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Grade 11
Unit 6
SCIENCE 1106
CHEMICAL REACTIONS, RATES
AND EQUILIBRIUM
CONTENTS
I. CHEMICAL REACTIONS . . . . . . . . . . . . . . . .
2
DETECTION . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
2
ENERGY REPRESENTATION . . . . . . . . . . . . . . . . . . . . . . .
9
II. REACTION RATES . . . . . . . . . . . . . . . . . . . . .
27
RATE VARIABLES . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
27
REACTION DIAGRAMS . . . . . . . . . . . . . . . . . . . . . . . . . . .
33
III. REACTION EQUILIBRIUMS . . . . . . . . . . . . .
44
EQUILIBRIUM MATHEMATICS . . . . . . . . . . . . . . . . . . . . . .
44
EQUILIBRIUM VARIABLES . . . . . . . . . . . . . . . . . . . . . . . .
55
Author:
Harold Wengert, Ed.D.
Editor:
Alan Christopherson, M.S.
Illustrators:
Alpha Omega Graphics
804 N. 2nd Ave. E., Rock Rapids, IA 51246-1759
© MM by Alpha Omega Publications, Inc. All rights reserved.
LIFEPAC is a registered trademark of Alpha Omega Publications, Inc.
All trademarks and/or service marks referenced in this material are the property of their respective owners. Alpha Omega Publications, Inc.
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SCIENCE 1106:
CHEMICAL REACTIONS, RATES
AND EQUILIBRIUMS
Have you ever wondered why a fire is hot, or
why the food you eat produces energy? In both
examples, a reorganization of molecules occurs. In
the fire, fuel reacts with oxygen to produce heat
plus different chemicals. In the case of food eaten,
digestion breaks up the molecules so that they can
react with oxygen to produce heat (energy) and
waste products. The reorganization of molecules
involves energy and energy changes.
You will explore how molecular reorganization
produces energy, how and why reactions occur, and
what controls the rate of a reaction. You may need
to review previous LlFEPACs. Also, be sure to have
the Periodic Table from SCIENCE 1101 handy.
OBJECTIVES
Read these objectives. The objectives tell you what you will be able to do when you have
successfully completed this LIFEPAC®.
When you have finished this LIFEPAC, you should be able to:
1. Describe how a reaction can be detected.
2. Interpret energy (enthalpy) diagrams.
3. Explain what factors affect the rates of reactions.
4. Explain Le Chatelier’s Principle.
5. Describe an equilibrium condition in a reaction.
6. Apply the Law of Chemical Equilibrium.
7. Explain how condition variables affect an equilibrium reaction.
Survey the LIFEPAC. Ask yourself some questions about this study. Write your questions here.
1
I. CHEMICAL REACTIONS
With the models of atoms, molecules, and ions
developed in previous LlFEPACs, we are now
ready to investigate in depth the systems of
chemical reactions. What is a chemical reaction?
From our investigations so far, we may generalize
that in a chemical reaction the atoms are
rearranged to form new substances with new
formulas. In other words, the products (substances
produced) are different in molecular makeup from
the reactants (substances started with).
SECTION OBJECTIVES
Review these objectives. When you have completed this section, you should be able to:
1. Describe how a reaction can be detected.
2. Interpret energy (enthalpy) diagrams.
VOCABULARY
Study these words to enhance your learning success in this section.
endothermic
exothermic
enthalpy
heat of reaction
precipitate
Note: All vocabulary words in this LIFEPAC appear in boldface print the first time they are used. If you are unsure
of the meaning when you are reading, study the definitions given.
DETECTION
We see many examples of chemical reactions
around us daily. The changing colors of the leaves,
the growing of new body cells, the tarnishing of
silverware, the fading of the color in paints, and the
burning of gasoline in a car are all examples of
chemical reactions. Some reactions are easy to
detect and others go on too slowly to be noticed.
How can chemical reactions be detected?
Do this experiment.
Solid formation. You studied in Science
LIFEPAC 1102 about distinguishing between
chemical, physical, and phase changes. Any
reaction that caused the starting materials to
completely lose their properties to form new and
distinct substances was a chemical change. The
next set of activities will help you to discover one
way a chemical reaction can be detected.
WEAR GOGGLES
These supplies are needed:
0.01 M sodium chloride, NaCl solution, table salt – 0.58 g/L of solution
0.01 M potassium chromate, K2CrO4 solution, dilute solution – 1.94 g/L of solution;
solid can be purchased at a chemical supply store
0.01 M silver nitrate, AgNO3 solution, about 1.7 g/L of solution or a diluted solution;
solid or solution can be purchased at a chemical supply store
several small test tubes
several medicine (eye) droppers or pipets
2
Follow these directions and answer the questions. Put a check in the box when each
step is completed.
❏ 1. Place ten drops of 0.01 M NaCl in a test tube and ten drops of 0.01 M potassium chromate,
K2CrO4, in another test tube.
1.1
What do the three solutions look like?
❏ 2. Put ten drops of silver nitrate, AgNO3, in each of the other two solutions.
1.2
What happened in Step 2?
1.3
Why do you believe a reaction did or did not occur?
❏ 3. In a clean test tube put ten drops each of K2CrO4 and AgNO3.
❏ 4. Slowly, a drop at a time, add NaCl until a significant change occurs.
1.4
What happened in Step 4?
1.5
What is your hypothesis for the observations and changes?
Color change. Another way to detect a chemical change is by comparing colors. When the leaves turn
colors in the fall, a chemical change has occurred. The next experiment will show how color is used in
reaction detection.
Do this experiment.
WEAR GOGGLES
These supplies are needed:
0.01 M acidified iron (II) sulfate, FeSO4 - 1.52 g/liter of solution and 1 ml
concentrated HCl; solid FeSO4 can be purchased at a chemical supply store
0.01 M potassium permanganate, KMnO4 - 1.58 g/liter of solution; solid KMnO4
can be purchased at a chemical supply store
0.01 M NaCl sodium chloride solution, table salt – 0.58 g/L of solution
0.01 M ammonium nitrate, NH4NO3 - 0.80 g/liter of solution; solid ammonium
nitrate can be purchased at a chemical supply store
several test tubes or glass baby-food jars
several medicine (eye) droppers or pipets
3
Follow these directions and answer the questions. Put a check in the box when each
step is completed.
❏ 1. Place about 3 ml of 0.01 M acidified iron (II) sulfate, FeSO4, in a test tube.
❏ 2. Add 6 drops of 0.01 M potassium permanganate, KMnO4, 1 drop at a time, shaking the test
tube after each additional drop.
1.6
What do the reactants look like?
1.7
What happens when the reactants are combined?
1.8
Why do you think a reaction has, or has not, occurred?
❏ 3. Add 10 drops of 0.01 M sodium chloride, NaCl, to 10 drops of 0.01 M ammonium nitrate,
NH4NO3.
1.9
What do the reactants look like?
1.10
What happens when the reactants are combined?
1.11
Why do you think a reaction has, or has not, occurred?
Gas formation. We have now seen that
chemical reactions can be detected by the change
in color and by the formation of a precipitate. A
third way is found to occur very frequently in
nature. The production of methane gas, CH4, when
living material decays in the absence of oxygen, is
a very important example of gas formation in a
chemical reaction. This example is the source of
natural gas, the gas used in kitchen stoves, home
furnaces, and factories. When the living materials
were covered during the time of Deluge (Genesis,
Chapters 6-8), the decaying organic materials
produced the methane gas. This gas is found both
in oil wells and in coal mines.
Our study will investigate this third way to
detect chemical reactions.
4
Do this investigation.
WEAR GOGGLES
These supplies are needed:
concentrated HCl (hydrochloric acid) - purchased from a building contractor, pool
supply, concrete mason, or chemical company; common name is muriatic acid.
glacial acetic acid - purchased from a chemical supply store
6 M HCl - made by putting 50 ml of concentrated HCl into H2O to make 100 ml
1 M HCl - made by putting 8.3 ml HCl into H2O to make 100 ml
0.1 M HCl - made by putting 10 ml of the 1 M HCl in H2O to make 100 ml
6 M acetic acid - made by putting 34.8 ml glacial acetic in H2O to make 100 ml
1 M acetic acid - made by putting 5.8 ml glacial acetic in H2O to make 100 ml
5 small chips of blackboard chalk (white is best) or calcium carbonate chips
5 small test tubes (purchased from a hobby shop or a chemical supply company)
Follow these directions and answer the questions. Put a check in the box when each
step is completed.
❏ 1. Set up the following experiment.
a.
3 ml
6M
HCl
b.
3 ml
6M
CH3COOH
c.
3 ml
1M
HCl
d.
3 ml
0.1 M
HCl
Add
1.12
What do the reactants look like?
1.13
What did each reaction look like?
a.
b.
c.
5
e.
3 ml
1M
CH3COOH
small chalk chip,
CaCO3(s) in each
test tube
d.
e.
1.14
Why do you think a reaction has, or has not, occurred?
1.15
We have seen three types of reactions so far in this LIFEPAC. In summary, how could you
detect if a reaction has occurred?
a.
b.
c.
Adult check ___________________
Initial
Date
Temperature change. The fourth class of
reactions is much more difficult to detect than the
previous three. This fourth class is really the basis
of the other three. Every chemical reaction involves
energy in some form, because all chemical
reactions involve either the breaking of old bonds
or the forming of new bonds or both. We know that
every chemical bond contains energy or it would
not exist. This result means that every chemical
Do this investigation.
reaction involves an energy change because the
bonds are changed. Energy tied up in the bonds of
atoms, molecules, ions, and the rest of the chemical
system is called enthalpy. The absolute amount of
this energy cannot be measured, but the change in
enthalpy from the reactants to products can be
measured. If the energy is lost (given off), the
reaction is exothermic; if energy is gained, the
reaction is endothermic.
WEAR GOGGLES
These supplies are needed:
solid sodium hydroxide, NaOH – lye, can be purchased in a grocery store, a hardware
store, a soap making supply store, or a chemical supply store
solid ammonium nitrate, NH4NO3 – can be purchased from a fertilizer store or a
chemical supply store
concentrated hydrochloric acid, HCl – can be purchased in a pool supply store or a
chemical supply store
phenolphthalein solution (or other indicator) – can be purchased from a chemical
supply store
thermometer to fit test tubes
water
forceps (tweezers)
test tubes with stoppers
6
Follow these directions and answer the questions. Put a check in the box when each
step is completed.
❏ 1. CAUTION: CONCENTRATED HCl IS VERY DANGEROUS TO SKIN, EYES, AND
CLOTHES. IF SPILLED, WIPE UP IMMEDIATELY WITH WET TOWELS OR SPONGE.
IF SPILLED ON SKIN OR CLOTHES, WASH FREELY WITH WATER AND INFORM
YOUR TEACHER OF THE ACCIDENT. Your teacher has one place in the room where this
acid is kept and used. Ask for its location and safety precautions. To 2 ml of tap water in
a test tube, add 10 drops, a drop at a time, of 12 M hydrochloric acid, HCl. Use a
thermometer to check for any temperature change. Save the diluted acid for future use in
Step 5.
1.16
What do the reactants look like?
1.17
What happened when the HCl was added to the H2O?
1.18
Why do you think a reaction has, or has not, taken place?
1.19
Whenever the temperature changes as the result of chemicals being added together, the
enthalpy of the products will differ from the enthalpy of the reactants.
a. Has a temperature change taken place?
b. Has an energy change taken place?
c. In order to get the solution back to the original temperature for an energy comparison with
the reactants, will you need to heat or cool the test tube?
d. Did this reaction release energy or take on added energy?
e. Therefore, is the enthalpy of the products greater or less than the enthalpy of the reactants?
H3O+ + Cl- + heat
Symbol Model:
HCl + H2O
Energy Model:
Reactants’ Enthalpy (E1)
Products’ Enthalpy (E2) + heat
❏ 2. To 3 ml of tap water in a test tube, add 2 small pellets of solid sodium hydroxide, NaOH(S).
DO NOT HANDLE SODIUM HYDROXIDE PELLETS WITH YOUR FINGERS — USE
FORCEPS. Place a stopper in the test tube and shake it gently. Use a thermometer to
check for a temperature change. Save this solution for Step 3.
1.20
What do the reactants look like?
1.21
What happened when NaOH was added to H2O?
1.22
Why do you think a reaction did, or did not, occur?
7
1.23
Analyze the changes in the NaOH - H2O system.
a. Has a temperature change taken place?
b. Has an energy change taken place?
c. To get the solution back to its original temperature, must you add or take away energy?
d. How does this compare to the previous HCl reaction?
e. Do the products or the reactants have more enthalpy?
f. Is this reaction exothermic?
1.24
Symbol Model:
NaOH(s)
Energy Model:
E1
Na+(aq) + OH-(aq)
(Complete this model.)
Note: Read all the directions before doing Steps 3 through 7.
❏ 3. Place 1 ml of the NaOH solution made in Step 2 in a test tube.
❏ 4. Add a few drops of phenolphthalein solution (an indicator solution).
❏ 5. Add, dropwise, the HCl solution made in Step 1 until the color changes.
❏ 6. Mix thoroughly between additional drops.
❏ 7. Use a thermometer to check for a temperature change.
1.25
What do the reactants look like?
1.26
What happens when the reactants are placed together?
1.27
Why do you believe a reaction has, or has not, occurred?
1.28
Describe the energy changes.
a. Has an energy change taken place?
b. How can you tell?
1.29
Describe the enthalpy change.
a. Is the NaOH - HCl reaction exothermic?
b. Do the products or reactants have more enthalpy?
8
1.30
Symbol Model:
Na+(aq) + OH-(aq) + H+(aq) + Cl-(aq)
Na+(aq) + Cl-(aq) + H2O
Energy Model:
(Complete this model.)
❏ 8. Place solid ammonium nitrate, NH4NO3, to a depth of about 1 cm in a test tube.
❏ 9. Add about 1 ml of water that is at room temperature.
1.31
Analyze the NH4NO3 - H2O system.
a. What happens to the temperature?
b. To compare the enthalpy of the products with the enthalpy of the reactants, both must be
at the same temperature. Would you add or take away energy to get the products at the
same temperature as the reactants?
c. Do the products or reactants have more enthalpy?
d. Is this reaction the same as the HCl and NaOH systems?
1.32
Describe the enthalpy of the NH4NO3 system.
a. Was the enthalpy of the products of exothermic reactions increased (+) or decreased (-)?
b. Was the enthalpy of the products of the dissolving NH4NO3 plus (+) or minus (-)?
c. What is the reason for your prediction in (b)?
1.33
Symbol Model:
NH4NO3
NH4+(aq) + NO3-(aq)
Energy Model:
(Make a complete energy model.)
Adult check ___________________
Initial
Date
ENERGY REPRESENTATION
Scientists are able to display conveniently
scientific data in graphic form. This display can be
by models, such as the energy and symbol models
we used in the previous section. Other ways are by
equations and diagrams. You will learn to
represent reactions by equations and diagrams,
and you will learn how to figure the enthalpy of
simple chemical reactions. These activities will
help you to understand how the breaking and
forming of bonds in chemical reactions produce or
absorb energy and how common substances make
good fuel for heating foods and houses.
Energy diagrams. Chemical reactions are
studied by the energy contained in the products in
comparison to the reactants. To study the absolute
amount of energy in the system is difficult because
not all of the energy can be measured. Therefore,
only the ∆H (change in energy) is of interest because
this change in energy is the only amount that can be
measured directly. What is really being described is
the difference between bond-forming and bondbreaking energies. This difference can be
summarized:
9
Exothermic Reaction
E
products
E
products
<E
(∆H is negative because the reaction produced energy which is lost to the
environment)
reactants
– E reactants = -∆H
(reactants
products + energy)
Endothermic Reaction
E
products
E
products
>E
(∆H is positive because energy must be added to the reactants to make the
products)
reactants
– E reactants = +∆H
(reactants + energy
products)
Do these activities.
1.34
List the four ways that a chemical reaction can be detected: (Hint: These were covered in
Section I.)
a.
b.
c.
d.
1.35
Consider the following reactions.
(Note: small subscript letters of (s), (l), and (g) indicate the states of matter of the
components. The subscript (aq) means the substance is in a water solution.)
Mg(s) + 1/2 O2(g)
MgO(s) + 146 kcal/mole
H2(g) + 1/2 O2(g)
∆H = -57.82 kcal/mole
H2O(g)
What type of reaction is represented by the previous two examples?
1.36
Consider the following reactions.
/2 H2(g) + 1/2 I2(g)
1
∆H = + 6.2 kcal/mole
HI(g)
21.0 kcal/mole + C(s) + 2S(s)
CS2(l)
What type of reaction is the previous two examples?
1.37
Consider the following reactions. Which reactions are endothermic?
a. H2(g) + 1/2 O2(g)
b. 1/2 N2(g) + 1/2 O2(g)
e. NH3(g)
∆H = +21.6 kcal/mole
NO(g)
c. 1/2 N2(g) + O2(g) + 8.1 kcal
d. 1/2 N2(g) + 3/2 H2(g)
∆H = -57.83 kcal/mole
H2O(g)
NO2(g)
NH3(g) + 11.0 kcal/mole
/2 N2(g) + 3/2 H2(g)
1
10
∆H = +11.0 kcal/mole
1.38
Study the following two reactions.
(A)
C
(graphite)
+ O2(g)
CO2(g) + 94.05 kcal
(B)
C
(diamond)
+ O2(g)
CO2(g) + 94.50 kcal
a. Do the substances in the box for reaction A and for B have the same amount of enthalpy
at the same temperature?
b. If your answer to a. is “yes” where does the difference in the ∆H values for the two
reactions come from?
Adult check ___________________
Initial
Date
Review Activities 1.16 to 1.20. The reaction between HCl and H2O can be represented in
the following ways:
H3O+ + Cl-
Symbol Model:
HCl + H2O
Energy Model:
Reactants’ Enthalpy (E1)
+ heat
Products’ Enthalpy (E2) + heat
Energy Stored
in Products
Energy Stored
in Reactants
+
Excess
Energy
Lost
Enthalpy or Bonding Energy
Increase
Reactants and products are compared at the same temperature.
0
Energy (Enthalpy) Diagram
Energy in
Reactants
HCl + H2O
-∆H = Heat is given off; this is an
exothermic reaction
-∆H
(products contain less enthalpy
than reactants)
Reaction Coordinate
The energy produced, ∆H, is equal to
-∆H
= Enthalpy of products - Enthalpy of reactants
= E2 - E1
11
Do the following activities.
Review Activities 1.20 through 1.24. Use that data and analysis for this activity. Using the
enthalpy diagram for HCl - H2O as a model, complete and label each part of the following
diagram for the NaOH - H2O reaction.
Enthalpy or Bonding Energy
Increase
1.39
Energy in Reactants
NaOH(s)
0
Reaction Coordinate
Review Activities 1.25 through 1.30. Complete and label the following enthalpy diagram for
the HCl - NaOH system.
Enthalpy or Energy of Bonding
1.40
0
Reaction Coordinate
12
1.41
Review Activities 1.31 through 1.33. Draw an enthalpy diagram for the dissolving of NH4NO3.
Label all of the parts for identification.
Adult check ___________________
Initial
Date
Heat of reaction. When some chemical changes
occur, the only easily observed indication of a
reaction is the energy absorbed or released. Actually
all chemical reactions involve a release or absorption
of energy. This energy is called the heat of reaction.
In this section of this LIFEPAC you will:
Why? Do we have an energy crisis? Why or why
not? The following discussion will help us to deal
with such questions as these.
As we begin to think about energy changes, let
us recall one of the regularities of nature with
which we are already familiar: The Law of
Conservation of Energy. This law reminds us that
as we observe chemical changes, energy may
change from one form to another but the amount of
energy in an isolated system does not change.
Energy is neither produced nor destroyed during a
chemical change. No chemical change creates
energy, only releases energy.
We already know that chemical changes must
involve the breaking and making of chemical
bonds. We also know that chemical bonds involve
energy. Perhaps the energy released or absorbed by
substances during a chemical change may be
explained by returning to a consideration of
chemical bonds.
a. Analyze the fact that all chemical changes
involve energy and see if this makes sense and
b. Try to discover a way to predict the energy
change in a reaction.
Have you ever thought about these forms of
energy: about calories or food, about having enough
energy for the game tonight or making weight for
wrestling? Is this energy the same as that used by
a plant, a 747, a model plane, or a truck? Why?
How can stuff—coal, gasoline, butane—be changed
into motion? From where does the “motion-energy”
come? Can you use water, sand, or glass as a fuel?
13
Do the following activities.
1.42
When bonds are formed, energy is lost by the atoms that are bonding. When two hydrogen
atoms bond to form H2, 104 kcal of energy are released for every mole of molecules formed
from the hydrogen atoms in the monatomic state.
+
H
Energy
H
Bond
Energy
H:H
When two chlorine atoms bond to form Cl2, 81 kcal are released for each mole of molecules
formed.
+
Cl
Energy
Cl
Bond
Energy
Cl
Cl
a. When bonds are formed, is energy gained or lost?
b. The energy released by the atoms when they bond to form a molecule is called the
c. Bond energy is expressed in kilocalories per mole of bonds formed. The bond energy for H2
is kcal per mole H2.
d. The strength of a single hydrogen bond holding two atoms together in one molecule would
be:
=
kcal/bond
6 x 1023 bonds/mole H2
Bond energies have been determined for many different bonds. In some cases they are only
approximations, but they are still useful for making some predictions. The following typical values will be
helpful to us as we proceed through this LIFEPAC.
14
Bond
Bond Energy (kcal/mole)
C-H
O - O (in O2)
C = O (in CO2)
H-O
H-H
C-C
C-O
Cl - Cl
99
119
192 (this value is for one double bond)
111
104
83
84
81
An example of a bond energy analysis for a reaction follows. The reaction 2H2 + O2
considered.
1
H-H
Bonds
Breaking
1 H-H
2 H-H
3 O-O
Total
2
+
H-H
Energy (kcal)
+104
+104
+119
+327
3
+
4
5
O-O
H-O-H
Bonds
Forming
Energy (kcal)
4
5
6
7
H-O
O-H
H-O
O-H
Total
-111
-111
-111
-111
-444
Net Energy = Breaking + Forming
= 327 kcal + (-444 kcal)
∆H = -117 kcal
= -117 kcal
6
+
7
H-O-H
+119
-111
+104
Enthalpy
Bond #
2H2O will be
-111
+104
-111
Net
Energy
(kcal)
-111
This analysis shows that when H2 and O2 are
combined to form H2O, energy is released. This
reaction is exothermic.
Energy Chart
Do these activities.
1.43
Try to predict the energy change involved in a familiar reaction, the burning of methane
(CH4) in a stove or furnace. Complete this equation for the combustion of methane:
a.
CH4 +
O2
CO2 +
H2O(l)
b. Now use the model shown above and prepare the structures for each of the reactant
molecules present in your equation. In the following space draw all of the structures for
the reactant molecules. Use dashes to represent each bond.
Adult check ___________________
Initial
Date
15
1.44
Now break down the reactant molecules bond by bond. Follow this procedure:
a. Imagine that every molecule you handle represents a mole of that kind of molecule.
b. Use the energy chart to make a graph to summarize what you are doing.
c. Record in the following table the energy needed to break each bond. Since we have to put
energy into the system to break bonds, the system gains energy. When we do this, we will
assign these bond energies a positive value.
List Each Bond Broken
Bond Energy (kcal/mole)
a.
b.
c.
d.
e.
f.
Total Energy Needed
g.
1.45
Now put the products together following the same procedure. Since energy is released when
bonds are formed, the system is releasing energy. We therefore assign these energies a
negative value.
Product Structures:
Bonds Formed
Bond Energy (kcal/mole)
a.
b.
c.
d.
e.
f.
g.
1.46
Total Energy Released
Net Energy Change for the entire reaction:
Total Energy
Absorbed
1.47
+
Total Energy
Released
=
In the following space draw a sketch of the graph you made on the energy chart. Label it
including the net energy change for the reaction.
16
1.48
The net energy change for a reaction is called the enthalpy change. Enthalpy is given the
symbol H. The enthalpy change is given the symbol ∆H.
a. The ∆H for the combustion of one mole of methane is
.
b. The negative sign on the enthalpy change for this reaction means the methane-oxygen
(lost, gained) energy.
system
c. This energy was
(absorbed from, released to) the environment.
Adult check ___________________
Initial
Date
The structures called reactants possess a certain
amount of potential energy (PE) because of the
arrangement of the atoms. This energy is called H1.
H1
E
N
T
H
A
L
P
Y
H2
the system to release heat. As long as the energy
remains stored, it is PE and not KE. A fuel does not
actually contain heat any more than a jack-in-thebox contains motion. Under the right conditions the
energy stored in a gallon of gasoline or in a jack-inthe-box can be released, and we can observe it as it
is on the move.
Imagine that you live in a suburb of Chicago. If
you stand on the school roof and contemplate the
buildings in the Loop, you will be brought “down to
earth.” You may have felt “on top of the world” when
you were on the observation deck of the Willis Tower
CH4 + 2O2
HEAT
&
LIGHT
CO2 + 2H2O(l)
The structures called products have less
potential energy. They also have stronger bonds, on
the average, and the molecules are more stable
than the reactants. The enthalpy (or PE) of the
products is called H2.
The enthalpy change is ∆ H = H2 - H1.
Energy was stored in the reactant structures
(H1). The products have less energy stored in them
(H2). The energy difference is released as energy of
motion (kinetic): heat and light.
Every molecule can be thought of as a
storehouse of energy. This is easy to believe when
we think of the molecules in a match, a candle, or a
bit of food. When a match burns, a noticeable
amount of heat and light is given off. Since the
released energy has the form of heat (kinetic
energy), chemists often talk about the heat content
of a system. By this term they mean the potential of
FIGURE 1
17
(commonly referred to as Sears Tower) only to be
laughed at by your neighbor who has just scaled Mt.
Everest. Even if you are just standing waiting for a
bus, you can take comfort in knowing the people in
Death Valley (altitude -280 ft.) have to “look up to
you.”
Do you realize how Death Valley rates a
negative altitude? It was just arbitrary. Someone
decided to measure altitude from sea level and
Death Valley got the short end of the meter stick. If
you want to get “up in the world,” just convince
everyone to start measuring altitude from Death
Valley.
No one knows for sure how “high” one
substance is above the lowest possible potential
energy value. Arbitrarily and conveniently,
chemists have decided to use as a reference point
the condition of the free elements in their natural
state. All the elements as they exist in nature at
25° C are assigned an enthalpy of zero.
180
170
C(g)
160
150
140
130
120
110
100
90
80
70
E
N
T
H
A
L
P
Y
60
50
O(g)
H(g)
40
30
NO
20
10
0
-10
-20
NO2
N2(g), O2(g), C(s), and so forth
NH3(g)
CH4(g)
CO
-30
-40
-50
H2O(g)
-60
-70
H2O(l)
-80
-90
CO2
-100
-110
FIGURE 2
18
Do this activity.
1.49
Study the graph represented in Figure 2.
a. How could it be useful to you?
b. How does it help to make sense of the reaction we have looked at concerning CH4 in 1.43
through 1.48?
This order is not necessarily the one in which
things actually happen. The events may (and
probably do) happen in a completely different
sequence. (Perhaps some new bonds are made
before all of the old bonds are broken.)
To prove a point that will be helpful in the
future, look once again at the reaction in a natural
gas burner or stove. A particular reaction
mechanism was assumed, namely:
1. Break all the bonds in CH4 and O2, then
2. Make all the bonds in CO2 and H2O.
Do these activities.
1.50
Try to work through the CH4 - O2 reaction following a different sequence of bond
breaking/making. Be creative. Use an energy chart if you want to.
a. Make an enthalpy diagram to describe your sequence. When you are finished, your
diagram may be similar to the one at the right.
E
Energy Diagram (Chart)
b. Determine the ∆H for the reaction.
19
1.51
If possible find someone who has a diagram unlike your own. Explain your sequence to him.
Listen to his explanation of his diagram.
b. Your ∆H is
a. Do you understand the other person’s diagram?
c. The other person’s ∆H is
.
.
d. What generalization might this similarity suggest about the ∆H of a chemical reaction and
the sequence or path of the reaction?
1.52
Jim has to walk down five floors to get from his apartment to Steve’s apartment. Joe has to
walk down two floors to get to Steve’s apartment (unless he takes the elevator).
a. How far does Jim have to walk to get from his place to Joe’s?
.
b. On what floor does Steve live?
c. Did you need to know answer (b.) to answer the first question?
d. Draw a picture to help someone understand how to answer the first question.
1.53
One way to make 1 mole of CO2: (I) C
(graphite)
Another way: (II) CO(g) + 1/2O2
+ O2
CO2(g) + 94.05 kcal
CO2(g) + 67.6 kcal
How much energy would be absorbed or released if we make 1 mole of CO(g) starting with 1
mole C (graphite) and 1 mole O2?
Hints: 1. We could write: (III) C (graphite) + O2(g)
CO(g) + 1/2 O2
2. Maybe Jim, Joe, and Steve could help us.
20
1.54
Sue has to walk down six flights to visit Sarah. Sharon walks up five flights from her
apartment to visit Sarah. If Sue wants to go from her apartment to Sharon’s what will she
have to do? Draw a picture to help someone understand how to solve this problem.
1.55
Given: (IV) NO(g) + 1/2 O2(g)
NO2(g) + 13.5 kcal
(V) 1/2 N2(g) + O2(g) + 8.1 kcal
NO2(g)
What would the ∆H for this reaction be? (VI) 1/2 N2(g) + 1/2O2(g) - NO(g)? Make an enthalpy
diagram to help.
1.56
Summarize the ∆H for the reactions we have just examined in 1.53 and 1.55.
a. ∆HI =
b. ∆HII =
c. ∆HIII =
d. ∆HIV =
e. ∆HV =
f. ∆HVI =
g. Could we make a generalization about ∆H based on these examples? Try it.
21
h. How does this generalization relate to your conclusion at the end of 1.51 d?
1.57
Water can be decomposed into hydrogen and oxygen. How much energy is needed to break
up 6.0 grams of H2O(I) at 25°C and one atmosphere of pressure into its gaseous elements at
the same conditions of temperature and pressure?
H2O(l)
1.58
∆H = +
H2(g) + 1/2 O2(g)
Given: C2H4 + 3O2
2CO2 + 2H2O
57.8 kcal
mole H2O
∆H = -331.6 kcal/mole C2H4
How many kilocalories would be given off when 12.5 grams of ethylene (C2H4) completely
burned?
Adult check ___________________
Initial
Date
Review the material in this section in preparation for the Self Test. The Self Test will check
your mastery of this particular section. The items missed on this Self Test will indicate specific
areas where restudy is needed for mastery.
22
SELF TEST 1
Answer true or false (each answer, 1 point).
1.01
1.02
A negative ∆H means an exothermic reaction.
Changing colors is not an indication of a chemical change.
1.03
Gas formation is a phase not chemical change.
1.04
Reactants are the starting materials in a chemical reaction.
1.05
Bond formation releases energy.
1.06
Endothermic means to lose energy.
1.07
The formation of a precipitate is an indication of a chemical change.
1.08
All reactions are exothermic.
1.09
A temperature change in a reaction indicates a physical change.
1.010
Breaking bonds is an endothermic process.
Circle the letter of the best response (each answer, 3 points).
1.011
Which of the following reactions is not one of the general types of chemical reactions studied?
a. Precipitate formation
b. Change of color of a solution when it is diluted
c. Dissolving NaOH(s) in H2O
d. Gas evolution
Questions 1.012-1.014 relate to the following information:
The two balanced equations (1) and (2) are for reactions in which gaseous carbon dioxide is
produced from the combustion of (1) solid carbon and (2) gaseous carbon monoxide.
1. C(s) + O2(g)
2. CO(g) + 1/2 O2(g)
1.012
CO2(g) + 94.0 kcal
CO2(g) + 67.6 kcal
On the basis of the information given in equation (1) and assuming no change in temperature
or pressure, one can correctly conclude that
a. the rate of reaction is rapid.
b. the total number of moles of products is the same as the total number of moles of
reactants.
c. the reaction is exothermic.
d. the weights of the products are greater than those of the reactants.
e. there will be an increase in the volume of the reactants and products taken together as the
reaction proceeds.
1.013
When 112 grams of carbon monoxide are consumed according to equation (2), which of the
following reactions occurs? (atomic weights: C = 12.0, O = 16.0)
a. 1.0 mole of carbon dioxide is produced.
b. 67.6 kcal of heat are generated.
c. 2.0 moles of oxygen are consumed.
d. 0.25 mole of carbon dioxide is produced.
e. 0.50 mole of oxygen is consumed.
23
1.014
Which of the following equations represents an exothermic reaction?
b. C
CS2 ∆H = 27,550 cal
C
(graphite)
c. C + 2S
(diamond)
2H2O + O2 ∆H = +58 kcal/mole H2O
d. CH4 + 2O2
CO2 + 2H2O + 212,800 cal
e. 2H2O
1.015
∆H = +0.45 kcal/mole C
2N2(g) + 3/2 H2(g)
a. NH3(g) + 12.0 kcal
The combustion of 22.4 liters of CO according to the reaction CO(g) + 1/2 O2(g)
kcal gives off how much heat?
a. 67.6 kcal
d. 1,514 kcal
b. 33.3 kcal
e. neither a,b,c, nor d
CO2(g) + 67.6
c. 135.2 kcal
1.016
In the reaction (1.015) how much heat is produced from 14 grams of CO?
a. 6,716 kcal
d. 1,514 kcal
b. 33.8 kcal
e. neither a, b, c, nor d
c. 135.2 kcal
1.017
∆H = -57.82 kcal/mole
Which of the following reactions are endothermic?
a. H2(g) + 1/2 O2(g)
H2O(g)
b. /2 N2(g) + O2(g) + 8.1 kcal
1
c. /2 N2(g) + /2 H2(g)
1
d. C
3
(diamond)
+ O2(g)
NO2(g)
∆H = -94.50 kcal/mole
NH3(g) + 11.0 kcal/mole
CO2
Complete the following questions (each answer, 3 points).
Using the diagram at the right, answer
Questions 1.018 through 1.021. ~ Z + Q
Z+Q
-∆H
Energy
X+Y
1.018
Which is more stable, Z + Q or X + Y?
1.019
Is the reaction, X + Y
1.020
Does the -∆H indicate that heat is absorbed or given off the system Z + Q
1.021
Diamond has a more compact and ordered structure than graphite. When converting graphite
to diamond, is the enthalpy increased or decreased?
Z + Q exothermic or endothermic?
24
X + Y?
Make the following calculations (each answer, 3 points).
1.022
Determine the ∆H value for each of the following reactions:
a. N2(g) + 3H2(g)
2NH3(g) + 22,000 cal
b. C2H4(g) + 3O2(g)
2CO2(g) + 2H2O(g) + 331 kcal
c. 2HI(g) + 2.4 kcal
O2(g) + C
∆H =
∆H =
P4O10 + 712 kcal
e. CO2(g) + 94.50 kcal
∆H =
∆H =
H2(g) + I2(g)
d. P4(g) + 5O2(g)
∆H =
(diamond)
Calculate the following answers (each answer, 5 points).
1.023
Given:
1. C(s) + O2(g)
2. CO(g) + 1/2 O2
CO2(g) + 94.0 kcal
CO2(g) + 67.6 kcal
On the basis of calculations using equations (1) and (2), find how much heat in kcal per mole
of carbon would be produced in the following reaction:
C(s) + 1/2 O2(g)
1.024
Given:
CO(g) +
C = O 192 kcal
O - H 111 kcal
kcal/mole carbon
C-H
O-O
99 kcal
119 kcal
H
Balanced
Equation
C=O+O-O
O=C=O+H-O-H
H
Calculate the ∆H for this reaction.
Score
Adult check
54
68
25
_______________________
_______________________
Initial
Date