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AP Matter Class Packet Unit 5

AP Thermochemistry Unit 3 Packet

Electrochemistry Review:

Thermodynamics Review

Students should be able to demonstrate an understanding of the following essential knowledge:

· 3.C.2 Net changes in energy for a chemical reaction can be endothermic or exothermic.

· 5.A.1 Temperature is a measure of the average kinetic energy of atoms and molecules.

· 5.A.2 The process of kinetic energy transfer at the particulate scale is referred to in this course as heat transfer, and the spontaneous direction of the transfer from a hot to a cold body.

· 5.B.2 When two systems are ion contact with each other and are otherwise isolated, the energy that comes out of one system is equal to the energy that goes into the other system. The combined energy of the two systems remains fixed. Energy transfer can occur through either heat exchange or work.

· 5.B.3 Chemical systems undergo three main processes that change their energy: heating/cooling, phase transitions, and chemical reactions.

· 5.B.4 Calorimetry is an experimental technique that is used to determine the heat exchanged/transferred in a chemical system.

· 5.C.1 Potential energy is associated with a particular geometric arrangement of atoms or ions and the electrostatic interactions between them.

· 5.D.1 Potential energy is associated with the interaction of molecules; as molecules draw near each other, they experience an attractive force.

Answer the following questions using the heat formula. Show work with units and significant figures.

1. How many Joules of energy are needed to change the temperature of 100.0 grams of water from 20.0C to 40.0C?

2. How many kilojoules of energy are needed to change the temperature of 15.0 grams of water from 35.0C to 75.0C?

3. If the temperature of water is changed from 10.0C to 35.0C by the addition of 350.0J, how many grams were heated?

4. If the temperature of water is changed from 100.0C to 250.0C by the addition of 5000.0J, how many grams were heated?

5. If 3500.0J of energy are applied to 150.0 grams of water at 50.0C, what is the final temperature?

6. Look at the rearranged equation for heat, solved for specific heat. What are the units of specific heat based on this rearranged equation?

7. What is the specific heat of silver if an 80.0 gram sample is heated from 24.0C to 49.0C by adding 468.2J?

8. What is the change in temperature when 3.00 grams of Iron (specific heat = 0.45J/gC) is subjected to 350.0.J of energy?

9. What mass of Aluminum (specific heat = 0.902 J/gC) can be heated from 25.0C to 90.0C with the addition of 100.0J of heat?

10. How many joules of heat must be released in order to change the temperature of 50.0 grams of air (specific heat 1.01 J/gC) from 35.0C to 25.0C?

Base answers to question 11 on diagram:

11. How much heat is added to change the

substance from the coldest to the warmest

pure liquid state?

AP Calorimetric Calculations


· 5.B.4 Calorimetry is an experimental technique that is used to determine the heat exchanged/transferred in a chemical system.

The first law of thermodynamics: Energy can never be created nor destroyed. Therefore, the energy of the universe is constant.

1. It takes 585J of energy to raise the temperature of 125.6g of Hg from 20.0 to 53.5C. Calculate the specific heat of Hg.

2. A 46.2g sample of Cu is heated to 95.4C and placed in a calorimeter containing 75.0g of water at 19.6C. The final temperature inside the calorimeter equals 21.8C. Calculate the specific heat of copper.

3. The specific heat of aluminum is 0.9000J/gC and the density is 2.71g/cm3.

a. Calculate the energy needed to raise the temperature of 4.36x105cm3 block from 22.8 to 94.6C.

b. Calculate the molar heat capacity of aluminum.

4. A piece of iron with a mass of 56.0g and specific heat of 0.45J/gC is placed in 155g of water at 21.0C. The final temperature is 33.5C. Calculate the original temperature of the iron.

Enthalpy



· 5.C.2 The net change during a reaction is the sum of the energy required to break the bonds in the reactant molecules and the energy released in forming the bond of the product molecules. The net change in energy may be positive for endothermic reactions where energy is required or negative for exothermic reactions where energy is released.

1. Consider the reaction:2Mg + O2 2MgOH=-1204kJ

a. Is this reaction endothermic or exothermic?

b. Calculate the heat transferred when 3.60g of Mg reacts with excess oxygen.

c. How many grams of MgO are produced during the enthalpy change of -96.0kJ?

2. Consider the reaction: 2AgBr + Cl2 2AgCl + Br2H=-55.2kJ

a. Calculate the heat transferred by 37.56g of silver bromide reacting with excess chlorine.

b. How many grams of liquid bromine are produced when the enthalpy change equals -106kJ?

c. Calculate the heat of reaction when 60.0g of solid silver chloride are produced.

3. Given the following information: C2H6 + 7/2 O2 2CO2 + 3H2O

a. What is the enthalpy change for the reverse reaction?

b. Which reaction (forward or reverse) is more favorable?

Enthalpy of Formation




1. The thermite reaction is highly exothermic and is used for welding :2Al(s) + Fe2O3 2Fe(s) + Al2O3

Calculate the heat of this reaction using enthalpies of formation.

2. Calculate the enthalpy of these reactions using enthalpies of formation:

a. 4FeO(s) + O2(g) 2Fe2O3(s)

b. SiCl4(l) + 2H2O(l) SiO2(s) + 4HCl(aq)

c. NH3(g) + HCl(g) NH4Cl(s)

d. MgO(s) + H2O(l) Mg(OH) 2 (s)

e. NH3(g) + HCl(g) NH4Cl(s)

f. C2H5OH(l) + 3O2 (g) 2CO2 (g) + 3H2O(g)

Reminder of an old Enthalpy Question:

3. Calculate the change in energy that accompanies the following reaction given the data below.

H2(g) + F2(g) → 2 HF(g) Bond Type Bond Energy H−H 432 kJ/mol F−F 154 kJ/mol H−F 565 kJ/mol

Hess’s Law



· 5.C.1 Potential energy is associated with a particular geometric arrangement of atoms or ions and the electrostatic interactions between them.


1. Given the following data, calculate the heat of S(s) + O2(g) SO2(g)

S(s) + 3/2 O2(g) SO3(g)H= -395.2kJ

2SO2(g) + O2(g) 2SO3(g)H= -198.2kJ

2. Given the following data, calculate the heat of C6H4(OH)2(aq) + H2O2 (aq) C6H4O2 (aq) + 2H2O(l)

C6H4(OH)2(aq) C6H4O2 (aq) + H2(g)H= 177.4kJ

H2 (g) + O2 (g) H2O2 (aq)H= -191.2kJ

H2 (g) + ½ O2 (g) H2O(g)H= -241.8kJ

H2O(g) H2O(l)H= -43.8kJ

3. Given the following data, calculate the heat of2CO2 (g) + H2O(g) C2H2 (g) + 5/2 O2 (g)

C2H6(g) C2H2 (g) + 2H2 (g)H= 850.5kJ

H2 (g) + ½ O2 (g) H2O(g)H= -641.2kJ

C2H6(g) + 7/2 O2 (g) 2CO2 (g) + 3H2O(g)H= -2547kJ

4. Given the following data, calculate the heat of NO(g) + O(g) NO2(g)

2O3 (g) 3O2(g)H= -427kJ

O2 (g) 2O(g)H= 495kJ

NO(g) + O3(g) NO2 (g) + O2 (g)H= -199kJ

Thermodynamics Practice

1. Pentane, C5H12, is a hydrocarbon used in the production of Styrofoam and is present in certain fuels.

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

(b)Calculate the volume of air at 25C and 1.00 atmosphere that is needed to burn completely 50.5 grams of pentane. Assume that air is 21.0 percent O2 by volume.

(c)The heat of combustion of pentane is -3,285.3 kJ/mol. Calculate the heat of formation, ΔHf, of pentane given that ΔHf of H2O(l) = -285.3 kJ/mol and ΔHf of CO2(g) = -393.5 kJ/mol.

(d)Assuming that all of the heat evolved in burning 50.5 grams of pentane is transferred to 10.0 kilograms of water (specific heat = 4.18 J/g.K), calculate the increase in temperature of water.

2. (a) The specific heat of fluorine gas is 0.037 J/gK. Calculate the molar heat capacity (in J/molK) of fluorine gas.

(b) The molar heat capacity of a compound with the formula C4H10SO is 43.6 J/molK. Calculate the specific heat, c, of this substance.

3.Given the following data:S(s) + 3/2 O2(g) SO3(g)∆Hfº =-395.2 kJ

2 SO2(g) + O2 2 SO3(g)∆Hfº=-198.2 kJ

Calculate ∆Hºrxn for the reaction:S(s) + O2(g) SO2(g)

Work


· 5.B.2 When two systems are ion contact with each other and are otherwise isolated, the energy that comes out of one system is equal to the energy that goes into the other system. The combined energy of the two systems remains fixed. Energy transfer can occur through either heat exchange or work.

Energy is often defined as the “ability to do work”. ∆E = q (heat) + w (work)

Algebraic signs of q:

· +q if heat absorbed

· –q if heat released

Algebraic sign of w:

· + w if work done on the system (i.e., compression)

· −w if work done by the system (i.e., expansion)

· When related to gases, work is a function of pressure. Pressure is defined as force per unit of area, so when the volume is changed work was either done on the gas or by the gas. work = −P∆V

1. Calculate ∆E for a system undergoing an endothermic process in which 15.6 kJ of heat flows and where 1.4 kJ of work is done on the system.

2. Calculate the work associated with the expansion of a gas from 46 L to 64 L at a constant external pressure of 15 atm.

3. A balloon is being inflated to its full extent by heating the air inside it. In the final stages of this process, the volume of the balloon changes from 4.00 × 106 L to 4.50 × 106 L by the addition of 1.3 × 108 J of energy as heat. Assuming that the balloon expands against a constant pressure of 1.0 atm, calculate ∆E for the process. (To convert between L ⋅ atm and J, use 1 L ⋅ atm = 101.3 J.)

Entropy


· 5.E.1 Entropy is a measure of the dispersal of matter and energy.

Entropy is the degree of disorder or randomness in a substance. The symbol for change in entropy is ΔS.

Solids are very ordered and have low entropy. Liquids and aqueous ions have more entropy because they move about more freely, and gases have an even larger amount of entropy. According to the Second Law of Thermodynamics, nature is always proceeding to a state of higher entropy.

Determine whether the following reactions show an increase or decrease in entropy.

1. 2KClO3(s) → 2KCl(s) + 3O2(g)________________________

2. H2O(l) → H2O(s)________________________

3. N2(g) + 3H2(g) → 2NH3(g)________________________

4. 2NaCl(s) → 2Na(s) + Cl2 (g)________________________

5. KCl(s) → KCl(l)________________________

6. CO2(s) → CO2(g)________________________

7. H+(aq) + C2H3O2-(aq) → HC2H3O3(l)________________________

8. C(s) + O2(g) → CO2(g)________________________

9. H2(g) + Cl2(g) → 2HCl(g)________________________

10. Ag+(aq) + Cl-(aq) → AgCl(s)________________________

11. 2N2O5(g) → 4NO2(g) + O2(g)________________________

12. 2Al(s) + 3I2(s) → 2AlI3(s)________________________

13. H+(aq) + OH-(aq) → H2O(l) ( str.acid-base rxn)________________________

14. 2NO(g) → N2(g) + O2(g)________________________

15. H2O(g) → H2O(l)________________________

Entropy Calculations


· 5.E.1 Entropy is a measure of the dispersal of matter and energy.

· 5.E.2 Some physical or chemical processes involve both a decrease in the internal energy of the components (ΔH<0) under consideration and an increase in the entropy of those components (ΔS>0). These processes are necessarily “thermodynamically favored” (ΔG<0).

The second law of thermodynamics: the universe is constantly increasing the dispersal of matter and energy.

The third law of thermodynamics: the entropy of a perfect crystal at 0 K is zero. [Not a lot of perfect crystals out there so, entropy values are RARELY ever zero—even elements] So what? This means the absolute entropy of a substance can then be determined at any temp. higher than zero K. (Handy to know if you ever need to defend why G & H for elements = 0. . . . BUT S does not!)

ΔS is + when dispersal/disorder increases (favored)

ΔS is – when dispersal/disorder decreases

NOTE: Units are usually J/(molrxn • K) (not kJ!)

1. Calculate the entropy change at 25°C, in J/(molrxn • K) for: 2 SO2(g) + O2(g) → 2 SO3(g) Given the

following data: SO2(g) 248.1 J/(mol• K) O2(g) 205.3 J/(mol• K) SO3(g) 256.6 J/(mol • K)

Gibbs Free Energy Calculations

2. 2 H2O + O2(g) → 2 H2O2 Calculate the free energy of formation for the oxidation of water to produce

hydrogen peroxide given the following information ∆G values: H2O ‒56.7 kcal/molrxn O2(g) 0 kcal/molrxn

H2O2 ‒27.2 kcal/molrxn

3. Cdiamond(s) + O2(g) → CO2(g) ∆G°= ‒397 kJ

Cgraphite(s) + O2(g) → CO2(g) ∆G°= ‒394 kJ

Calculate ∆G° for the reaction Cdiamond(s)→Cgraphite(s)

4. 2 SO2(g) + O2(g) → 2 SO3(g) The reaction above was carried out at 25°C and 1 atm. Calculate ∆H°, ∆S°, and

∆G° using the following data:

5. The overall reaction for the corrosion (rusting) of iron by oxygen is 4 Fe(s) + 3 O2(g) → 2 Fe2O3(s) Using the following data, calculate the equilibrium constant for this reaction at 25°C

Gibbs Free Energy

For a physical or chemical reaction to be spontaneous, the sign of ΔG (Gibbs Free Energy) must be negative. The mathematical formula for this value is:


· 5.E.2 Some physical or chemical processes involve both a decrease in the internal energy of the components (ΔH<0) under consideration and an increase in the entropy of those components (ΔS>0). These processes are necessarily “thermodynamically favored” (ΔG<0).

· 5.E.3 If a chemical or physical process is not driven by both entropy and enthalpy changes, then the Gibbs free energy change can be used to determine whether the process is thermodynamically favored.

· 5.E.4 External sources of energy can be used to drive change in cases where the Gibbs free energy change is positive.

· 5.E.5 A thermodynamically favored process may not occur due to kinetic constraints (kinetic vs. thermodynamic control).

ΔG = ΔH – TΔS

where ΔH = change in enthalpy or heat of reaction

ΔS = change in entropy or randomness

T = temperature in Kelvin

ΔH

ΔS

ΔG

-

+

+

-

-

-

+

+

Complete the table for the sign of ΔG; +, - or undetermined. When conditions allow for an undetermined sign of ΔG, temperature will

decide spontaneity. (temp. dependent)

1. The conditions in which ΔG is always negative is when ΔH is _________________ and ΔS is _________________.

2. The conditions in which ΔG is always positive is when ΔH is _________________ and ΔS is _________________.

3. When the situation is indeterminate, a low temperature favors the ( entropy / enthalpy ) factor, and a high temperature favors the ( entropy / enthalpy ) factor.

Answer Problems 4-6 with always, sometimes or never.

4. The reaction: 2KClO3(s) + heat → 2KCl(s) + 3O2g) will _________________ be spontaneous.

5. The reaction: 2H2(g) + O2(g) → 2H2O(l) + heat will _________________ be spontaneous.

6. The reaction: heat + H2O(s) → H2O(l) will _________________ be spontaneous.

7. What is the value of ΔG if ΔH = -32.0 kJ/mol, ΔS = +25 J/molK and T = 293 K? _________________

8. Is the reaction in Problem 7 spontaneous? ______ Explain ______________

9. What is the value of ΔG if ΔH = +12.0 kJ/mol, ΔS = - 5 J/molK and T = 290 K? _________________

10. Is the reaction is Problem 9 spontaneous? _________________

AP Problems

1. For the reaction above, H = +22.1 kilocalories per mole at 25C

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

(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.

2. Standard Heat ofAbsolute

Formation, Hf,Entropy, S,

Substancein kJ mol-1 in J mol-1 K-1

C(s)0.005.69

CO2(g)-393.5213.6

H2(g)0.00130.6

H2O(l)-285.8569.91

O2(g)0.00205.0

C3H7COOH(l)?226.3

The enthalpy change for the combustion of butyric acid at 25C, Hcomb, is -2,183.5 kilojoules per mole. The combustion reaction is

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

(a)From the above data, calculate the standard heat of formation, Hf, for butyric acid.

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

(c)Calculate the standard entropy change, Sf, for the formation of butyric acid at 25C. The entropy change, S, for the combustion reaction above is -117.1 J K-1 at 25C.

(d)Calculate the standard free energy of formation, Gf, for butyric acid at 25C.

3. When crystals of barium hydroxide, Ba(OH)2.8H2O, are mixed with crystals of ammonium thiocyanate, NH4SCN, at room temperature in an open beaker, the mixture liquefies, the temperature drops dramatically, and the odor of ammonia is detected. The reaction that occurs is the following:

Ba(OH)2.8H2O(s) + 2 NH4SCN(s) Ba2+ + 2 SCN- + 2 NH3(g) + 10 H2O(l)

(a)Indicate how the enthalpy, the entropy, and the free energy of this system change as the reaction occurs. Explain your predictions.

(b)If the beaker in which the reaction is taking place is put on a block of wet wood, the water on the wood immediately freezes and the beaker adheres to the wood. Yet the water inside the beaker, formed as the reaction proceeds, does not freeze even though the temperature of the reaction mixture drops to -15C. Explain these observations.

4. Br2(l) Br2(g)

At 25C the equilibrium constant, Kp, for the reaction above is 0.281 atmosphere.

(a)What is the G298 for this reaction?

(b)It takes 193 joules to vaporize 1.00 gram of Br2(l) at 25C and 1.00 atmosphere pressure. What are the values of H298 and S298 for this reaction?

(c)Calculate the normal boiling point of bromine. Assume that H and S remain constant as the temperature is changed.

(d)What is the equilibrium vapor pressure of bromine at 25C?

5. BCl3(g) + NH3(g) Cl3BNH3(s)

The reaction represented above is a reversible reaction.

(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 proceeds 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 (c); that is, the specific temperature at which the direction of the spontaneous reaction changes? Explain.

6. Cl2(g) + 3 F2(g) 2 ClF3(g)

ClF3 can be prepared by the reaction represented by the equation above. For ClF3 the standard enthalpy of formation, Hf, is -163.2 kilojoules/mole and the standard free energy of formation, Gf, is -123.0 kilojoules/mole.

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

(b)Calculate the standard entropy change, S, 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 joules per mole-Kelvin and 281.50 joules per mole-Kelvin, 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.

7. 2 C4H10(g) + 13 O2(g) 8 CO2(g) + 10 H2O(l)

The reaction represented above is spontaneous at 25C. Assume that all reactants and products are in their standard state.

(a)Predict the sign of S for the reaction and justify your prediction.

(b)What is the sign of G for the reaction? How would the sign and magnitude of G be affected by an increase in temperature to 50C? Explain your answer.

(c)What must be the sign of H for the reaction at 25C? How does the total bond energy of the reactants compare to that of the products?

(d)When the reactants are place together in a container, no change is observed even though the reaction is known to be spontaneous. Explain this observation.

8. 2H2S(g) + SO2(g) 3 S(s) + 2 H2O(g)

At 298 K, the standard enthalpy change, H for the reaction represented above is -145 kilojoules.

(a)Predict the sign of the standard entropy change, S, 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 G 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?

9. 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 30C 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, of propane given that Hf of H2O(l) = -285.3 kJ/mol and Hf of CO2(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.K), calculate the increase in temperature of water.

10. Lead iodide is a dense, golden yellow, slightly soluble solid. At 25C, lead iodide dissolves in water forming a system represented by the following equation.

PbI2(s) Pb2+ + 2 I-H = +46.5 kilojoules

(a)How does the entropy of the system PbI2(s) + H2O(l) change as PbI2(s) dissolves in water at 25C? Explain

(b)If the temperature of the system were lowered from 25C to 15C, what would be the effect on the value of Ksp? Explain.

(c)If additional solid PbI2 were added to the system at equilibrium, what would be the effect on the concentration of I- in the solution? Explain.

(d)At equilibrium, G = 0. What is the initial effect on the value of G of adding a small amount of Pb(NO3)2 to the system at equilibrium? Explain.

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