Thermochemistry
Review
1. Stoichiometry
#g |
| ΔHrxn
| Mm | n
Exothermic Reaction (enthalpy is a product):
where n = moles in the reaction for the specific substance
H2 (g) + Cl2 (g)  2 HCl (g) + ΔH
Endothermic Reaction (enthalpy is a reactant): 2 HBr (g) + ΔH  H2 (g) + Br2 (g)
2. Hess’s Law
 Given a list of reactions
 Flip reaction = flip the sign of ΔH
 Multiply or Divide reaction coefficients = Multiply or Divide ΔH
 Add up all ΔH to get ΔHrxn
3. Standard Heats of Formation
 ΔHrxn = ΣH products – ΣH reactants
 ΔSrxn = ΣS products – ΣS reactants
 ΔGrxn = ΣG products – ΣG reactants
 (Don’t forget to multiply by # of moles)
4. Bond Enthalpies
 Given the energies to break bonds
 Draw out all molecules (including the number of molecules)
 (+) Energy for breaking bonds: Need energy to break bonds
 (-) Energy for forming bonds: Release energy when forming bonds
 Add up all energies of all bonds to get ΔHrxn
5. Entropy and Gibbs Free Energy
 ΔG = ΔH – TΔS
Symbol
Description
ΔH
Enthalpy
ΔS
Entropy
ΔG
Gibbs Free Energy
Negative Value (-)
Exothermic (release)
Less disorder
Spontaneous
Positive Value (+)
Endothermic (Absorb)
More disorder
Non-spontaneous
© 2011, Robert Ayton. All rights reserved.
www.mrayton.com
© 2011, Robert Ayton. All rights reserved.
www.mrayton.com
Thermochemistry Problem Set
Stoichiometry
1.
2.
ZnCO3 (s)  ZnO (s) + CO2 (g)
The decomposition of zinc carbonate, as shown in the reaction above, at constant pressure requires the addition of
71.5 kJ of heat per mole of ZnCO3.
a. Include the heat in the reaction above to show a balanced thermochemical equation for this reaction.
b. Draw an enthalpy diagram for the reaction.
CH4 (g) + 2 O2 (g)  CO2 (g) + 2 H2O (l)
ΔH = -891 kJ
Consider the reaction above:
a. Calculate the heat released when 8.0 g of methane is burned in excess oxygen gas.
b. Calculate the heat released when 10.0 L of methane gas is burned at 1.2 atm and 25°C in the presence of
excess oxygen gas.
c. Calculate the heat released when 16.0 g of methane and 32.0 g of oxygen gas are reacted?
3.
CH3OH (g)  CO (g) + 2 H2 (g)
ΔH = +90.6 kJ
Consider the reaction above:
a. Is heat absorbed or evolved in the course of this reaction?
b. Calculate the amount of heat transferred when 64.0 g of CH3OH is decomposed at constant pressure.
c. For a given sample of CH3OH, the enthalpy change on the reaction is 22.65 kJ. How many grams of
hydrogen gas are produced?
d. What is the value of ΔH for the reverse of the previous reaction? How many kilojoules of heat are
released when 42.0 g of CO (g) reacts completely with H2 (g) to form CH3OH (g) at constant pressure?
4.
H2 (g) + F2 (g)  2 HF (g)
Gaseous hydrogen and fluorine combine in the reaction above to form hydrogen fluoride with an enthalpy change
of -540 kJ. What is the value of the heat of formation of one mole of HF (g)?
(A) -1080 kJ/mol
(B) -540 kJ/mol
(C) -270 kJ/mol
(D) +270 kJ/mol
(E) +540 kJ/mol
5.
2 H2 (g) + O2 (g)  2 H2O (l)
ΔH° = - 572 kJ
How much heat is released when 8.0 g of hydrogen gas is reacted with excess oxygen gas?
(A) -286 kJ
(B) -572 kJ
(C) -1144 kJ
(D) -1716 kJ
(E) -2288 kJ
© 2011, Robert Ayton. All rights reserved.
www.mrayton.com
Hess’s Law
6. Calculate the enthalpy change for the reaction
P4O6 (s) + 2 O2 (g)  P4O10 (s)
Given the following enthalpies of reaction:
P4 (s) + 3 O2 (g)  P4O6 (s)
ΔH = -1640.1 kJ
P4 (s) + 5 O2 (g)  P4O10 (s) ΔH = -2940.1 kJ
7. From the enthalpies of reaction
2 H2 (g) + O2 (g)  2 H2O (g) ΔH = -483.6 kJ
3 O2 (g)  2 O3 (g)
ΔH = +284.6 kJ
Calculate the heat of the reaction: 3 H2 (g) + O3 (g)  3 H2O (g)
8. From the enthalpies of reaction
H2 (g) + F2 (g)  2 HF (g)
ΔH = -537 kJ
C (s) + 2 F2 (g)  CF4 (g)
ΔH = -680 kJ
2 C (s) + 2 H2 (g)  C2H4 (g) ΔH = +52.3 kJ
Calculate ΔH for the reaction of ethylene with F2:
C2H4 (g) + 6 F2 (g)  2 CF4 (g) + 4 HF (g)
9. Given the data
N2 (g) + O2 (g)  2 NO (g)
ΔH = +180.7 kJ
2 NO (g) + O2 (g)  2 NO2 (g)
ΔH = -113.1 kJ
2 N2O (g)  2 N2 (g) + O2 (g)
ΔH = -163.2 kJ
Use Hess’s Law to calculate ΔH for the reaction: N2O (g) + NO2 (g)  3 NO (g)
10. C (s) + O2 (g)  CO2 (g)
ΔH° = -390 kJ/mol
H2 (g) + ½ O2 (g)  H2O (l) ΔH° = -290 kJ/mol
2 C (s) + H2 (g)  C2H2 (g) ΔH° = + 230 kJ/mol
Based on the information given above, what is ΔH° for the following reaction?
C2H2 (g) + O2 (g)  2 CO2 (g) + H2O (l)
(A) -1300 kJ
(B) -1070 kJ
(C) -840 kJ
(D) -780 kJ
(E) -680 kJ
11. C (s) + 2 H2 (g)  CH4 (g)
ΔH° = x
C (s) + O2 (g)  CO2 (g)
ΔH° = y
H2 (g) + ½ O2 (g)  H2O (l) ΔH° = z
Based on the information given above, what is ΔH° for the following reaction?
CH4 (g) + 2 O2 (g)  CO2 (g) + 2 H2O (l)
(A) x + y + z
(B) x + y – z
(C) z + y – 2x
© 2011, Robert Ayton. All rights reserved.
www.mrayton.com
(D) 2z + y – x
(E) 2z + y – 2x
© 2011, Robert Ayton. All rights reserved.
www.mrayton.com
Enthalpies of Formation
Substance
CH4 (g)
CO2 (g)
CCl4 (l)
Fe2O3 (s)
FeCl3 (s)
ΔH°f (kJ/mol)
-74.8
-393.5
-139.3
-822.16
-400
Substance
H2O (g)
H2O (l)
HCl (g)
H2S (g)
KOH (s)
K2CO3 (s)
ΔH°f (kJ/mol)
-241.8
-285.8
-92.3
-20.17
-424.7
-1150.18
Substance
N2O4 (g)
S8 (g)
SO2 (g)
SO3 (g)
SiO2 (s)
SiCl4 (l)
ΔH°f (kJ/mol)
9.66
102.3
-296.9
-395.2
-910.9
-640.1
12. Using values from the Enthalpies of Formation found above, calculate the value of ΔH° for each of the following
reactions:
a. N2O4 (g) + 4 H2 (g)  N2 (g) + 4 H2O (g)
b. 2 KOH (s) + CO2 (g)  K2CO3 (s) + H2O (g)
c. SO2 (g) + 2 H2S (g) 
S8 (g) + 2 H2O (g)
d. Fe2O3 (s) + 6 HCl (g)  2 FeCl3 (s) + 3 H2O (g)
13. Using values from the Enthalpies of Formation found above, calculate the value of ΔH° for each of the following
reactions:
a. 2 SO2 (g) + O2 (g)  2 SO3 (g)
b. 2 H2 (g) + O2 (g)  2 H2O (g)
c. CH4 (g) + 4 Cl2 (g)  CCl4 (l) + 4 HCl (g)
d. SiCl4 (l) + 2 H2O (l)  SiO2 (s) + 4 HCl (g)
Mg (s) + O2 (g) + H2 (g)  Mg(OH)2 (s)
14. Calculate the standard enthalpy of formation of solid Mg(OH) (for the reaction above), given the following data:
2 Mg (s) + O2 (g)  2 MgO (s)
ΔH° = -1203.6 kJ
Mg(OH)2 (s)  MgO (s) + H2O (l)
ΔH° = +37.1 kJ
2 H2 (g) + O2 (g)  2 H2O (l)
ΔH° = -571.7 kJ
15.
3 C2H2 (g)  C6H6 (g)
What is the standard enthalpy change, ΔH°, for the reaction represented above?
(ΔH°f of C2H2 (g) is 230 kJ mol-1; ΔH°f of C6H6 (g) is 83 kJ mol-1)
(A) -607 kJ
(B) -147 kJ
(C) -19 kJ
(D) +19 kJ
(E) +773 kJ
© 2011, Robert Ayton. All rights reserved.
www.mrayton.com
Bond Enthalpies
16. Using the Bond Enthalpies table, draw out the molecules and calculate the ΔH for each of the following gas-phase
reactions:
a. C2H4 + HOOH  HOCH2CH2OH
b. C2H4 + HCN  CH3CH2CN
c. 2 NCl3  N2 + 3 Cl2
17. Using the Bond Enthalpies table, draw out the molecules and calculate the ΔH for each of the following gas-phase
reactions:
a. CHBr3 + Cl2  CClBr3 + HCl
b. HSCH2CH2SH + 2 HBr  BrCH2CH2Br + 2 H2S
c. NH2NH2 + Cl2  2 NH2Cl
18. Using the Bond Enthalpies table, draw out the molecules and calculate the ΔH for each of the following gas-phase
reactions:
a. 2 CH4 (g) + O2 (g)  2 CH3OH (g)
b. H2 (g) + Br2 (g)  2 HBr (g)
c. 2 H2O2 (g)  2 H2O (g) + O2 (g)
d. 3 ethene (g)  cyclohexane
© 2011, Robert Ayton. All rights reserved.
www.mrayton.com
19.
2 H2 (g) + O2 (g)  2 H2O (g)
Based on the following information given in the table below, what is ΔH° for the above reaction?
Bond
Average bond energy
(kJ/mol)
H-H
500
O=O
500
O-H
500
(A) -2000 kJ
(B) -1500 kJ
(C) -500 kJ
(D) +1000 kJ
(E) +2000 kJ
Entropy and Gibbs Free Energy
20. For each of the following pairs, indicate which substance possesses the larger standard entropy:
a. 1 mol of O2 (g) at 300°C, 0.01 atm or 1 mol of N2 (g) at 50°C, 0.01 atm
b. 1 mol of H2O (g) at 100°C, 1 atm or 1 mol of H2O (l) at 100°C, 1 atm
c. 0.5 mol of CO2 (g) at 298 K, 5-L volume or 0.5 mol C3H8 (g) at 298 K, 20-L volume
d. 100 g of Na2SO4 (s) at 30°C or 100 g Na2SO4 (aq) at 30°C
21. Predict the sign of entropy change of the system for each of the following reactions:
a. 2 SO2 (g) + O2 (g)  2 SO3 (g)
b. Ba(OH)2 (s)  BaO (s) + H2O (g)
c. CO (g) + 2 H2 (g)  CH3OH (l)
d. FeCl2 (s) + H2 (g)  Fe (s) + 2 HCl (g)
e. Molten Fe solidifies
f. LiCl (s) is formed from Li (s) and Cl2 (g)
g. A precipitate is formed through two aqueous solutions
Substance
Al (s)
AlCl3 (s)
Cl2 (g)
CH4 (g)
C2H6 (g)
H2 (g)
ΔH°f (kJ mol-1)
0
-705.6
0
-74.8
-84.68
0
ΔS° (J mol-1 K-1)
28.32
109.3
222.96
186.3
229.5
130.58
Substance
H2O (l)
HCl (g)
MgCl2 (s)
Mg(OH)2 (s)
NH3 (g)
N2H4 (g)
ΔH°f (kJ mol-1)
-285.8
-92.3
-641.6
-924.7
-46.2
95.4
ΔS° (J mol-1 K-1)
69.91
186.7
89.6
63.24
192.5
238.5
22. Calculate ΔH°, ΔS°, and ΔG° reaction values for the following reactions by using the thermochemical data chart
above at 298 K. Also, indicate whether the reaction is spontaneous or nonspontaneous.
a. N2H4 (g) + H2 (g)  2 NH3 (g)
b. 2 Al (s) + 3 Cl2 (g)  2 AlCl3 (s)
c. Mg(OH)2 (s) + 2 HCl (g)  MgCl2 (s) + 2 H2O (l)
d. 2 CH4 (g)  C2H6 (g) + H2 (g)
© 2011, Robert Ayton. All rights reserved.
www.mrayton.com
23. Which of the following processes are spontaneous, and which are nonspontaneous:
a. The melting of ice cubes at -5°C and 1 atm pressure.
b. The dissolution of sugar in a cup of hot coffee.
c. The reaction of nitrogen atoms to form N2 molecules at 25°C and 1 atm.
d. The formation of CH4 and O2 molecules from CO2 and H2O at room temperature and 1 atm of pressure.
24. A certain reaction has ΔH° = -19.5 kJ and ΔS° = + 42.7 J/K.
a. Is the reaction exothermic or endothermic?
b. Does the reaction lead to an increase or decrease in disorder?
c. Calculate ΔG° for the reaction at 298 K.
d. Is the reaction spontaneous at 298K under standard conditions?
25. A particular reaction is spontaneous at 450 K. The enthalpy change for the reaction is +34.5 kJ. What can you
conclude about the sign and magnitude of ΔS for the reaction?
26. For a particular reaction, ΔH = -32 kJ and ΔS = -98 J/K. Assume that ΔH and ΔS do not vary with temperature.
a. At what temperature will the reaction be at equilibrium?
b. If the temperature is increased from that in part (a), will the reaction occur?
27. A certain reaction is nonspontaneous -25°C. The entropy change for the reaction is 95 J/K. What can you
conclude about the sign and magnitude of ΔH?
28.
2 Al (s) + 3 Cl2 (g)  2 AlCl3 (s)
The reaction above is not spontaneous under standard conditions but becomes spontaneous as the temperature
decreases toward absolute zero. Which of the following is true at standard conditions?
(A) ΔS and ΔH are both negative.
(B) ΔS and ΔH are both positive.
(C) ΔS is negative, and ΔH is positive.
(D) ΔS is positive, and ΔH is negative.
(E) ΔS and ΔH are both equal to zero.
29. Which of the following reactions is entropy increasing?
(A) 2 SO2 (g) + O2 (g)  2 SO3 (g)
(B) CO (g) + H2O (g)  H2 (g) + CO2 (g)
(C) H2 (g) + Cl2 (g)  2 HCl (g)
(D) 2 NO2 (g)  2 NO (g) + O2 (g)
(E) 2 H2S (g) + 3 O2 (g)  2 H2O (g) + 2 SO2 (g)
30. When solid NH4SCN is mixed with solid Ba(OH)2 in a closed container, the temperature drops and a gas is
produced. Which of the following indicates the correct signs for ΔG, ΔH, and ΔS for the process?
ΔG
ΔH
ΔS
(A)
(B)
+
(C)
+
+
(D)
+
+
© 2011, Robert Ayton. All rights reserved.
www.mrayton.com
(E)
+
-
-
Free Response
31. Liquid water is formed according to the following equation below.
2 H2 (g) + O2 (g)  2 H2O (l)
(a) Calculate the standard enthalpy change, ΔH, for the reaction represented in the equation above. (The molar
enthalpy of formation, ΔHf, for H2O (l) is -285.8 kJ mol-1 at 298 K)
(b) Calculate the amount of heat, in kJ, that is released when 6.0 g of H2 (g) is burned in air.
(c) What is the algebraic sign for the entropy of the reaction represented in the equation above? Justify your
answer.
(d) Clearly justify why this reaction above is spontaneous using thermodynamic principles.
(e) Given the molar enthalpy of vaporization, ΔHvap, for H2O (l) is 44.0 kJ mol-1 at 298 K, what is the standard
enthalpy change, ΔH, for the reaction 2 H2 (g) + O2 (g)  2 H2O (g)? (Hint: develop a reaction for the heat
of vaporization of liquid water)
N2 (g) + 3 F2 (g)  2 NF3 (g)
ΔH = -264 kJ mol-1
ΔS = 278 J K-1 mol-1
32. The following questions relate to the reaction represented by the chemical equation and information above.
(a) Calculate the value of the standard free energy change, ΔG, for the reaction.
(b) Calculate the standard enthalpy change, ΔH, that occurs when a 0.250 mol sample of NF3 (g) is formed from
N2 (g) and F2 (g) at 1.00 atm and 298 K.
(c) How many bonds are broken for N2 and F2, respectively, as well as how many bonds are formed in NF3 (all
according to the chemical reaction above)?
(d) Use both the information above and the table below of average bond enthalpies to calculate the average
enthalpy of the N≡N bond.
Bond
N≡N
N-F
F-F
Average Bond Enthalpy (kJ mol-1)
?
272
141
CO (g) + ½ O2 (g)  CO2 (g)
33. The combustion of carbon monoxide to produce carbon dioxide is represented in the equation above.
(a) Determine the value of the standard enthalpy change, ΔH°rxn, using the information below.
C (s) + ½ O2 (g)  CO (g)
ΔH° = -110.5 kJ mol-1
C (s) + O2 (g)  CO2 (g)
ΔH° = -393.5 kJ mol-1
(b) Determine the value of the standard entropy change, ΔS°rxn, using the information below.
CO (g)
S° = 197.7 J K-1 mol-1
CO2 (g)
S° = 213.7 J K-1 mol-1
© 2011, Robert Ayton. All rights reserved.
www.mrayton.com
O2 (g)
S° = 205.1 J K-1 mol-1
(c) Determine the standard free energy change, ΔG°rxn, for the reaction at 298 K. Include units with your answer.
(d) Is the reaction spontaneous under these standard conditions? Justify your answer.
© 2011, Robert Ayton. All rights reserved.
www.mrayton.com