CHM 1020 Path 4 Chapter 9 Study Pack Part IV

G. ___ First Law of Thermodynamics & Related Terms- Answer

G1. ____ Discussion Questions-Chapter 9 Answer

G2. ____Specific Heat Problem- Answer Sample 2

G3. ____Enthalpy Change with Phase Change Prob- Answer Sample 2

H. ____ Enthalpy Change in Chemical Reaction- Answer

H1.   ___Bond Making/Breaking Problem- Answer

H2. ____Introduction to Entropy and Spontaneity-

Part G: First Law of Thermodynamics and Related Terms

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.

 

You can 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?

You must understand the concept of an thermodynamic system and its surroundings: :

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.

 

 

There are three 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.

 

 

 

 

 

 

 

 

 

 

Chapter 9:

Part G First Law of Thermodynamics & Related Terms

 

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 G1: 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)      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

Hreaction = Σ(Hf (products) - ∆Hf (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 Hf.  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 (-)

Chapter 9:

Part G1 Discussion Questions

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 Hf . 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

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

 

 

 

 

 

 

 

 

 

 

 

 

 

Chapter 9

Part G2: Specific Heat/First Law Problems

 

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

 

 

 

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 G3: Enthalpy Change with Phase Change

 

 

 

 

 

 

 

 

 

 

 

 

Chapter 9

Part G4: Enthalpy change with Phase Change/Ice Cube Problem

 

Phase Change:

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:

 

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 G5: Enthalpy Change in Chemical Reaction

 

 

 

 

Chapter 9

Part G4: Enthalpy change in Chemical Reactions

 

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) 3 CO2 (g) + 4 H2O (l) ∆H = -2220 kJ

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Answer:

-50,500 kJ

 

Part G0: Bomb/Coffee Cup Calorimeter Problem:

 

 

 

 

 

Chapter 9

Part G0: Calorimeter Problems

 

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 G5: Hess Law /Heats of Reaction Problems

My favorite problems for Chapter 9 involve Hess's Law. 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. 

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.

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Chapter 9

Part G5: Hess Law of Constant Heat Summation

 

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

 

 

Chapter 9

Part G5: Hess Law of Constant Heat Summation Problem#3

 

Using the following equations (if necessary)

 

2CH4 (g) + 3 O2 (g) 2 CO (g) + 4 H2O (l) ∆H˚ = -1215 kJ

 

 

2C (s) + O2 (g) 2 CO (g) ∆H˚ = -221 kJ

 

 

C (s) + O2 (g) CO2 (g) ∆H˚ = -394 kJ

 

 

to calculate the enthalpy change for the reaction

 

CH4 (g) + 2 O2 (g) CO2 (g) + 2 H2O (l) ∆H˚ = ?

 

 

 

 

 

 

 

Part G6: Standard Enthalpies of Formation Problem-

We need to introduce 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:

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:

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.

 

 

 

 

Part H: Standard Enthalpies of Formation

 

Calculate H for the reaction:

 

2 Al (s) + 1 Cr2O3 (s) 1 Al2O3 (s) + 2 Cr (s)

f (Al2O3 (s) ) = -1676 kJ/mol f (Cr2O3 (s) ) = - 1128 kJ/mol

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Answer: -548 kJ

 

 

 

 

 

 

Chapter 9:

Part G6: 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) 2 NO2 (g) + 3 H2O (l) ∆H˚ = -698 kJ

 

Calculate the standard molar enthalpy of formation of ammonia if

 

f (H2O(l) ) = -286 kJ/mol f (NO2(g) ) = + 33 kJ/mol

 

 

 

 

 

 

 

 

 

 

 

 

Part I: Bond Making/Breaking Problem-

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 a video in Chapter 9.

Since the author assumes you do not know Dot Structures of Molecules (which you do from Chapter 6). There is table 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)

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Chapter 9 Part G7: Heat of Reaction from Bond Energies

 

Methane burns in oxygen to produce heat for homes by the following reaction:

 

CH4 + 2 O2 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)

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Chapter 9 continued

Part G7: Heat of Reaction from Bond Energies

 

Propylene burns in oxygen to produce heat by the following reaction:

 

2 C3H6 + 9 O2 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):

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Chapter 9:

Part H: Introduction to Entropy and Spontaneity-

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 C4H10 + C4H8 + C3H6

 

Is ΔS for this reaction likely to be positive or negative?

 

 

Explain!

 

 

 

 

Chapter 9

Part H1: Introduction to Free Energy and Spontaneity-

 

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