ap chemistry: chapters 6 & 17 notes...

16
AP Chemistry: Chapters 6 & 17 Notes Thermochemistry 6.1 The Nature of Energy In this chapter, we will study energy in terms of chemical potential energy and how that energy can change form to accomplish work. We will focus specifically on heat and energy transfer. Energy- the capacity to do work or to produce heat Work- force acting over a distance o It involves a transfer of energy 1st Law of Thermodynamics- also known as the Law of Conservation of Energy. States that energy can be converted from one form to another but it can be neither created nor destroyed. o The total amount of energy in the universe is constant. Energy can be classified in two ways: o Potential energy- energy due to position or composition (included chemical potential energy) o Kinetic energy- energy due to the motion of an object Kinetic energy is dependent on the mass and velocity of an object Heat- involves a transfer of energy between two objects due to a temperature difference. o Heat always moves from warmer matter to cooler matter. Temperature- a property that reflects random motions of the particles of a particular substance o Exothermic- reaction which releases heat Energy flows out of the system Potential energy is changed to thermal energy Products have lower potential energy than reactants o Endothermic- reaction which absorbs heat Energy flows into the system Thermal energy is changed into potential energy Products have higher PE than reactants In terms of a chemical reaction, the system is our reaction. The surroundings are everything else, including things like the container the reaction occurs in, the room it sits in, etc. Internal energy (E) of a system is the sum of the kinetic and potential energies of all the particles in a system. Thermodynamic quantities always consist of a number and a sign (+ or -). The sign represents the systems point of view. (Engineers use the surroundings point of view) o Exothermic -q (systems energy is decreasing) o Endothermic +q (systems energy is increasing) o Example: Calculate E if q = -50 kJ and w = +35kJ. For a gas that expands or is compressed, work can be calculated by: o w = -PV o units: L . atm, conversion 101.3 J/L-atm o Example: Calculate the work if the volume of a gas is increased from 15 mL to 2.0 L at a constant pressure of 1.5 atm. KE = ½ mv 2 m = mass in kg v = velocity in m/s units are J, since J = (kg . m 2 )/s 2 E = q + w E is the change in the systems internal energy q represents heat w represents work usually in J or kJ

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Page 1: AP Chemistry: Chapters 6 & 17 Notes Thermochemistrymhs-ap-chemistry-2014-2015.weebly.com/.../01_thermo_student_note… · AP Chemistry: Chapters 6 & 17 Notes Thermochemistry ... Simple

AP Chemistry: Chapters 6 & 17 Notes

Thermochemistry

6.1 The Nature of Energy In this chapter, we will study energy in terms of chemical potential energy and how that energy can change form to

accomplish work. We will focus specifically on heat and energy transfer.

Energy- the capacity to do work or to produce heat

Work- force acting over a distance

o It involves a transfer of energy

1st Law of Thermodynamics- also known as the Law of Conservation of Energy. States that energy can be

converted from one form to another but it can be neither created nor destroyed.

o The total amount of energy in the universe is constant.

Energy can be classified in two ways:

o Potential energy- energy due to position or composition (included chemical potential energy)

o Kinetic energy- energy due to the motion of an object Kinetic energy is dependent on the mass and velocity

of an object

Heat- involves a transfer of energy between two objects due to a

temperature difference.

o Heat always moves from warmer matter to cooler matter.

Temperature- a property that reflects random motions of the particles of a particular substance

o Exothermic- reaction which releases heat Energy flows out of the system Potential energy is changed to thermal energy

Products have lower potential energy than

reactants

o Endothermic- reaction which absorbs heat Energy flows into the system Thermal energy is changed into potential energy

Products have higher PE than reactants

In terms of a chemical reaction, the system is our reaction. The

surroundings are everything else, including things like the

container the reaction occurs in, the room it sits in, etc.

Internal energy (E) of a system is the sum of the kinetic and

potential energies of all the particles in a system.

Thermodynamic quantities always consist of a number and a sign

(+ or -). The sign represents the systems point of view.

(Engineers use the surroundings point of view)

o Exothermic -q (systems energy is decreasing)

o Endothermic +q (systems energy is increasing)

o Example: Calculate E if q = -50 kJ and w = +35kJ.

For a gas that expands or is compressed, work can be calculated by:

o w = -PV

o units: L.atm, conversion 101.3 J/L-atm

o Example: Calculate the work if the volume of a gas is increased from 15 mL to 2.0 L at a constant pressure of 1.5 atm.

KE = ½ mv2

m = mass in kg

v = velocity in m/s

units are J, since J = (kg.m

2)/s

2

E = q + w E is the change in the system‟s internalenergy

q represents heat

w represents work usually in J or kJ

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6.2 Enthalpy and Calorimetry

Enthalpy (H) concerns the heat energy in a system.

H = q at constant pressure only

At constant pressure, the terms heat of reaction and

change in enthalpy are used interchangeably.

The change in the enthalpy of a system can be calculated using:

o H= H products - H reactants

For an exothermic reaction, H is negative

For an endothermic reaction, H is positive

o Example: For the reaction 2Na + 2H2O 2NaOH + H2 , H = -368 kJ Calculate the heat

change that occurs when 3.5 g of Na reacts with excess water.

Calorimetry- the science of measuring heat flow in a chemical reaction.

o It is based on observing the temperature change when a body absorbs or discharges heat.

o The instrument used to measure this change is the calorimeter.

The heat capacity of H2O is 4.18 J/g oC

Constant pressure calorimetry-

pressure remains constant during the

process

Simple calorimetry- used to

determine heats of reaction

The primary reaction to calculate

heat changes in a system is the

“Mcat” equation.

o Heat lost = -heat gained

o Energy released as heat = (heat capacity) (mass of solution )

(increase in temp)

o Example: A 30.0 g sample of water at 280. K is mixed with 50.0 g

water at 330K. Calculate the final temperature of the mixture

assuming no heat loss to the surroundings.

Heat Capacity (C) C = heat absorbed

Increase in temp.

C = J/g°C J/mol°C

H = E + PV E is internal energy

P is pressure

V is volume

q = mCT J = (J/g

oC )(g)(°C)

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o Example: A coffee cup calorimeter contains 150 g H2O at 24.6oC. A

110 g block of molybdenum is heated to 100oC and then placed in the

water in the calorimeter. The contents of the calorimeter come to a

temperature of 28.0oC. What is the heat capacity per gram of

molybdenum?

In terms of calorimetry, we can describe certain properties of the reaction as an:

o Extensive property- this depends on the amount of substance (ex. Heat of reaction)

o Intensive property- doesn‟t depend on the amount of substance (ex. Temperature)

Calorimetry can be done in a closed, rigid container. This is called constant volume calorimetry

o Ex. Flashbulb in a camera or a bomb calorimeter o No work can be done since the volume doesn‟t change

o Heat evolved = T x heat capacity of calorimeter (energy required to change the temp 1oC)

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r

6.3 Hess’s Law

Hess‟s Law- States that the change in enthalpy from products to

reactants, or H, is the same whether the reaction occurs in one

step or in several steps.

o H is not dependent on the reaction pathway.

o The sum of the H for each step equals the H for the

total reaction.

o If a reaction is reversed, the sign of H is reversed. o If the coefficients in a reaction are multiplied by an

integer, the value of H is multiplied by the same integer.

o Example: Given the following reactions and their respective enthalpy changes, calculate H for the

reaction: 2C + H2 C2H2.

o Example: The heat of combustion of C to CO2 is -393.5 kJ/mol of CO2, whereas that for combustion

of CO to CO2 is -283.0 kJ/mol of CO2. Calculate the heat of combustion of C to CO.

6.4 Standard Enthalpies of Formation

Standard enthalpy of formation (H f ) – The change in enthalpy that accompanies the formation of one

mole of a compound from its elements with all substances in their standard states at 25oC.

The degree sign on a thermodynamics function indicates that the process it represents has been carried out at

standard state conditions.

o Standard States

for gases, pressure is 1 atm

for a substance in solution, the concentration is 1 M

for a pure substance in a condensed state (liquid or solid), the standard state is the pure liquid

or solid.

for an element, the standard state is the form in which the element exists under conditions of

1 atm and 25oc.

Values of Hf are found in your handout or Zumdahl Appendix 4

Horeaction = Hf

oproducts - Hf

oreactants

o Example: The standard enthalpy change for the reaction CaCO3(s) CaO + CO2(g) is 178.1 kJ

Calculate Hof for CaCO3

C2H2 + 5/2 O2 2CO2 + H2O H = -1299.6 kJ/mol C2H2

C + O2 CO2 H = -393.5 kJ/mol C

H2 + ½ O2 H2O H = -285.9 kJ/mol H2

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8.8 Covalent Bond Energies

During reactions, bonds of the reactants

are broken and the bonds of the product

are formed. The overall energy change

is based on the difference in energy of

the bonds broken vs the bonds formed.

H = nHbonds broken - nHbonds formed

Where “n” is the number of bonds and

D is the bond energy perm mole of

bonds. Bond energies are given in the

table to the right from Zumdahl.

Bond energies vary based on the

elements involved, but double bonds

have more energy than single bonds

BUT NOT TWICE AS MUCH.

Example: Use bond energies to

calculate the change in energy that accompanies the following reaction:

H2 + F2 2HF

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o S

G f f

Chapter 17: Thermochemistry

17.1 & 17.2 Spontaneous Processes and Entropy and the Second Law of Thermodynamics The first Law of Thermodynamics states that energy is neither created nor destroyed; it is constant in the universe.

We can measure energy changes in chemical reactions to help determine exactly what is happening and how it is

occurring.

Keep in mind that thermodynamics deals with the reactants and products while kinetics deals with how reactants

become products.

Spontaneous process- occurs without outside intervention (ex. Rusting)

o This may be fast or slow (CO2 sublimes vs. iron rusting)

Entropy (S)- a measure of randomness or disorder

o This is associated with probability. (there are more ways for something to be disorganized than

organized)

o Entropy increases going from a solid to a liquid to a gas and when solutions are formed.

o Entropy increases in a reaction when more atoms or molecules are formed.

o Entropy increases with increasing temperature.

2nd Law of Thermodynamics- In any spontaneous process there is always an increase in the entropy of the

universe. The energy of the universe is constant but the entropy of the universe is increasing.

17.3 and 17.4 Free Energy

Free energy (G)- the amount of energy available to do work.

o Free energy change is a measure of the spontaneity of a reaction. It is the

maximum work available from the system. o A spontaneous reaction carried out as constant temperature

and pressure has a negative G. For example, when ice melts

H is positive (endothermic), S is positive and G = 0 at

0˚C.

17.5 Entropy Changes in Chemical Reactions

Third Law of Thermodynamics- The entropy of a perfect crystal at 0 K is zero. o

reaction o products

o reactants

o Ex. Given the following standard molar entropies, calculate So

for the reaction:

2Al(s) + 3MgO(s) 3Mg(s) + Al2O3(s) Al(s) = 28.0 J/K MgO(s) = 27.0 J/K Mg(s) = 33.0 J/K Al2O3(s) = 51.0 J/K

17.6 Free Energy and Chemical Reactions

Standard free energy change (Go) -The change in free energy that occurs if the reactants in their standard

states are converted to products in their standard states. o o

o Go

=Gf products - Gf reactants at standard conditions o

o Gf for a free element in its standard state is zero.

o Ex. Given the equation N2O4(g) 2NO2(g) and the following data, calculate Go.

o for N2O4(g) = 97.82 kJ/mol, G

o for NO2(g) = 51.30 kJ/mol

G = H -TS

= S - S

S H G Spontaneous?

+ - - Yes, always

+ + ? Yes, at high T

- - ? Yes, at low T

- + + No, never

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G = Go

+ RT ln (Q) Q = reaction quotient (partial pressure of products/reactants raised to the

power of their coefficients)-only pressures of gases are included. T = temperature in Kelvin

R = gas constant 8.314 J/K.mol

Go

= free energy change at 1 atm (be sure to change to Joules!)

The Gibbs-Helmholtz equation works like the Gibbs Free Energy equation, just at standard state conditions.

o Go

=Ho

-TSo

(When working this, change the units for S to kJ).

o Ex. For the given reaction and the following information, calculate Go

at 25°C.

2PbO(s) + 2SO2(g) 2PbS(s) + 3O2(g)

17.7 The Dependence of Free Energy on Pressure

In an ideal gas, enthalpy does not depend on pressure. Entropy, however, does depend on pressure because it depends upon volume. Gases in large volumes have greater

entropy than in a small volume. Gases at a low pressure have greater entropy than at a high pressure. Because

entropy depends on pressure, G of an ideal gas depends on

its pressure.

If we incorporate PV = nRT with the

equation for free energy, we end up with

an equation to calculate free energy in

relation to temperature and pressure

variables.

o Ex. Calculate G at 298K for the following reaction if the reaction mixture consists of 1.0 atm N2, 3.0 atm H2, and 1.0 atm

NH3. N2(g) + 3H2(g) 2NH3(g) Go

= -33.32 kJ/mol

17.8 Free Energy and Equilibrium

The equilibrium point in terms of kinetics occurs when the forward and reverse reactions were occurring

at an equal rate. In terms of free energy, the equilibrium point occurs at the lowest value of

free energy available to the reaction system.

These two definitions are the same!

o G = Gproducts - Greactants = 0

o Ex. Given for the reaction Hg(l)

Hg(g) that Ho

= 61.3 kJ/mol and So

= 100.0 J/mol.K, calculate the normal

boiling point of Hg.

We can utilize the previous two equations (G = Gproducts - Greactants = 0 and G = Go + RT ln (Q) to

form an equation that describes the relationship between free energy and the value of the equilibrium

constant.

Go

= -RT ln(K) (the “rat link” equation); R = 8.3145 J/mol K

o When Go

= 0, free energy of reactants and products are equal when all components are in

their standard states. During a phase change, G = 0.

o When Go

< 0, Go

products < Go

reactants The reaction is not at equilibrium, K > 1 since

pressure of products is > 1 and the pressure of reactants is < 1.

o When Go

> 0, Go reactants < G

o products The reaction is not at equilibrium, K < 1 since

pressure of products is < 1 and the pressure of reactants is > 1.

o Ex. Calculate the approximate standard free energy for the ionization of hydrofluoric acid, HF (Ka

= 1.0 x 10-3

), at 25°C.

We can use G =Go

+ RT ln(Q) to calculate the direction that a reaction will shift to reach equilibrium.

o Free energy is energy available to do useful work. Wmax = G

Ho

(kJ/mol) So

(J/mol.K)

PbO (s) -218.0 70.0

SO2 (s) -297.0 248.0

PbS (s) -100.0 91.0

O2 (g) ? 205.0

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combustion

f

comb comb

AP Chemistry: Thermochemistry Worksheet #1

#1

When 1.000 gram of propane gas, C3H8, is burned at 25oC and 1.00

atmosphere, H2O(l) and CO2(g) are formed with the evolution of

50.29 kilojoules.

a) Write a balanced equation for the combustion reaction.

b) Calculate the molar enthalpy of combustion, Ho , of propane.

c) Calculate the standard molar enthalpy of formation, Ho

, of propane gas.

d) Calculate the entropy change, So , for the reaction and account for the sign of S

o .

2. From each of the following pairs of substances, choose the one expected to have the greater absolute entropy.

Explain your choice in each case. Assume 1 mole of each substance.

1) He(g) at 1 atmosphere or He(g) at 0.05 atmosphere, both at the same temperature

2) Mg(s) at 0oC or Mg(s) at 150

oC, both at the same pressure

3. For the reaction shown, Ho

= +22.1 kilocalories per mole at 25oC. PCl5(g) PCl3(g) + Cl2(g)

a) Does the tendency of reactions to proceed to a state of minimum energy favor the formation of products of this

reaction? Explain.

Compound Ho

f (kJ/mol) So

(J/K.mol)

H2O (l) -286 70

CO2(g) -393.5 214.5

O2(g) 0.0 205

C3H8(g) ? 270

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comb

f

b) Does the tendency of reactions to proceed to a state of maximum entropy favor the formation of the products of

this reaction? Explain.

c) State whether an increase in temperature drives this reaction to the right, to the left, or has no effect. Explain.

d) State whether a decrease in the volume of the system at constant temperature drives this reaction to the right, to

the left or has no effect. Explain.

4. CO(g) + 2H2(g) CH3OH(l) Ho

= -128.1 kJ The data in the table shown were determined at 25°C.

a) Calculate Go

for the reaction above at 25°C.

b) Calculate Keq for the reaction above at 25°C.

c) Calculate So

for the reaction above at 25°C.

o o o

d) In the table above, there are no data for H2. What are the values of Hf , Gf , and of the absolute entropy, S ,

for H2 at 25°C?

5. The enthalpy change for the combustion

of butyric acid at 25°C, Ho , is -2183.5

kJ/mole. The combustion reaction is:

C3H7COOH(l) + 5O2(g) 4CO2(g) + 4H2O(l)

a) From the data above, calculate the

standard heat of formation, Ho

, for butyric acid at 25°C.

Hfo(kJ/mol) Gf

o(kJ/mol) S

o(J/K

.mol)

CO(g) -110.5 -137.3 +197.9

CH3OH(l) -238.6 -166.2 +126.8

Substance Standard Heat of Formation,

Hfo, in kJ/mol

Absolute Entropy,

S°, in J/K.mol

C(s) 0.00 5.69

CO2(g) -393.5 213.6

H2(g) 0.00 130.6

H2O(l) -285.85 69.96

O2(g) 0.00 205.0

C3H7COOH(l) ? 226.3

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f

f

b) Write a correctly balanced equation for the formation of butyric acid from its elements.

c) Calculate the standard entropy change, So

, for the formation of butyric acid at 25°C. The entropy change, S°,

for the combustion reaction above is -117.1 J/K at 25°C.

d) Calculate the standard free energy of formation, G o, for butyric acid at 25°C.

6.

a) Write a separate balanced chemical equation for

the combustion of each of the following: C(s),

H2(g), and C2H5OH(l). Consider the only products

to be CO2(g) and/or H2O(l).

b) In principle, ethanol can be prepared by the following reaction. 2C(s) + 2H2(g) + H2O(l) C2H5OH(l)

Calculate the standard enthalpy change, H°, for the preparation of ethanol, as shown in the reaction above.

c) Calculate the standard entropy change, S°, for the reaction given in part (b).

d) Calculate the value of the equilibrium constant at 25°C for the reaction represented by the equation in part (b).

Substance Enthalpy of Combustion,

H°, (kJ/mol)

Absolute Entropy,

S°, (J/mol.K)

C(s) -393.5 5.740

H2(g) -285.8 130.6

C2H5OH(l) -1366.7 160.7

H2O(l) ? 69.91

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AP Chemistry: Thermochemistry Worksheet #2

1. The reaction shown is a reversible reaction. BaCl2(g) + NH3(g) Cl3BNH3(s)

a. Predict the sign of the entropy change, S, as the reaction proceeds to the right. Explain your prediction.

b. If the reaction spontaneously proceeds to the right, predict the sign of the enthalpy change, H. Explain

your prediction.

c. The direction in which the reaction spontaneously precedes changes as the temperature is increased above a

specific temperature. Explain.

d. What is the value of the equilibrium constant at the temperature referred to in part c; that is, the specific

temperature at which the direction of the spontaneous reaction changes? Explain.

2. ClF3 can be prepared by the reaction represented by this equation: Cl2(g) + 3 F2(g) 2 ClF3(g) For ClF3 the o o

standard enthalpy of formation, Hf , is -163.2 kJ/mol and the standard free energy of formation, G f, is -123.0 kJ/mol.

a. Calculate the value of the equilibrium constant for the reaction at 298K.

b. Calculate the standard entropy change, So, for the reaction at 298K.

c. If ClF3 were produced as a liquid rather than as a gas, how would the sign and the magnitude of S for the

reaction be affected? Explain.

d. At 298K the absolute entropies of Cl2(g) and ClF3(g) are 222.96 J/mol.K and 281.50 J/mol

.K, respectively.

i. Account for the larger entropy of ClF3(g) relative to that of Cl2(g).

ii. Calculate the value of the absolute entropy of F2(g) at 298K.

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3. The tables below contain information for determining thermodynamic properties of the reaction shown.

C2H5Cl(g) + Cl2(g) C2H4Cl2(g) + HCl

a. Calculate the H° for the reaction above, using the table of average bond dissociation energies.

b. Calculate the S° for the reaction at 298 K, using data from either table as needed.

c. Calculate the value of Keq for the reaction at 298K.

d. What is the effect of an increase in temperature on the value of the equilibrium constant? Explain your

answer.

4. A student performs an experiment to determine the molar

enthalpy of solution of urea, H2NCONH2. The student places

91.95 g of water at 25°C into a coffee cup calorimeter and

immerses a thermometer in the water. After 50 s, the student

adds 5.13 g of solid urea, also at 25°C, to the water and

measures the temperature of the solution as the urea dissolves.

A plot of the temperature data is shown to the right.

a. Determine the change in temperature of the solution

that results from the dissolution of the urea.

b. According to the data, is the dissolution of urea in water an endothermic process or an exothermic process?

Substance Standard Free Energies of Formation

at 298K, Gof, at 298K (kJ/mol)

C2H4Cl2(g) -80.3

C2H5Cl(g) -60.5

HCl(g) -95.3

Cl2(g) 0

Bond Average Bond Dissociation

Energies at 298K, (kJ/mol)

C-H 414

C-C 347

C-Cl 377

Cl-Cl 243

H-Cl 431

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Justify your answer.

c. Assume that the specific heat capacity of the calorimeter is negligible and that the specific heat capacity of

the solution of urea and water is 4.2 J g−1

°C−1

throughout the experiment.

i. Calculate the heat of dissolution of the urea in joules.

ii. Calculate the molar enthalpy of solution, ΔHsoln, of urea in kJ mol−1

.

d. Using the information in the table shown, calculate the value of the

molar entropy of solution, ΔSsoln, of urea at 298 K. Include units with

your answer.

e. The student repeats the experiment and this time obtains a result for of urea that is 11 percent below the

accepted value. Calculate the value of that the student obtained in this second trial.

f. The student performs a third trial of the experiment but this time adds urea that has been taken directly from

a refrigerator at 5°C. What effect, if any, would using the cold urea instead of urea at 25°C have on the

experimentally obtained value of ΔHsoln? Justify your answer.

5. At 25°C the equilibrium constant, Kp, for the reaction shown is 0.281 atmosphere. Br2(l)Br2(g)

a. What is the G°298 for this reaction?

b. It takes 193 joules to vaporize 1.00 gram of Br2(l) at 25°C and 1.00 atmosphere pressure. What are the

values of H°298 and S°298 for this reaction?

c. Calculate the normal boiling point of bromine. Assume that H° and S° remain constant as the

temperature is changed.

Accepted Value

ΔHsoln of urea 14.0 kJ mol-1

ΔGsoln of urea -6.9 kJ mol-1

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f f

d. What is the equilibrium vapor pressure of bromine at 25°C?

6. Propane, C3H8, is a hydrocarbon that is commonly used as fuel for cooking.

a. Write a balanced equation for the complete combustion of propane gas, which yields CO2(g) and H2O(l).

b. Calculate the volume of air at 30°C and 1.00 atmosphere that is needed to burn completely 10.0 grams of

propane. Assume that air is 21.0 percent O2 by volume.

c. The heat of combustion of propane is -2,220.1 kJ/mol. Calculate the heat of formation, Hf

o, of propane

given that H o of H2O(l) = -285.3 kJ/mol and H

o of ClO2(g)= -393.5 kJ/mol.

d. Assuming that all of the heat evolved in burning 30.0 grams of propane is transferred to 8.00 kilograms of

water (specific heat = 4.18 J/g°C), calculate the increase in temperature of the water.

Page 15: AP Chemistry: Chapters 6 & 17 Notes Thermochemistrymhs-ap-chemistry-2014-2015.weebly.com/.../01_thermo_student_note… · AP Chemistry: Chapters 6 & 17 Notes Thermochemistry ... Simple

Extra Practice for the Extra Crazy Student

1. The information given refers to the following reaction: C2H2(g) + 2 H2(g) C2H6(g)

a. If the value of the standard entropy change, ΔS°, for the reaction is -232.7 joules per mole Kelvin, calculate

the standard molar entropy, S°, of C2H6 gas.

b. Calculate the value of the standard free-energy change, ΔG°, for the reaction. What does the sign of ΔG°

indicate about the reaction above?

c. Calculate the value of the equilibrium constant, K, for the reaction at 298 K.

d. Calculate the value of the C [triple bond] C bond energy in C2H2 in kilojoules per mole.

2. At 298 K, the standard enthalpy change, ΔH°, for the reaction represented below is -145 kilojoules.

2 H2S(g) + SO2(g) <===> 3 S(s) + 2 H2O(g)

a. Predict the sign of the standard entropy change, ΔS°, for the reaction. Explain the basis for your prediction.

b. At 298 K, the forward reaction (i.e., toward the right) is spontaneous. What change, if any, would occur in

the value of ΔG° for this reaction as the temperature is increased? Explain your reasoning using

thermodynamic principles.

c. What change, if any, would occur in the value of the equilibrium constant, Keq, for the situation described in

(b)? Explain your reasoning. d. The absolute temperature at which the forward reaction becomes nonspontaneous can be predicted. Write

the equation that is used to make the prediction. Why does this equation predict only an approximate value

for the temperature?

3. The reaction represented below is spontaneous at 25°C. Assume that all reactants and products are in their

standard states. C4H10(g) + 13 O2(g) 8 CO2(g) + 10 H2O(l)

a. Predict the sign of ΔS° for the reaction and justify your prediction.

b. What is the sign of ΔG° for the reaction? How would the sign and magnitude of ΔG° be affected by an

increase in temperature to 50 °C? Explain your answer.

c. What must be the sign of ΔH° for the reaction at 25°C? How does the total bond energy of the reactants

compare to that of the products?

d. When the reactants are placed together in a container, no change is observed even though the reaction is

known to be spontaneous. Explain this observation.

Substance S° (J/mol.K) ΔH°f (kJ/mol)

C2H2(g) 200.9 226.7

H2(g) 130.7 0.00

C2H6(g) ? -84.7

Bond Bond Energy (kJ/mol)

C-C 347

C=C 611

C-H 414

H-H 436

Page 16: AP Chemistry: Chapters 6 & 17 Notes Thermochemistrymhs-ap-chemistry-2014-2015.weebly.com/.../01_thermo_student_note… · AP Chemistry: Chapters 6 & 17 Notes Thermochemistry ... Simple