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 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

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,
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:

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

reaction = Σ(f (products) - ∆f (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 f.  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:

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        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:

1. heat capacity of the aluminum block.

1. molar heat capacity of aluminum.

1. specific heat capacity of aluminum.

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

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

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?

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

-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)

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:

rxn  = ∑[∆f (products)]-  ∑[∆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

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