came up with was the metric system and was
based upon divisions of 10 using the same prefixes for length, mass and volume.
While the United States have not fully adopted the metric system, scientists in all countries rely on
the metric system to communicate clearly without having to convert from one country’s system to
another’s.
Instruction:
1. Note that each card has four sides, each with a different value.
2. You must construct a grid which will be 4 cards wide and 4 cards tall.
3. Each side of each card must be touching another card with an equal value. (For example,
the side of one card has 7 liters printed on it, while another card has 700 centiliters. These
two sides can touch each other, because 7 liters is equal to 700 centiliters. In another
example, a card whose side reads 20 milligrams can touch another card whose side reads
0.020 grams (20 mg = 0.20 g).
4. Continue to place cards together appropriately until all cards are in their proper place.
Remember, every side of every card must be matched correctly.
5. When your grid is complete, call the teacher for a signature.
Teacher’s Signature
Page 14 of 67
[Chapter One: The Science of Chemistry ] [2011]
Vocabulary: CHAPTER ONE
 Chemical reaction
 Reactants
 Products
 Matter
 Volume
 Physical Property
 Chemical Property
 Quantitative
 Qualitative
 Atom
 Pure Substance
 Mass
 Molecule
 Weight
 Element
 Physical Change
 Compound
 Chemical Change
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[Chapter One: The Science of Chemistry ] [2011]
 Mixture
 Colloid
 Homogeneous
 Suspension
 Heterogeneous
 Density
Page 16 of 67
[Chapter One: The Science of Chemistry ] [2011]
Two Types of Data:
Qualitative: _______________________________________________________________________________________________
Quantitative: _____________________________________________________________________________________________
Matter: __________________________________________________________________________________________________________
Type of Matter: _______________
1. Arrangement:
2. Movement:
Type of Matter:
Type of Matter: _______________
_______________
1. Arrangement:
1. Arrangement:
2. Movement:
2. Movement:
Fixed:
State
Solid
Indefinite:
Shape
Volume
Compress
Flow
Liquid
Gas
Physical Properties: ______________________________________________________
 Examples of Physical Properties
Chemical Properties: ______________________________________________________
 Examples of Chemical Properties
Page 17 of 67
[Chapter One: The Science of Chemistry ] [2011]
Physical Changes (∆)
 Definition: __________________________________________________
 Things to look for:
o
o
o
 Examples of Physical Changes (∆)
Chemical Changes (∆)
 Definition: _______________________________________________
 Things to look for to indicate a CHEMICAL CHANGE
o
o
o


o

Examples of Chemical Changes (∆)
Classifying Matter
Pure Substances : _________________________________________________
 Examples
Mixtures: ________________________________________________________
 Example
Metric System
Common Base Units

Length:

Mass:

Volume:
Page 18 of 67
[Chapter One: The Science of Chemistry ] [2011]
Prefixes
K
H
D
B
D
C
M
Density
What is density? __________________________
Calculating Density
Page 19 of 67
[Chapter Two: Matter & Energy ] [2011]
Chapter Two
Matter and Energy
Warmup Activity
For each of the following activities, identify whether it is matter, not matter or
not sure
.
Peanut butter
The human brain
Water
Music
Fish
Carbon dioxide
Light
Air
Garbage
Yourself
Time
An idea
Motion
Tree
Heat
Temperature
FOCUS ACTIVITY
1. If you had to define the word “energy” for someone, how would you explain it? List all the
sources of energy that you can think of.
2. What two types of energy did you learn about in physics? Please define these.
3. You body requires energy to work. Where does this energy come from? What unit do we
use when talking about food energy?
4. Where does the energy that runs your house come from? Explain the steps required to
get this energy from its source to your house.
Page 20 of 67
[Chapter Two: Matter & Energy ] [2011]
1
MODULE 2Phase Change Diagram
25
G
20
Temperature (celsius)
15
E
10
F
D
5
0
B
-5
C
A
-10
-15
Make sure you use the terms kinetic energy, potential energy, molecules moving faster,
molecules moving slower, molecules moving farther apart, molecules moving closer
together.
1. What is the freezing point (in oC)?
2. What is the boiling point (in oC)?
3. What letter represents where the
solid is being warmed?
4. What letter represents where the
liquid is being warmed?
5. What letter represents where the
vapor is being warmed?
6. What letter represents the melting of
the solid?
7. What letter represents the
vaporization of the liquid?
8. What letter(s) show a change in
potential energy?
11. What letter(s) show the molecules
moving farther apart?
12. When going from B to C what
happens?
13. When going from A to D what
happens?
14. When going from F to A what
happens?
15. When going from G to E what
happens?
9. What letter(s) show a change in
kinetic energy?
10. What letter(s) show the molecules
moving faster?
Page 21 of 67
[Chapter Two: Matter & Energy ] [2011]
SECTION 1: ENERGY (PG. 38 TO 45)
Examples/Evidence:
Examples/Evidence
Always involved in two types
of:
Energy
Is
transferred
by:
Can be either
Define:
Define:
Examples:
Examples:
Page 22 of 67
[Chapter Two: Matter & Energy ] [2011]
EXOTHERMIC VS. ENDOTHERMIC Venn Diagram (pg. 40)
Compare and contrast
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[Chapter Two: Matter & Energy ] [2011]
Graphic Organizer: Changes in State (pg. 378)
Types of Matter
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[Chapter Two: Matter & Energy ] [2011]
Section: Energy (pg. 38-45)
Complete each statement below by writing the correct term in the space provided.
1. The capacity to do work is __________________.
2. A change in matter from one form to another without a change in chemical
properties is a ________________ change.
3. A change that occurs when one or more substances change into a new
substances with different properties is a ____________________ change.
4. A change in matter in which energy is absorbed is a(n) __________________
process.
5. A change in matter in which energy is released is a(n) _________________
process.
6. Energy must be added to a solid to melt it. This addition gives the particles
______________ energy, allowing them to move out of the crystalline structure.
7. To freeze a substance, energy must be ________________ from the substance.
Write the answers to the following questions in the space provided.
8. State the law of conservation of energy.
9. What is heat?
10. Define temperature.
11. What is the difference between heat and temperature.
Determine whether the process is EXOTHERMIC or ENDOTHERMIC.
1.
2.
3.
4.
5.
6.
Freezing water
Evaporation of water
Condensing water
Baking bread
Nuclear fission where energy is released
Magnesium and hydrochloric acid is mixed in a beaker, the beaker feels
warm to the touch
7. Melting ice cubes
8. A candle flame
Page 25 of 67
[Chapter Two: Matter & Energy ] [2011]
Phase Change word sort
Organize the following terms under the two categories listed below.
Boiling
Change of state
Gases
Liquid
Moving closer together
Moving farther apart
Moving faster
Moving slower
Solid
Temperature changes
No temperature change
_________________________________________________________________________________________
KINETIC ENERGY CHANGE
POTENTIAL ENERGY CHANGE
Page 26 of 67
[Chapter Two: Matter & Energy ] [2011]
q = heat energy
Q= m C∆T
m = mass
T = temperature
∆T = (Tfinal – Tinitial)
c= specific heat
1. A 15.75-g piece of iron absorbs 1086.75 joules of heat energy, and its temperature changes from
25°C to 175°C. Calculate the specific heat of iron.
2. How many joules of heat are needed to raise the temperature of 10.0 g of aluminum from 22°C to
55°C, if the specific heat of aluminum is 0.90 J/g°C?
3. To what temperature will a 50.0 g piece of glass raise if it absorbs 5275 joules of heat and its
specific heat capacity is 0.50 J/g°C? The initial temperature of the glass is 20.0°C.
4. Calculate the specific heat of a piece of wood if 1500.0 g of the wood absorbs 6.75×104 joules of
heat, and its temperature changes from 32°C to 57°C.
5. 100.0 mL of 4.0°C water is heated until its temperature is 37°C. If the specific heat of water is
4.18 J/g°C, calculate the amount of heat energy needed to cause this rise in temperature.
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[Chapter Two: Matter & Energy ] [2011]
6. 25.0 g of mercury is heated from 25°C to 155°C, and absorbs 455 joules of heat in the process.
Calculate the specific heat of mercury.
7. What is the specific heat of silver metal if 55.00 g of the metal absorbs 47.3 calories of heat and
the temperature rises 15.0°C?
8. If a sample of chloroform is initially at 25°C, what is its final temperature if 150.0 g of chloroform
absorbs 1000 joules of heat, and the specific heat of chloroform is 0.96 J/g°C?
Page 28 of 67
[Chapter Two: Matter & Energy ] [2011]
Ethanol 101
What is ethanol?
Gasoline is not the only fuel that can power cars! Cars can actually be powered by corn, sugar cane,
wheat, potatoes, cellulosic materials and other feedstocks (unprocessed natural products used in
manufacturing) that are converted into ethanol. Ethanol is a clear, grain alcohol produced from
renewable sources that can be used as fuel. Ethanol (CH3CH2OH) is an organic compound that is
considered an alcohol since it has a hydroxyl group (OH) attached to a carbon.
Pure ethanol is generally mixed with a percentage of gasoline when used as a motor fuel. There are
two common types of ethanol that are used in vehicles- E10 and E85. E10 is a blend of 10% ethanol
and 90% gasoline. All cars are capable of using E10 as fuel. E85, on the other hand, is 85% ethanol
and 15% gasoline. E85 is considered an “alternative fuel” for use in Flexible Fuel Vehicles (FFV’s).
These vehicles can either run on E85 or regular gasoline.
Ethanol is not a new fuel source. Henry Ford’s Model T was designed to run on ethanol.
How is ethanol made?
In simple terms, ethanol is made by fermenting starch or sugar from feedstock, and then distilling it
into alcohol. Most of the ethanol produced in the US is distilled from corn. In Brazil, sugar cane is
used to make ethanol. Ethanol can also be made from cellulosic biomass. Cellulosic materials
include corn stalks and husks, wheat and barley straw, rice or sugar cane bagasse (fibre from
crushing the sugar cane), willow and polar trees, switchgrass, and municipal waste. Cellulosic
ethanol is made the same way; however, the sugars in cellulose are more complex than those found
in corn, and need to be separated carefully for the process to work and require more effort to break
down. The ethanol produced by cellulose is the same as that made from corn.
How much corn does it take to make ethanol?
According to the Department of Energy, it takes one bushel of corn to make 2.5 gallons of ethanol.
Since only part of the corn kernel is used to make the ethanol, one bushel of corn also yields;


1.6 lbs of corn oil
10 lbs of high protein feed


2.6 lbs of corn meal
31.5 lbs of starch used for beverages and
sweeteners
According to the American Coalition for Ethanol, 2.81 billion gallons of ethanol were produced in
2003. Over the past few years, ethanol production has increased:



2004- 12% of nation’s corn crop was used to produce 3.4 billion gallons of ethanol
2005- 14% of the nation’s corn crop was used to produce 4 billion gallons of ethanol
2006- an estimated 20% of the nation’s crop was used to produce about 5 billion gallons of ethanol
Approximately 50 million gallons of ethanol are used to make E85.
Does ethanol have a positive net energy balance?
A fuel source has a positive net energy balance if it produces more energy than is required to make
it. The energy needed to create ethanol from corn includes the energy required to grow the corn,
and the energy required to convert the corn into alcohol, which most likely involves the use of fossil
fuels such as oil or coal.
According to the Department of Energy, ethanol has a positive net energy balance, that means that
it provides about 25% more energy than is used to make it, including growing the corn, harvesting
Page 29 of 67
[Chapter Two: Matter & Energy ] [2011]
it, and distilling it into alcohol. It takes about 0.74 million Btu of “Fossil Energy Input” to yield 1
million Btu of ethanol.
Why use ethanol as a fuel?
There are many environmental and economic benefits to using ethanol as a fuel source. Ethanolblended fuel has a high oxygen content, which enables it to burn more completely and create less
pollution. It reduces the carbon monoxide and hydrocarbon emissions that come from a vehicle’s
tailpipe. Likewise, according to the Department of Energy’s Argonne National Laboratory, ethanolblended fuels reduced the CO2 equivalent greenhouse gas emissions by 7.8 million tons in 2005.
That’s like removing the greenhouse gas emissions of over one million cars from the road.
Increasing the use of ethanol-blended fuels produced in the U.S. helps to reduce the country’s
dependence on oil and foreign suppliers. The American Coalition for Ethanol reports that one barrel
of ethanol (42 gallons) can displace 1.2 barrels of petroleum at the refinery. While the 4 billion
gallons of ethanol produced in 2005 equals only about 3% of the total U.S. gas consumption for a
year ( about 140 billion gallons) there’s potential for ethanol. The government has suggested a
proposal to help reduce present gasoline consumption by blending it with ethanol and other
alternative fuels. The plan seeks for a 20% reduction in gasoline usage by 2017. Improvements in
production, technology, and awareness may yield higher results in the future.
Is there a downside to ethanol?
The production and use of renewable ethanol-blended fuels is increasing. However, the number of
production facilities and supplies or availability of E85 are currently limited. In addition, while
ethanol is cost-efficient to produce, there is a drawback to E85- there is approximately a 10%
decrease in energy per mile for E85 relative to gasoline, meaning that a car does not get as many
miles from a gallon of E85 as it does from a gallon of gasoline.
Answer the following questions based on the information about ethanol.
1. What is the difference between E10 and E85?
2. In the U.S., what is the primary source for ethanol production? In Brazil?
3. If it takes one bushel of corn to make 2.5 gallons of ethanol, how many bushels did it take to make
ethanol produced in the U.S. in 2005?
4. What does it mean that ethanol “has a positive net energy balance”?
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[Chapter Two: Matter & Energy ] [2011]
5. Is ethanol energy efficient to produce? Why or why not?
6. What do you consider to be the pros and cons of using ethanol as a fuel?
Pros
Cons
Page 31 of 67
[Chapter Two: Matter & Energy ] [2011]
HEAT OF VAPORIZATION/FUSION PROBLEMS
Q = H* m
Hvaporization = 2260 J/g
H fusion = 333 J/g
m = mass (g)
1. Calculate the energy required to melt 250 grams of ice.
2. Determine the energy released when 1000 grams of steam is condensed. (2,260,000)
3. If 975 grams of water is frozen, how much energy is released?
4. I am trying to vaporize 225 grams of iodine. How much energy will I need, if the heat of
vaporization for iodine is 0.163 kJ/gram? (36,675)
5. 49.5 grams of water is boiled. How much energy will it take until all of the water is
vaporized?
6. How much energy is required to melt 100 grams of iron, if the heat of fusion for iron is 0.25
kJ/gram? (25,000)
7. A 21.0 g sample of water is cooled from 34°C to 28°C. How many joules of heat were
removed from the water? (527.2)
Page 32 of 67
[Chapter Two: Matter & Energy ] [2011]
8. An 18.7 g sample of platinum metal increases in temperature by 2.3°C when 5.7 joules of
heat are added. What is the specific heat of platinum?
9. The specific heat of silver is 0.24 J/g°C. How much heat must be added to a silver block of
mass 86 grams to raise its temperature by 9.0°C? (185.76)
10. How much heat is required to raise the temperature of 68.0 grams of AlF3 from 25°C to
80.0°C? The specific heat of AlF3 is 0.8948 J/g°C.
11. How much heat is required to raise the temperature of 789 grams of acetic acid, CH3COOH,
from 25°C to 82.7°C? The specific heat of acetic acid is 2.57 J/g°C. (117,000)
12. How much heat is released when 42.8 grams of calcium carbide cools from 74.2°C to
11.5°C? The specific heat for calcium carbide is 0.982 J/g°C.
13. In order to make tea, 1000 grams of water is heated from 22° to 99°C. How much energy is
needed? (322,168)
Page 33 of 67
[Chapter Two: Matter & Energy ] [2011]
MULTISTEP PROBLEMS
1. Calculate the amount of energy required to take 125 grams of H2O from -15C to 45C.
2. Determine how much energy is release when 67 grams of water at 0C, is frozen, then
cooled to 25C. (25,611)
3. I have heated 50 grams of ice from -11C, until it is steam at 125C. How much energy did
this require?
4. How much energy will I need to heat 112 grams of water from 50C to 150C? (287,862)
5. If I cool 18 grams of H2O from 200C to -25C, will I be releasing or absorbing energy?
Calculate the amount of energy.
Page 34 of 67
[Chapter Two: Matter & Energy ] [2011]
•
What is Energy?
– Energy definition:
__________________________________________________
– Examples:
________________________________________________________
– Energy ____________ is always involved when there is a change in
matter (whether physical or chemical changes in matter)
–

Energy is never __________________________________________- just
transferred
Two Types of Energy
– Kinetic
 Energy of motion
 In chemistry this refers to
______________________________________________________
– Potential
 Stored energy or energy of position
 In chemistry this refers to
______________________________________________________

Energy gets transferred by:
– Light, Chemicals, Motion (Mechanical), Electrical, Sound or Heat
– Heat:
____________________________________________________________
– Heat Energy is always transferred in a very specific way:
____________________________________________________________

What is Temperature?
– The measurement of kinetic energy of a substance
(___________________________________________________________)
Page 35 of 67
[Chapter Two: Matter & Energy ] [2011]

What do the particles look like in each phase of matter?
Solid
Gas
Liquid
Heat Energy
Exothermic
 Definition:

Examples: __________________________________
Endothermic
 Definition:

Examples: ___________________________________
Phase Changes
Page 36 of 67
[Chapter Two: Matter & Energy ] [2011]
Phase Change Graph
Calculating the Energy absorbed or released (Joules) during a temperature
change.
Heat Energy = Cp x m x ΔT
Heat Energy (Joules)
Cp= Specific Heat (J/g°C)
m = mass (grams)
ΔT= change in temperature (°C)
Examples:
Examples:
Page 37 of 67
[Chapter Two: Matter & Energy ] [2011]
Calculating the Energy absorbed or released (Joules) during a phase change.
Energy of freezing/melting = Hf (Heat of Fusion for substance) X mass of substance
Energy of evaporation/condensation = Hv (Heat of Vaporization) X mass of substance
Heat of Fusion for water = 333 J/g
Heat of Vaporization for Water= 2260 J/g
Examples:
Examples:
Alternative Energy
Non-renewable Resources: ___________________________________________________
Examples: _____________________________________________________________
Renewable Resources: _______________________________________________________
Examples: _____________________________________________________________
Wind: _______________________________________________________________________
Uses: _________________________________________________________________
Pros:
Cons:
Geothermal: _________________________________________________________________
Uses: __________________________________________________________________
Pros:
Cons:
Solar: _______________________________________________________________________
Uses: __________________________________________________________________
Pros:
Cons:
Page 38 of 67
[Chapter Two: Matter & Energy ] [2011]
Fuel Cells: ___________________________________________________________________
Uses: __________________________________________________________________
Pros:
Cons:
Ethanol: _____________________________________________________________________
Uses: __________________________________________________________________
Pros:
Cons:
Nuclear Energy: _____________________________________________________________
Uses: __________________________________________________________________
Pros:
Cons:
Page 39 of 67
Chapter Three: Atoms and Moles 2011
Chapter Three
Atoms and Moles
INTRODUCTORY ACTIVITY: SEEING ATOMS
Purpose: To view elements on both a macroscopic and sub-microscopic level and to note similarities and
differences.
Instructions:
Go to your assigned lab station. At your station you will find selected samples of elements as well as pictures
showing atomic models of those elements. Observe both the sample elements and their models and use
your observations to answer the questions below.
The word macroscopic refers to things we are able to see with the naked eye. When you look at the
elements in their containers, you are viewing them macroscopically. Please record observations for each
of the four elements at your station.
Element
Macroscopic
Microscopic
Magnesium
Hydrogen
Carbon
Sulfur
1. Do you note any similarities between the samples of elements? What distinct differences do you
observe?
2. Now observe the pictures representing models of these elements. How are the models similar to
each other? How are they different?
[Chapter Three: Atoms and Moles ] [2011]
BLACK BOX ACTIVITY
You will be performing a “black box” activity. In this activity, you will attempt to determine the interior
structure for a “box” without actually looking inside. In other words, you will be using indirect observation(s)
to make a hypothesis.
Procedure:
1. Obtain a disc from your teacher. Note the number of the disc and record it in the table below.
2. Do not open the disc! You must make all of your observations without looking inside this disc. Make
observations regarding this disc. Record these observations in the table below.
3. Based on your observations, draw a picture of what you think the inside of the disc looks like.
4. Repeat steps 1-3 until you have “tested” at least four discs.
Hypothesis
Disc Number
Observations ( in words)
(picture of what you think it looks
like)
Page 41 of 67
[Chapter Three: Atoms and the Mole ] [2011]
RUTHERFORD-MARBLE ACTIVITY
(Measuring what you cannot see)
Directions: The keys to science are observation and measurement, which are often used together in
experiments. Experiments are carried out to test hypotheses that attempt to explain the world around us.
Also, experiments can lead to new hypotheses.
Chemistry experiments are often designed to gather information about what cannot be observed directly.
The purpose of this activity is to demonstrate how an experiment can provide information about something
that cannot be seen.
Objectives:
 Record data for repeated trials of an experiment
 DETERMINE THE UNKNOWN SIZE OF AN OBJECT WITHOUT DIRECT MEASUREMENT
 Compare the difference between the calculated and the actual sizes
Equipment:
7 marbles
meter stick
masking tape
Procedures:
1. Make a masking tape line, 60 cm long, on the floor. Use the meter stick to mark the tape at 5, 15, 25,
35, 45, and 55 cm. Also, use a small piece of tape to mark a spot about 1 meter away from the
center of the tape.
2. Place marbles along the long piece of tape, one at each of the marked spots. Place one marble on
the small piece of tape.
3. With eyes closed, one team member will roll the single marble toward the line of marbles. The
second team member will note whether this marble hits any of the other marbles or misses. The hit
or miss will be tallied on the data sheet. (If the marble misses the entire line of tape completely, the
trial is disregarded and another tallied.)
4. All of the marbles are returned to their original positions and the process is repeated for a total of 75
tallied trials.
5. After 75 trials, the team members trade places and 75more trials are completed .
DATA
Trial
Team member 1
Team member 2
Hits
Misses
Page 42 of 67
[Chapter Three: Atoms and the Mole ] [2011]
SECTION TWO: STRUCTURE OF ATOMS – ATOMIC SYMBOLS (pg. 84-89)
Will be different
for two
isotopes.
Will be the
same for two
isotopes
Relates to
neutrons only.
Relates to
protons only.
The higher of
the two
numbers
Related to the
number of
electrons
Related to the
number of
neutrons
Related to the
number of
protons
Map: Complete the following semantic map based on your reading:
Atomic
Number
Mass
Number
PRACTICE PROBLEMS:WRITING NUCLIDE SYMBOLS
Write nuclide symbols for each of the following isotopes:
Element
# of protons
# of neutrons
Oxygen
8
8
Oxygen
8
9
Oxygen
8
7
Hydrogen
1
0
Hydrogen
1
2
Nickel
28
27
Silver
47
50
Barium
56
60
Lithium
3
4
Beryllium
4
4
Nuclide symbol
Page 43 of 67
[Chapter Three: Atoms and the Mole ] [2011]
MORE NUCLIDE SYMBOLS/ SUBATOMIC PARTICLES
1. For each of the following, indicate the number of protons, neutrons, and electrons:
Nuclide Symbol Protons Neutrons Electrons
40
18
Ar
197
Au
79
69
30
Zn
2
25
12 Mg
3
15
7
N
Ba-138
Hg-201
Cs1- - 134
Page 44 of 67
[Chapter Three: Atoms and the Mole ] [2011]
NUCLIDE SYMBOLS CONTINUED
Complete the following table:
Element
Mass Number
Atomic Number
48
22
Boron
Number of
Protons
5
Phosphorus
31
Xenon
130
Number of
Neutrons
Number of
Electrons
(when neutral)
6
54
Helium
2
2
14
Radium
13
138
92
Silicon
28
146
14
1
Magnesium
25
0
12
30
Lanthanum
139
29
57
Complete the following table:
Element
Mass
Number
Ca
40
Ta
179
Mn2+
56
Si
Ce1+
142
Sb?
123
Cu
64
Atomic
Number
Number of
Protons
Number of
Neutrons
Number of
Electrons
15
14
Charge
57
-3
Page 45 of 67
[Chapter Three: Atoms and the Mole ] [2011]
Nuclide Symbol Summary Questions:
1. If you are given the symbol for an element and its mass number, what information
can you determine? (Assume the element is neutral.)
2. How are the atomic number and the mass number the same? How are they
different?
3. How are K-40 and K-41 the same? How are they different? What do we call two
elements like these?
COMPLETE YOUR MAP HERE:
..
Word List:
Atoms consist of:
Which make up the:
Which make up the:
Volume
Mass
Nucleus
Electrons
Energy levels
Ground state
Excited state
Move closer
Move farther away
Give off light/energy
Absorb light/energy
And are found in:
Electrons can be found:
from which they can
Page 46 of 67
[Chapter Three: Atoms and the Mole ] [2011]
Sublevel
# of orbitals
# of electrons
possible in each
orbital
(symbol)
s
p
d
f
Energy Level
(n)
1
2
3
4
5
6
7
# of
Type of
Sublevel Sublevels
s
s, p
# of electrons
possible in each
orbital
2
2,6
Total # of electrons
possible in this
sublevel
Total # of
electrons possible
in this main level
2
8
The answer comes to us in the form of a rule called the aufbau principle which states that electrons will
always fill up an atom from lowest energy to highest energy. In other words, the first electrons will go in the
lowest sublevel of the lowest energy level and work their way out from their.
A quick reminder, in terms of energy:
For energy levels 1<2<3<4 etc.
For sublevels s<p<d<f
Page 47 of 67
[Chapter Three: Atoms and the Mole ] [2011]
Complete electron configurations for the following atoms:
oxygen
boron
helium
sulfur
magnesium
Using the arrow diagram, complete electron configurations for the following elements:
1. titanium
2. calcium
3. germanium
4. strontium
5. hafnium
6. copper
7. aluminum
8. lead
9. selenium
10. iron
Page 48 of 67
[Chapter Three: Atoms and the Mole ] [2011]
1s
2s
2p
3s
3p
3d
4s
4p
4d
4f
5s
5p
5d
5f
6s
6p
6d
6f
7s
7p
7d
7f
Page 49 of 67
[Chapter Three: Atoms and the Mole ] [2011]
GREEK MODEL
Time Period: ___________
_________________________ was a Greek Philosopher who first proposed the idea of a small
solid and indestructible particle that composes all matter.
Democritus called this building block the ___________________.
Democritus
Aristotle
vs.
MORE RECENT ATOMIC THEORIES
The first individual to propose an atomic theory based on experimentation was
_______________________________.
Time period: ____________________
1. All matter consists of tiny particles. Dalton, like the Greeks, called these particles
_____________________.
2. Atoms of one element can neither be ________________ nor changed into atoms of any
other element.
3. All atoms of the same element are ____________________ in mass, size, and other
properties.
4. Atoms of one element _____________________________________________
from atoms of other elements.
5. In ____________________, atoms of different elements combine in simple, whole
number ratios.
Page 50 of 67
[Chapter Three: Atoms and the Mole ] [2011]
Thomson’s Model of the Atom
Time Period:____________________

Description of Experiment

What happened?

Conclusions from the Cathode Ray Tube (CRT)

Analogy for Model: ___________________________
From his experimentation Thomson
developed a model of the atom
called the plum-pudding model:
Ernest Rutherford
Draw his model
here
Time Period: _________________________
Description of the experiment (GOLD FOIL EXPERIMENT)
Page 51 of 67
[Chapter Three: Atoms and the Mole ] [2011]
What happened?
1.
2.
3.
Conclusions
1.
2.
3.
Analogy for Model: __________________________
Draw a picture of Rutherford’s Model of the Atom
James Chadwick
Time Period: _____________________
Discovered the neutral subatomic particle in the nucleus: ________________________
Bohr’s Model of the Atom
Time Period: ________________________
Description of Model:
Draw a picture of the model
Analogy: _____________
Page 52 of 67
[Chapter Three: Atoms and the Mole ] [2011]
Quantum Mechanical Model (Electron Cloud Model)
Description of Model
Time Period: _____________________
Draw a picture of the model:
Analogy: __________________
Structure of the Atom
Proton
Location:
Charge:
Neutron
Location:
Charge:
Electron
Location:
Charge:
Nuclide Symbols
Three Parts:
40
19
𝐾
Isotopes
Definition: _______________________________________________________
Ions
Definition: _______________________________________________________
Examples:
Page 53 of 67
[Chapter Three: Atoms and the Mole ] [2011]
Properties of Electrons
Most of the atom’s mass is made up of ____________________________.
Most of the atom’s volume is made up of _________________________.
Electrons are found in the __________________________.
The ______________________ the electron is to the nucleus the _________________ energy.
Ground State: _________________________________
Excited State:
Electron Configuration: locating & mapping electrons
Energy Levels (n)
- Farther away from the nucleus the ____________________
Sublevels/ orbitals
- n =1 , # sublevels are _________
- n=2, # sublevels are __________
- n=3 , # sublevels are __________
- n=4, # sublevels are ___________
- n= 5, # sublevels are __________
- n=6, # sublevels are __________
- n= 7, # sublevels are __________
An orbital is __________________________________________
The MAXIMUM number of electrons in ANY orbital is _____________
Page 54 of 67
[Chapter Three: Atoms and the Mole ] [2011]
4 different shapes of Orbitals
Orbital
Energy
Level
# of orbitals
# of sublevels
Type of
sublevel
Maximum # of
TOTAL # of
electrons in EACH electrons in orbital
orbital
# of electrons Total # of
in EACH orbital electrons in
EACH LEVEL
1
2
3
4
5
6
7
Page 55 of 67
[Chapter Four: The Periodic Table ] [2011]
Chapter Four
The Periodic Table of Elements
FOCUS ACTIVITY
To get a grasp of what this chapter will be about, you must understand the meaning of the word “periodic”. If
something is said to be periodic, it means that it has a repeating pattern in its arrangement. Many of the
concepts we deal with on a daily basis can be termed periodic.
1. Consider a monthly calendar. How are the days of the month arranged? In other words, are they in
one long row?
2. If the days of the month are not in one long row, how do they decide when to form a new row?
3. This is an example of periodic. Can you think of something else we might classify as periodic?
Page 56 of 67
[Chapter Four: The Periodic Table ] [2011]
SIMULATION: ARRANGING A PERIODIC TABLE
Card Deck #
1. Obtain a deck of cards from your teacher. Write the number from the back of your cards in the space
provided above. Observe your cards. Notice the differences in each card. There are six characteristics that
can describe each card. For example: one characteristic is the number of holes each card has. List five
more characteristics you observe:
2. Arrange your cards in table fashion with rows and columns. Your arrangement must make sense.
This means that there must be order and reason. You will know you have an appropriate table when
there is a pattern for all six characteristics. NOTE: There is one piece missing from your deck of
cards. If constructed properly, your table will have a spot for that missing piece to fit in.
3. Study the pattern for each characteristic and describe in the table below.
Property
Observed Pattern
4. Using the patterns you have identified, determine the six properties of your missing piece. Write these
properties in the table below. When you feel that you have properly identified your missing piece, ask the
teacher to check your answer.
Property
Description for missing piece
Page 57 of 67
[Chapter Four: The Periodic Table ] [2011]
SECTION 1: HOW ARE ELEMENTS ORGANIZED? (pg. 116-118)
Contribution
Developed by:
The Periodic Table
Developed by:
Contribution
Contribution
Page 58 of 67
[Chapter Four: The Periodic Table ] [2011]
SKILL REVIEW
Use your notes from Chapter One to find the necessary formula. Calculate the
density of the following:
1. Determine the density of a substance with a mass of 19 grams and a volume of 25 ml.
2. A lump of metal X weighing 15 grams is placed in a graduated cylinder containing 10 ml of water.
The water rises to the 20 ml mark. What is the volume of the lump of metal X.? Calculate the
density of metal X.
SKILL REVIEW
Complete electron configurations for the following elements; revisit your notes from CHAPTER THREE if
necessary.
1. sodium
2. copper
3. arsenic
VALENCE ELECTRONS
Identifying valence electrons:
Practice Questions: Go up to the electrons configurations you just completed and write the
electron configuration of just the valence electrons.
Sodium
Copper
Arsenic
Page 59 of 67
[Chapter Four: The Periodic Table ] [2011]
COMPLETING THE TABLE: UNDERSTANDING WHY THE ELEMENTS REPEAT






Your teacher will be assigning you elements that you are to do an electron configuration
for.
After determining the correct electron configuration, identify the valence electrons for each
element.
On the next page is a blank periodic table. Find the block(s) which represents each of
your assigned elements and write the valance electron configuration for each of the
element in the appropriate block.
Do the same on the board, writing your valence electron configurations in the appropriate
place.
As other people put their answers on the transparency, copy them onto your periodic table.
Then answer the following questions:
Questions:
1. Look at the columns on your periodic table. What trend do you notice in the valance
electrons?
2. What trend do you notice in the rows of your periodic table?
3. The element Francium is not shown on your partial periodic table. However you can
predict it’s valence electron configuration based on what you have observed. What do you
think the valance configuration for francium would be?
Page 60 of 67
[Chapter Four: The Periodic Table ]
2011]
Electron Configuration Pattern in the Periodic Table
Page 61 of 67
[Chapter Four: The Periodic Table ]
Periodical Table Trends:
Electron Configuration
Density
Valence Electrons
Family Names
s,p,d,f BLOCKS
2011]
Nonmetals
Metalloids
Metals
Page 62 of 67
[Chapter Four: The Periodic Table ]
SECTION 1: HOW ARE ELEMENTS ORGANIZED CONTINUED (pg. 119-122)
Complete a word sort using the following words:
Energy levels
Group
Horizontal
Period
Periodic law
Similar properties
Valence electrons
Vertical
SECTION 2: TOUR OF THE PERIODIC TABLE (pg. 124 to131)
After reading pages 124 to 128, complete the following table:
Name of Group
Location on the
periodic table
Main characteristics
Page 63 of 67
[Chapter Four: The Periodic Table ]
SECTION 2: METALS VS. NON METALS (PAGES 128 TO 131)
Compare the Characteristics of Metals and Nonmetals
METALS
NONMETALS
Page 64 of 67
[Chapter Four: The Periodic Table ]
Section: How are Elements Organized? (pgs116-122)
Answer the following questions in the space provided.
1. Why do Li, Na, K, Rb, Cs, and Fr all react with Cl in a 1:1 ratio forming substances
with similar properties?
2. Explain the method that John Newlands used to organize the elements.
3. What method did Dmitri Mendeleev use to arrange his periodic table?
4. Why did Mendeleev have gaps in his table? How did he use these gaps?
5. What was Henry Moseley’s contribution to the periodic table?
6. Why was Moseley able to resolve the discrepancies in Mendeleev’s table when
Mendeleev could not?
7. Explain the importance of valence electrons.
8. Why do elements with similar properties appear at regular intervals in the
periodic table?
9. How is the electron configuration similar for each element in a group?
10. How is the electron configuration similar for each element in a period?
Page 65 of 67
[Chapter Four: The Periodic Table ]
Section: Tour of the Periodic Table (pgs. 124-131)
Complete each statement below by choosing a term from the following list.
Terms may be used more than once.
main-group elements
alkaline earth metals
halogens metals
alkali metals
transition metals
hydrogen
noble gases
1.The ___________ have a single electron in the highest occupied energy level.
2. The _____________ are in the s- and p-blocks of the periodic table.
3. All the ________________ have two valence electrons and get to a stable electron configuration
by losing two electrons.
4. Unlike the main-group elements, each group of the ________________ does not have the identical
outer electron configuration.
5. The _________________, the most reactive group of non-metals, achieve stable electron
configurations by gaining one electron.
6. The ________________ have a full set of electrons in their outermost energy level.
7. The __________________are very stable and have low reactivity.
8. The __________________ are highly reactive and readily form salts with metals.
9. In general, the ____________________are metals that are less reactive than the alkali metals and
the alkaline earth metals.
10. The ____________________ are metals that lose one electron when they react with water to
form alkaline solutions.
11. Most elements are ___________________________.
12. With its one valence electron,___________________ reacts with many other elements.
Answer the following questions in the space provided.
13. Which groups compose the main-group elements?
14. Why are the main-group elements called the representative elements?
15. Why are Group 2 elements less reactive than Group 1 elements?
16. Using electron configurations, explain why the halogens readily react with the
alkali metals to form salts.
Page 66 of 67
[Chapter Four: The Periodic Table ]
PA..
Section: Where Did the Elements Come From? (pgs. 142-147)
Complete each statement below by writing the correct terms or terms.
1. Most of the atoms in living things are of just six elements ______________, ______________,
_______________, _________________, ________________, and ________________.
2. Immediately after the big bang, temperatures were extremely high and only
_________________________could exist.
3. As the universe began to cool, energy was converted to____________________, in the form of
__________________ ,________________ and _________________ .
4. As the universe continued to cool, these particles joined and formed the first
two elements, __________________ and _________________________ .
5. The temperatures in stars get high enough to fuse _______________ nuclei with one another,
forming elements of still higher atomic numbers.
6. Massive atoms such as iron and nickel form by repeated __________________.
7. When a massive star has converted almost all its core hydrogen and helium
into heavier elements, it collapses and blows apart in an explosion called a _______________
forming elements heavier than iron.
8. The nuclear reaction that changes one nucleus into another by radioactive
disintegration or by bombardment with other particles is called___________________.
9. Elements that chemists have created are called ____________________ elements.
10. The special equipment that scientists use to create elements are called_______________.
Page 67 of 67