CHM
2045C Name: _________________
Module Seven Homework Packet-Jespersen-Chapter 6
Module Seven:
Thermochemistry & States of Matter (Chap
8 McMurry) ppt
Assignment Outline (Kotz)
A. ____(03) First Law of
Thermodynamics & Related Terms-6.5
Answer
B. ____(03) Discussion Questions-Chapter
6 Answer
C. ____(04) Specific Heat Problem-Section
6.3 Answer Sample 2
D. ____(05) Enthalpy Change with
Phase Change Prob-Sect 8.7 Answer Sample 2
E. ____(05) Enthalpy Change in
Chemical Reaction-Section 6.4 Answer
F. ____ (00) Bomb Calorimeter
Section 6.6 Answer
G. ____(10) Hess Law
/Heats of Reaction Problem-Section 6.8 Answer
H. ____(10) Standard
Enthalpies of Formation Problem-Section 6.9 Answer
I. ____(10) Bond Making/Breaking
Problem-Section 18.10 Answer
J. ____ (00)
Introduction to Entropy and Spontaneity-Section 18.3
L. ____ (00)
Introduction to Free Energy and Spontaneity-Section 18.4
______(50) Total = ______%
Part A: First Law of Thermodynamics and Related Terms
To begin the study of the energies of
chemical reactions, first read section 6.1 and 6.2. There are three laws of
thermodynamics, of which only the first is presented in CHM 2045C in Chapter 6
(Section 6.5). In CHM 2046C in Chapter 18 we have to wait to be exposed
to the 2nd and 3rd laws of Thermodynamics. After studying Entropy in Section
18.3, the 2nd Law is discussed in Section 18.4. Then Gibbs Free Energy is introduced in Section 18.4, so that Section
18.5 focuses on the 3rd Law of thermodynamics.
The book clarifies energy concepts from
CHM 1025C in Sections 6.1 and 6.2. Energy can be classified as kinetic and
potential. Kinetic energy is energy associated with motion, while Potential
energy is stored energy and can be converted to kinetic energy. The sum
of all the kinetic and potential energy in the universe is the Total Energy of
the Universe.
Examples of kinetic energy are:
Thermal Energy,
Mechanical Energy,
Radiant Energy,
Electrical Energy, and
Sound.
Examples of Potential Energy are:
Gravitational Energy,
Nuclear Energy,
Chemical Potential Energy, and
Electrostatic Energy.
I have placed a list of the forms of
energy in the chapter 6 vocabulary list with examples: The list is found at:
http://www.fccj.us/chm2045/Kotz6eVocabulary/45Chap6Vocabulary.htm
On the bottom of page 254 and top of 255 the discussion of the Law of Conservation of
Energy is reviewed. You can also go to the following web sites for forms and
conservation of energy:
http://www.eia.doe.gov/kids/energyfacts/science/formsofenergy.html
You should be able to identify forms of energy associated with energy changes.
What is the difference between heat and
temperature? Bottom of page 255 and page 256 discusses the two.
You must understand the concept of an thermodynamic system and
its surroundings explained in Section 6.3 on
page 258. and the directionality of heat transfer on
pages 259-261. The book introduces to
the units of energy, the Joule and the calorie on page 254-255.
For Part A, you need to be able to
define the following:
1. State the first law of
thermodynamics:
The
energy of the universe is a constant. [Law of Conservation of Energy-Energy can
neither be created or destroyed, but can be converted from form to form]
2. Explain a Thermodynamic System
and Its Surroundings
A thermodynamic
system is defined as the object, or collection of objects,
being studied.
The surroundings
include everything outside the system that exchange energy with
the system.
In
section 6.3 there is a discussion about the types of systems: Open Systems;
Closed Systems; and Isolated Systems:
Open Systems can gain or lose
energy across their boundaries.
Closed Systems can absorb or
release energy, but not mass, across the boundary. The mass of a closed system
is a constant, no matter what happens inside.
Isolated Systems cannot exchange
matter or energy with its surroundings. (Adibatic)
3. Define endothermic and exothermic
processes.
In the exothermic process heat is transferred from a
system to the surroundings.
An endothermic
process is the opposite of an exothermic process: heat is
transferred from surroundings of the system.
Module
Seven:
Part A First Law of
Thermodynamics &
Related Terms 5 points
State the first law of thermodynamics:
Explain a Thermodynamic System and Its
Surroundings:
Define endothermic and exothermic
processes:
State the second law of
thermodynamics:
State the third law of thermodynamics:
Part B: Discussion Questions
After studying the remainder of the
chapter you should be able to answer any two of the following discussion
questions:
1. What is the standard state of an element or compound substance and
give an example?
The standard
state of an element or a compound is defined as the most stable
form of the substance in the physical state that exists at a pressure of 1 bar
and specified temperature (usually 25oC or 298 K).
For example
∆Hof for CO2
(g):
At 25 oC
and 1 bar, the standard state of carbon is solid graphite, the most stable form
of this element and the most stable form of oxygen is O2 (g)
C(s) + O2(g) à CO2(g)
∆Hof = -393.5 kJ
2. Why does water have a high specific heat capacity? What does this mean?
The specific heat of water is much
larger than for most substances because of the unusually strong bonds between
the water molecules (Look up the hydrogen bond in later chapters). These
intermolecular bonds are progressively broken as more and more heat is added. What this means is that a considerable quantity of
heat is required to heat water and considerable amount of heat must be
transferred out of the water before it cools down appreciably.
3. Write four different
mathematical expressions for the 1st Law of
Thermodynamics. How are they related?
Some expressions for the 1st
law are:
∆E = q + w
where
∆E refers to the system
qin = qout
heat
gained = heat lost
∆E = zero
where
∆E refers in this case to the entire universe
∆Hºreaction = Σ(∆Hºf (products) - ∆Hºf (reactants) )
All these expressions represent an energy balance,
reflecting the fact that energy can neither be created nor destroyed.
4. Where does the energy come from in
an endothermic process? And where does it go?
In an endothermic process energy is
required. There are two sources for this energy: the energy may come from the
surroundings if the system is heated; or the energy could come from the system
itself if the kinetic energy of the atoms and molecules of the system is
reduced. In this case, the temperature of the system decreases. Unless the
system is isolated (well-insulated), there will be a movement of the energy
between the system and its surroundings to reestablish thermodynamic
equilibrium.
5. Define standard molar enthalpy of formation ∆Hºf.
Why is the standard enthalpy of formation of a pure element
in its most stable form defined as zero (page 284 and see table 6.2 p 285)?
The standard molar enthalpy of formation of a
substance is the enthalpy change for a reaction in which one mole of the
substance in its standard state is made from its constituent elements in their
standard states.
For a substance that is an element,
such a reaction represents no change, and therefore then enthalpy change must
be zero because the element (or atom) already exists in nature and can not be
assembled by man from its building blocks of subatomic particles. Elements are
defined as the smallest unit of matter that has the chemical properties of that
matter. It cannot be subdivided into it building blocks by any chemical means.
Therefore, we state energy change begins with putting atoms together to make
molecules of compounds.
6. Define a spontaneous reaction.
How can you tell whether a reaction is spontaneous?
A spontaneous reaction is a
reaction that happens by itself. It may happen quickly or very very slowly but it does happen. Calculations in
thermodynamics can be done to determine whether or not a reaction is
spontaneous. However, if a process or reaction does happen itself, you can be certain it is spontaneous.
7. A system can exchange energy with its surroundings either by transferring heat or by doing work. This is expressed by the following equation: Δ E = q + w
Fill in the chart with correct signage:
Change Sign Conversion Effect of Esystem
Work done on the system by surroundings w > 0 (+) E increases (+)
Work done by the system on surroundings w < 0 (-) E decreases (-)
Heat transferred to system from surroundings q > 0 (+) E increases (+)
Heat transferred from system to surroundings q < 0 (-) E decreases (-)
Module Seven:
Part B Discussion
Questions 5 points
For the exam, your
instructor will select four of the following questions for you to write the
answers:
1. What is the standard state of an element or compound
substance?
2. Why does water have a high specific heat capacity? What does this mean?
3. Write four different
mathematical expressions for the 1st Law of Thermodynamics. How are they related?
4. Where does the energy come
from in an endothermic
process? And where does it go?
5. Define standard molar enthalpy of formation ∆ Hºf . Why is the standard enthalpy of formation of a
pure element in its most stable form defined as zero?
6. Define a spontaneous reaction. How
can you tell whether a reaction is spontaneous?
7. A system can exchange energy with its surroundings either by
transferring heat or by doing work. This is expressed by the following
equation: Δ E = q +
w
Fill in the chart with correct signage:
Change
Sign Conversion Effect of Esystem
Work done on the system by
surroundings
Work done by the system on
surroundings
Heat transferred to system
from surroundings
Heat transferred from system
to surroundings
Part C: Specific Heat Problem
From Parts C through Part I we will focus
on the mathematical problems which are presented in Chapter 6 (and section 18.10-Bond
Energies).
Sections 6.3 and 11.6 present the
problems associated with Part C. Read these sections. Study Example 6.1
page 261, Example 6.2 pages 262-3, and Example
6.3 p 263.
Then work the Exercises: 6.4, 6.5 and
6.6 page 264. At the end of the chapter there are 6 problems for you to
work for Part C: Measuring Problems #6.21-6.26 page 293.
Under section 6.5 there is the discussion
of changes in the heat content of the system using the equation 6.10 on page 269.
One of the Part C problems will be a simple calculation of heat content change
using this equation. but the key is to understand the
signs applied:
If the internal energy of a
thermodynamic system is decreased by 300 J when 75 J
of work is done on the system, how much heat was transferred, and in which
direction, to or from the system. See section 6.4 p 253
Change
Sign
Conversion Effect of Esystem
Work done on the system by
surroundings w > 0 (+)
E increases
Work done by the system on
surroundings w < 0 (-) E decreases
Heat transferred to system from
surroundings q > 0 (+) E
increases
Heat transferred from system to
surroundings q < 0 (-) E decreases
Δ E = q + w
Given Δ E= - 300 J w
= + 75 J
-
300J = q + (+75J)
q =
- 375 J of heat was transferred from the system to the surroundings
Module
Seven
Part
C: Specific Heat/First Law Problems 05
points
If the temperature of a 50.0 gram
block of aluminum increases by 10.9 K when heated by 500 joules, calculate the:
If the internal energy of a
thermodynamic system is decreased by 300 when 75 J of
work is done on the system, how much heat was transferred, and in which
direction, to or from the system
Answers:
a.
heat capacity of the aluminum block = 45.9 J/K
b.
molar heat capacity of aluminum 24.8 J/K mol
c. specific heat capacity of aluminum =
0.917 J/Kg
-375 J was transferred From the system
Part D: Enthalpy Change with Phase Change
Section 11.3 presents the problems associated with
Part D. Read this section. Study Example 11.2 page 539.
Then work the Exercises: 11.10 and
11.11 page 540. At the end of the chapter 11 there are 6 problems for you
to work for Part D: Energy and Changes of State #11.50-11.56 page 568.
Module
Seven
Part
D: Enthalpy change with Phase Change/Ice Cube Problem
5 points
Phase
Change (3 points):
1.
Calculate the amount of heat necessary to melt 27.0 grams of ice at 0oC, if the heat of fusion of ice is 333 J/g.
If I had the same amount of water at 100oC,
calculate the amount of heat required to boil 27.0 grams of water if the heat of vaporization of water is 2256 J/g?
How much heat is required to raise the
temperature of the 27 grams of water at
0oC to 100oC, if the specific heat of water is 4.184 J/goC
Ice
Cube Problem (2 points):
If 27.0 grams of ice at 0oC is added
to an insulated cup of water containing
123 grams of water at 50oC. What will be the final thermodynamic
equilibrium temperature of the water/ice mixture assuming no heat is lost to
the surroundings?
Part E: Enthalpy Change in Chemical Reaction
Section 6.4 and 6.9 demonstrates the
Enthalpy change in Chemical reactions. Look at example 6.9 p283 and work
6.21 and 6.22 p284.
At the end of the chapter there are four
problems which resemble Part E calculation: Enthalpy Problems #6.63-6.66 page 294-5.
Module
Seven
Part
E: Enthalpy change
in Chemical Reactions 5 points
If the
enthalpy change for the combustion of propane
gas, C3H8 (g) is -2220kJ/mol
propane. What quantity of heat is released when 1.00 kg of propane is burned?
C3H8
(g) + 5 O2 (g) à 3 CO2 (g) + 4 H2O
(l) ∆H
= -2220 kJ
Answer:
-50,500 kJ
Part F: Bomb/Coffee Cup Calorimeter
Problem: Section 6.6
Section 6.6 is devoted to measuring the
heat gains and losses through a calorimeter. Two types of calorimeters are
discussed. One which we can do in our lab, coffee cup calorimeter, is
sometime called a Constant Pressure calorimeter since the measurements are
carried out at lab pressure. Read pages 272-273 and study the discussion on pages
273-274-275 which describes the coffee cup example. In this lab we assume there
is no heat loss to the surroundings and we measure ∆H for the reaction.
The second type of calorimeter, which is
usually not performed in lab untill Physical Chemistry courses, is a Bomb
calorimeter. Results are much more accurate with this calorimeter. Since the
"bomb" is confided to a closed space, this calorimeter is known as a
Constant Volume Calorimeter. Read pages 270-272. Study Example 6.5 on page 271,
then work Exercise 6.11 and 6.12 on page 272.
A thought question for the Discussion
Board would be in the coffee cup calorimeter we are measuring ∆H, while a Bomb calorimeter we are measure ∆E. Why? At the end of the chapter you should work problems #6.63-6.66
as additional examples. On the exam you will have one calculation, either of
the two calorimeters.
Module
Seven
Part F: Calorimeter Problems 5 points
Benzoic acid (C6H5COOH) is sometimes used as a standard to
determine the heat capacity of a bomb calorimeter (constant volume). The
calorimeter is an insulated containing with 1.20 kg of water. When 1.32
g of benzoic acid is burned in a calorimeter that is being calibrated , the temperature rises from 20.93 oC to 22.93 oC. What is the heat capacity of the calorimeter?
The heat of combustion of benzoic acid (qv)
is -26.42 kJ/g
capacity of the calorimeter? The heat of
combustion of benzoic acid (qv) is -26.42 kj/g
Part G: Hess Law /Heats of Reaction Problems
My favorite problems for Chapter 6 involve Hess's Law: Section 6.8. Hess's
Law states : if a reaction is the sum of two or more
other reactions, then ΔH for the overall process is the sum of the ∆H values for those
reactions. Read the section and discussion on pages 277-281. The
interesting part of these problems is to look at a set of reactions, two,
three, four, or five. Then look at the desired reaction for which the ∆H is unknown. The fun part is
these sets of reactions are kind of a puzzle maze.
You can do two things to a reaction to help you solve the problem.
1. You can rewrite the problem by writing the reverse reaction, making
the products now the reactants and the reactants now the products. All you have
to so is change the sign of ∆H.
2. You can multiple any reaction through by a coefficient or a fraction,
and all you have to do is also multiple that ∆H by the same coefficient or fraction.
Once you have rearranged the equations in a set, they should add up to
the unknown or net equation.
Study example 6.8 on page 281 and work Exercise 6.19 on page 282-3.
Every college chemistry text has a neat set of Hess's Law problems. Our text
has ten examples at the end of the chapter Problems #6.79-6.88 on pages 296-297.
Read
Jespersen 7th Section 6.8 Pages 277-283
See Example 6.11 Page 287
Try
Practice Exercises 6.25-6.27 page 288
There
are 10 good Problems in Additional Exercises Pages 296-7 Problems 6.79-6.88
Module
Seven
Part
G: Hess Law of Constant Heat Summation
10 points
Using
the following equations:
S (s) + 3/2 O2 (g) → SO3 (g) ∆
Ho = -395.2 kJ
2 SO2 (g) +
O2 (g) → 2 SO3 (g) ∆ Ho = -198.2 kJ
calculate the ∆ Ho for the reaction:
S (s) +
O2 (g) → SO2 (g)
Given
the following equations:
B2O3 (s) + 3H2O (g) →
B2 H6 (g)
+ 3 O2 (g) ∆ Ho = +2035 kJ
H2O (l) → H2O (g) ∆ Ho
= +44 kJ
2 B (s)
+ 3 H2 (g) →
B2 H6
(g) ∆ Ho
= +36 kJ
H2 (g) + 1/2 O2 (g)
→
H2O
(l) ∆ Ho = -286 kJ
Calculate
the ∆ Ho for the reaction:
2 B (s)
+ 3/2 O2 (g)
→ B2 O3 (s)
Answer: -1273 kJ
Module Seven
Part
G: Hess Law of Constant Heat Summation
Problem#3
Using the following
equations (if necessary)
2CH4 (g) +
3 O2 (g) à 2
CO (g) + 4 H2O (l) ∆H˚ = -1215 kJ
2C (s)
+ O2 (g) à 2 CO (g) ∆H˚ = -221 kJ
C (s)
+ O2 (g) à CO2 (g) ∆H˚ = -394 kJ
to calculate the enthalpy change for the
reaction
CH4 (g) +
2 O2 (g) à CO2 (g) + 2 H2O (l) ∆H˚
= ?
Part H: Standard Enthalpies of Formation Problem-Section
6.9
Section 6.9 introduces the concept
standard molar enthalpies of formation and the standard state. Always an
interesting question for the discussion board is a statement for the discussion
board is found on page 266:
The standard enthalpy of formation
for an element in its standard state is defined as ZERO. Why?
To calculate the enthalpy change for a
reaction you will use the following formula:
From equation 6.14 page 287:
∆H˚rxn = ∑[∆H˚f (products)]- ∑[∆H˚f (reactants)]
Look closely at the problems. The trick is
always information seems to be left out. The standard enthalpies for formation
of elements are never given in the problem because its value is ZERO. Don't be
fooled. Also be careful about you signs as you make you summations. Study carefully
Example 6.11 on page 287 then work Exercises 6.25-6.27 on page 288.
The author feels strong about
this problem type. At the end of the chapter he provides you with 8 homework type
problems: #6.89-96 page 297
Book Reference:
Section 6.9 Pages 283-288
See Worked Example
6.9 and 6.10 page 283-284
Try Practice
Exercises 6.23-6.24 Page 284
Try End of Chapter
Additional Problems 6.89-696 Page 297
Part H: Standard Enthalpies of Formation 5 points
Calculate
∆ H for the reaction:
2 Al (s) + 1 Cr2O3
(s) à
1 Al2O3 (s) + 2 Cr (s)
∆H˚f (Al2O3 (s) ) = -1676 kJ/mol ∆H˚f (Cr2O3
(s) ) = - 1128 kJ/mol
Answer: -548 kJ
Module
Seven:
Part H:
Standard Enthalpies of Formation continued
When ammonia is
oxidized to nitrogen dioxide and water, the quantity of heat released equals
349 kJ per mole of ammonia:
2NH3 (g) +
7/2 O2 (g) à 2
NO2 (g) + 3 H2O
(l) ∆H˚ = -698 kJ
Calculate the standard molar enthalpy of
formation of ammonia if
∆H˚f (H2O(l) )
= -286 kJ/mol ∆H˚f
(NO2(g) ) = + 33 kJ/mol
Part I: Bond Making/Breaking Problem-Section 18.10
In addition to the several calculations
above where you are trying to find the Enthalpy of Reaction, there is another
method which is explained in Section 18.10 of Chapter 18.
In Section 6.4 the author
suggest this method, but decided to wait until Chapter 18 to fully
discuss. Since the author assumes you do not know Dot Structures of Molecules
(which you do from Module 4). You should Read Section 8.5 where the physical
properties of covalent bonds are discussed: Bond Order, Bond Length and Bond
Energy. Then the author waits untill chapter 18 for another interesting
calculation. Go to Section 18.10 See the subsection Bond Energy which begins on
page 890. On page 892 there is table 18.4 which gives the 'average' bond
energies for many types of covalent bonds, especially organic molecules. The
skill need here is to be able to sketch the dot structure (stick structure is
ok) of each reactant and each product. The formula which you will use to
calculate the heat of reaction:
ΔHorxn=
Σ (bonds broken) – Σ (bonds formed)
Try Exercise 18.30 and 18.31 on page 893.
There are several end of chapter problems as examples
of Part I Problems. Work #18.105-18.112 on Pages 900-901 for
additional practice.
Module 7 Part I: Heat of
Reaction from Bond Energies
10 points
Methane burns in oxygen to produce heat for homes by the following reaction:
CH4 +
2 O2 è CO2 + 2 H2O
If the average bond energies in kJ/mol are:
O-O 146
O=O 498
O-H 463
C-C 346
C-H 413
C-O 358
C=O 745
H-H 436
Calculate ∆Horxn for the reaction:
ΔHorxn
= Σ (bonds broken) – Σ (bonds formed)
Module 7
Part I: Heat of
Reaction from Bond Energies continued
Propylene burns in
oxygen to produce heat by the following reaction:
2 C3H6 +
9 O2 è 6 CO2 + 6
H2O
If the average bond
energies in kJ/mol are:
O-O 146
O=O 498
O-H 463
C-C 346
C-H 413
C-O 358
C=O 745
H-H 436
C=C 134
ΔHorxn
= Σ (bonds broken) – Σ (bonds formed)
Calculate ∆Horxn for the reaction (Hint draw the dot/stick
structures of the compounds):
Module
Seven:
Part J:
Introduction to Entropy and Spontaneity-
Section 8.12, 8.13 5 Points
What does entropy
measure?
How is it possible
for a reaction to be spontaneous yet endothermic?
Tell whether the entropy changes for the following process are likely
to be positive or negative?
(a) The fizzing of a newly opened can of soda?
(b) The growth of a plant from seed?
One of the steps in the cracking of petroleum into gasoline involves
the thermal breakdown of large hydrocarbon molecules into smaller ones. For
example the following reaction might occur:
C11H24 à C4H10 + C4H8 +
C3H6
Is
ΔS for this
reaction likely to be positive or negative?
Explain!
Module
Seven
Part L: Introduction
to Free Energy and Spontaneity-Section 8.13 5 points
What are the two terms that makeup the free-energy change for the
reaction, ΔG,
and which of the two is usually more important?
Tell whether reactions with the following values of ΔH and ΔS are spontaneous or non spontaneous
and whether they are exothermic or endothermic?
(a) Δ H = -48 kJ; ΔS = +135 J/K at 400K
(b) Δ H = -48 kJ; ΔS = -135 J/K at 400K
(c) Δ H = +48 kJ; ΔS = +135 J/K at 400K
(d) Δ H = +48 kJ; ΔS = -135 J/K at 400K