AP Chemistry
AP Chemistry 7: Thermodynamics Name __________________________
A. Enthalpy (H): Bond Energy (5.3 to 5.5, 8.8)
1. chemical reactions typically involve breaking bonds between reactant atoms and forming new bonds
2. breaking bonds takes energy ∴ chemical system gains bond energy; surroundings lose energy (heat, etc.)
3. forming bonds releases energy ∴ chemical system loses energy, surroundings gain energy
4. change in energy called “change in enthalpy”—ΔH
a. when energy required to break bonds > energy released to form new bonds, +ΔH (endothermic)
1. products at a higher energy state than reactants (weaker bonds)
2. surroundings lose energy (cool down)
b. when energy required to break bonds < energy released to form new bonds, –ΔH (exothermic)
1. products at a lower energy state than reactants (stronger bonds)
2. surroundings gain energy (heat up)
5. thermochemical equation
a. chemical equation with ΔH
1. listed to the right of equation
2. included as reactant (endothermic) or product (exothermic)
b. ΔH can be used in dimensional analysis process
6. ΔH from calorimetry
a. reactants are put in an insulated container filled with water, where heat is exchanged between reactants and water, but no heat is lost
b. by conservation of energy: ΔHreaction = –Qwater
1. Q = mcΔT for simple coffee cup calorimeter—aqueous reactions
a. m = mass of water
b. c = specific heat of water (4.18 J/g•K)
c. ΔT = change in temperature (Tf – Ti)
temperature can stay in oC, since
1 oC = 1 K (don't add 273 to ΔToC!)
2. Q = (C + mc)ΔT for “bomb" calorimeter
a. C = “bomb constant” accounts for all non-water components that change temperature
b. all other letters are the same as the simple calorimeter
7. ΔH using bond energy (B.E.) data
|Bond Energies in (kJ/mol) |
|Single |Multiple |
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B. Entropy (S): Disorder (19.2)
1. atoms/molecules have inherent disorder depending on
a. number of atoms—more internal motion = disorder
b. spacing of molecules—farther apart = disorder
c. speed of molecules—faster = disorder
2. predict increase in disorder for physical changes (+ΔS)
a. spread out: evaporation, diffusion and effusion
(solution: spread out solute and solvent (+ΔS), but bond solute-solvent (-ΔS) ∴ ?, but usually +ΔS)
b. motion: melting and boiling
3. predict increase in disorder for chemical changes (+ΔS): moles gaseous products > moles gaseous reactants
C. Thermodynamic Data (5.6 to 5.7, 19.4)
|Species |ΔHfo (kJ/mol) |So (kJ/mol•K) |
|Al |0.0 |+0.0283 |
|Al3+ |-531.0 |-0.3217 |
|Al2O3 |-1675.7 |+0.0509 |
1. standard heat of formation (ΔHfo) data
a. ΔHo for the formation of one mole of compound from its elements at standard temperature (25oC)
Al: Al(s) → Al(s) ∴ no reaction
Al3+: Al(s) → Al3+ + 3 e-
Al2O3: 2 Al(s) + 3/2 O2(g) → Al2O3(s)
b. ΔHfo for elements in natural state = 0.0
c. more negative = more stable (harder to decompose)
2. standard entropy (So) data
a. amount of disorder compared to H+ (simplest form of matter), which is zero by definition
b. listed in J/mol•K on AP exam, so you will have to convert to kJ/mol•K for most calculations
3. calculations using the thermodynamic data chart
a. altering ΔHfo
1. opposite sign for the reverse reaction
C + 2 Cl2 → CCl4 = –139.4 kJ
∴ CCl4 → C + 2 Cl2 = +139.4 kJ
2. multiply by number of moles (coefficient)
1 mole CCl4= –139.4 kJ
∴ 2 mole CCl4 = –278.8 kJ
b. calculate ΔH for a reaction using ΔHfo
1. ΔH ≈ ΔHo = ΔHfoproducts – ΔHforeactants
2. Hess’s Law: ΔH for a multi-step reaction equals the sum of ΔH for each step
| CH4(g) → C + 2 H2 |-(-74.8) |
| C + O2 → CO2(g) |-393.5 |
|+ 2 H2 + O2 → 2 H2O(g) |2(-241.8) |
|CH4(g) + 2 O2 → CO2(g) + 2 H2O(g) -802.3 |
c. calculate ΔS for a reaction using So
ΔS ≈ ΔSo = Soproducts – Soreactants
D. Gibbs Free Energy (G): Overall Energy State (19.5 to 19.6)
1. combination of enthalpy and entropy: G = H + TS
2. for a chemical or physical change: ΔGo = ΔHo – ToΔSo
a. To = 298 K
b. where T ≠ 298 K: ΔG ≈ ΔHo – TΔSo
3. determining if a process is spontaneous (ΔG < 0)
a. lower potential energy (-ΔH)—chemical reactions
b. greater disorder (+ΔS)—physical changes
c. depends on temperature
1. threshold temperature (Tthreshold)
2. occurs when ΔG = 0 ∴ Tthreshold = ΔHo/ΔSo
d. summary chart
|ΔH |ΔS |Spontaneous Process (ΔG 0, ΔS < 0 or ΔS ≈ 0.
| |> 0 |≈ 0 |< 0 |
|Melting ice at 0oC | | | |
|CH4(g) + 2 O2(g) → CO2(g) + 2 H2O(l) | | | |
|CH4(g) + 2 O2(g) → CO2(g) + 2 H2O(g) | | | |
|Distilling alcohol-water mixture | | | |
Thermodynamic Data
9. 2 Na2O2(s) + 4 HCl(g) → 4 NaCl(s) + 2 H2O(l) + O2(g)
Determine ΔH from the thermochemical reactions below.
2 Na2O2(s) + 2 H2O(l) → 4 NaOH(s) + O2(g) ΔH1 = -126 kJ
NaOH(s) + HCl(g) → NaCl(s) + H2O(l) ΔH2 = -179 kJ
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10. C2H2(g) + 5 N2O(g) → 2 CO2(g) + H2O(g) + 5 N2(g)
Determine ΔH from the thermochemical equations below.
2 C2H2(g) + 5 O2(g) → 4 CO2(g) + 2 H2O(g) ΔH1 = -2512 kJ
N2(g) + ½ O2(g) → N2O(g) ΔH2 = 104 kJ
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11. NO(g) + O(g) → NO2(g)
Determine ΔH for the above reaction using the following thermochemical equations.
NO(g) + O3(g) → NO2(g) + O2(g) ΔH1 = -198.9 kJ
O3(g) → 3/2 O2(g) ΔH2 = -142.3 kJ
O2(g) → 2 O(g) ΔH3 = 495.0 kJ
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12. a. Write the equation for the combustion of methanol, CH3OH. (all reactants and products are gaseous).
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b. Calculate ΔH using ΔHfo values.
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1.00 g of methanol is burned in a bomb calorimeter that contains 1200 g of water. The temperature increases 3.4 K.
c. Calculate the heat generated by the combustion reaction.
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d. Calculate the calorimeter constant of the bomb.
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13. Ca(s) + SO3(g) + 2 H2O(l) → CaSO3•2 H2O(s)
ΔH = -795 kJ and ΔS = -0.2535 kJ/K for the reaction.
a. Calculate ΔHfo for CaSO3•2 H2O.
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b. Calculate So for CaSO3•2 H2O.
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Gibbs Free Energy
14. Consider the reaction at 25oC:
Cu(s) + 4 H+(aq) + 2 NO3-(aq) → Cu2+(aq) + 2 NO2(g) + 2 H2O(l).
a. Calculate ΔHo using ΔHfo values.
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b. Calculate ΔSo using So values.
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c. Calculate ΔGo.
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15. NH4NO3(s) → NH4+(aq) + NO3-(aq)
Determine the following for the above reaction.
a. Is the reaction exothermic or endothermic?
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b. Is there an increase or decrease in entropy?
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c. Is the reaction spontaneous at 25oC?
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16. 2 SO2(g) + O2(g) → 2 SO3(g)
a. Calculate ΔHo.
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b. Calculate ΔSo.
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c. Calculate ΔG at 400 K.
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d. Determine the temperature range where the reaction is spontaneous.
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17. C2H5OH(l) + 3 O2(g) → 2 CO2(g) + 3 H2O(l)
a. Calculate ΔHo.
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b. Calculate ΔSo.
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c. Calculate ΔG at 20oC.
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d. At which temperature (if any) will the reaction be spontaneous?
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18. When H2SO4(l) is dissolved in water, the temperature of the mixture increases. Predict the sign of ΔH, ΔS and ΔG for this process (justify your answer).
| |+/– |Justification |
|ΔH | | |
|ΔS | | |
|ΔG | | |
19. C2H5OH(l) → C2H5OH(g)
Calculate the boiling point (threshold temperature) given the information: ΔH = 37.95 kJ and ΔS = 0.1078 kJ/K.
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Heat of Reaction Lab
20. Use calorimetry to determine ΔH for a series of reactions, compare the results with thermodynamic data, and combine the results to verify Hess' law.
Heat about 75 mL of water to about 70oC. Place a Styrofoam cup in a 250-mL beaker. Add 50.0 mL cold tap water to the cup. Record the temperature, TC. Measure out 50.0 mL of the hot water and place in a second Styrofoam cup. Record the temperature, TH. Pour the hot water into the cold water, cover the cup, insert the thermometer in the hole, and mix gently. Record the temperature every 20 seconds for 3 minutes. Discard the water.
a. (1) Record the temperatures.
|TC | |TH | |
|time|20 |40 |60 |80 |100 |120 |
|(s) | | | | | | |
| | |Time (s) |
(3) Use the y-intercept to determine Tmix.
|Tmix (y-intercept) | |
b. Calculate the calorimeter constant C using data above.
(1) Calculate the average of the hot and cold temperatures.
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(2) Calculate the bomb constant
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Place a Styrofoam cup in a 250 mL beaker. Add 50.0 mL of 3.00 M NaOH. Record the temperature, To. Pour 50.0 mL of 3.00 M HCl into the NaOH, cover, insert the thermometer, and mix gently. Record the temperature every 20 seconds for 3 minutes. Discard the mixture.
c. (1) Record the temperatures.
|To | |
|time|20 |40 |60 |80 |100 |120 |
|(s) | | | | | | |
| | |Time (s) |
(3) Use the y-intercept to determine Tmix.
|Tmix (y-intercept) | |
d. Calculate ΔHreaction per mole of reactant based on the calorimetry data.
|ΔT (K) | |
|Qwater (kJ) | |
|ΔHreaction/mole | |
e. Calculate ΔHreaction per mole of reactant based on ΔHfo.
|OH-(aq) + H+(aq) → H2O(l) |
|ΔH | |
|% | |
Place a Styrofoam cup in a 250 mL beaker. Add 50.0 mL of 3.00 M NH4Cl. Record the temperature, To. Pour 50.0 mL of 3.00 M NaOH into the NH4Cl, cover, insert the thermometer, and mix gently. Record the temperature every 20 seconds for 3 minutes. Discard the mixture.
f. (1) Record the temperatures.
|To | |
|time|20 |40 |60 |80 |100 |120 |
|(s) | | | | | | |
| | |Time (s) |
(3) Use the y-intercept to determine Tmix.
|Tmix (y-intercept) | |
g. Calculate ΔHreaction per mole of reactant based on the calorimetry data.
|ΔT (K) | |
|Qwater (kJ) | |
|ΔHreaction/mole | |
h. Calculate ΔHreaction per mole of reactant based on ΔHfo.
|NH4+(aq) + OH-(aq) → NH3(aq) + H2O(l) |
|ΔH | |
|% | |
Place a Styrofoam cup in a 250 mL beaker. Add 50.0 mL of 3.00 M NH3. Record the temperature, To. Pour 50.0 mL of 3.00 M HCl into the NH3, cover, insert the thermometer, and mix gently. Record the temperature every 20 seconds for 3 minutes. Discard the mixture.
i. (1) Record the temperatures.
|To | |
|time|20 |40 |60 |80 |100 |120 |
|(s) | | | | | | |
| | |Time (s) |
(3) Use the y-intercept to determine Tmix.
|Tmix (y-intercept) | |
j. Calculate ΔHreaction per mole of reactant based on the calorimetry data.
|ΔT (K) | |
|Qwater (kJ) | |
|ΔHreaction/mole | |
k. Calculate ΔHreaction per mole of reactant based on ΔHfo.
|NH3(aq) + H+(aq) → NH4+(aq) |
|ΔH | |
|% | |
l. Show that the chemical equations and ΔH from
part (d) – part (g) = part (j).
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Practice Quiz
Multiple Choice (no calculator)
Briefly explain why the answer is correct in the space provided.
|1 |2 |3 |4 |
|Bond Energy (kJ/mole) |150 |240 |210 |
(A) - 870 (B) - 390 (C) +180 (D) + 450
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2. C2H4(g) + 3 O2(g) → 2 CO2(g) + 2 H2O(g)
For the reaction, ΔH is -1,300 kJ. What is the value of ΔH, in kJ, if the combustion produced liquid water rather than water vapor? (ΔH for H2O(l) → H2O(g) is 45 kJ/mol)
(A) -1,300 (B) -1,210 (C) -1,345 (D) -1,390
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3. CH4 (g) + 2 O2(g) → CO2(g) + 2 H2O(l) ΔHo = -900 kJ
What is the standard heat of formation of CH4, in kJ/mol, as calculated from the data below?
(ΔHfoH2O = -300 kJ/mol, ΔHfoCO2 = -400 kJ/mol)
(A) -200 (B) -100 (C) 100 (D) 200
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4. H2(g) + ½ O2(g) → H2O(l) ΔHo = x
2 Na(s) + ½ O2(g) → Na2O(s) ΔHo = y
Na(s) + ½ O2(g) + ½ H2(g) → NaOH(s) ΔHo = z
What is ΔH for the reaction below?
Na2O(s) + H2O(l) → 2 NaOH(s)
(A) x + y + z (B) x + y – z
(C) x + y - 2z (D) 2z - x - y
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5. Which is true when ice melts at its normal melting point?
(A) ΔH < 0, ΔS > 0, ΔG = 0
(B) ΔH < 0, ΔS < 0, ΔG > 0
(C) ΔH > 0, ΔS < 0, ΔG < 0
(D) ΔH > 0, ΔS > 0, ΔG = 0
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6. Which of the following reactions has the largest positive value of ΔS per mole of Cl2?
(A) H2(g) + Cl2(g) → 2 HCl(g)
(B) Cl2(g) + O2(g) → Cl2O(g)
(C) Mg(s) + Cl2(g) → MgCl2(s)
(D) 2 NH4Cl(s) → 4 H2(g) + Cl2(g)
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7. Ice is added to hot water in an insulated container, which is then sealed. What has happened to the total energy and the total entropy when the system reaches equilibrium?
(A) Energy and entropy remain constant
(B) Energy remains constant, entropy decreases
(C) Energy remains constant, entropy increases
(D) Energy decreases, entropy increases
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8. N2(g) + 3 H2(g) → 2 NH3(g)
The above reaction is thermodynamically spontaneous at 298 K, but becomes nonspontaneous at higher temperatures. Which of the following is true at 298 K?
(A) ΔG, ΔH, and ΔS are all positive.
(B) ΔG, ΔH, and ΔS are all negative.
(C) ΔG and ΔH are negative, but ΔS is positive.
(D) ΔG and ΔS are negative, but ΔH is positive.
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9. 3 C2H2(g) → C6H6(g)
What is the standard enthalpy change, ΔHo, for the reaction represented above?
(ΔHfoC2H2 is 230 kJ•mol-1; ΔHfoC6H6 is 80 kJ•mol-1)
(A) -610 kJ (B) 150 kJ (C) -770 kJ (D) 610 kJ
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10. When solutions of NH4SCN and Ba(OH)2 are mixed 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?
(A) –ΔG –ΔH –ΔS (B) –ΔG +ΔH –ΔS
(C) –ΔG +ΔH +ΔS (D) +ΔG –ΔH +ΔS
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11. X(s) Δ X(l)
Which of the following is true for any substance undergoing the process represented above at its normal melting point?
(A) ΔS < 0 (B) ΔH = 0
(C) ΔH = TΔG (D) ΔH = TΔS
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12. For a reaction, ΔHo = -150 kg/mol and ΔSo = -50 J/mol•K. Which statement is true about this reaction?
(A) It is spontaneous at high temperature only.
(B) It is spontaneous at low temperature only.
(C) It is spontaneous at all temperatures.
(D) It is non-spontaneous at all temperatures.
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Free Response (calculator)
1. Consider the combustion of butanoic acid at 25oC:
HC4H7CO2(l) + 5 O2(g) → 4 CO2(g) + 4 H2O(l)
ΔHo= -2,183.5 kJ
|Substance |ΔHfo (kJ/mol) |So (kJ/mol•K) |
|CO2(g) |-393.5 |0.2136 |
|H2O(l) |-285.8 |0.0699 |
|O2(g) |0.0 |0.2050 |
|C3H7COOH(l) |? |0.2263 |
a. Calculate ΔHfo, for butanoic acid.
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b. Calculate ΔSo for the combustion reaction at 25oC.
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c. Calculate ΔGo for the combustion reaction at 25oC.
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d. What is the spontaneous temperature range?
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2. Consider the synthesis reaction:
N2(g) + 3 F2(g) → 2 NF3(g)
(ΔHo298 = -264 kJ mol-1, ΔSo298 = -278 J K-1 mol-1)
a. Calculate ΔGo298 for the reaction.
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b. For what temperature range is the reaction spontaneous?
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c. Calculate the heat released when 0.256 mol of NF3(g) is formed from N2(g) and F2(g) at 1.00 atm and 298 K.
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d. Calculate the F–F bond energy using the information above and the bond energies
(N≡N = 946 kJ/mol, N–F = 272 kJ/mol).
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3. The combustion of carbon monoxide is represented by the equation: CO(g) + ½O2(g) → CO2(g)
a. Determine ΔHo for the reaction above using the values.
C(s) + ½ O2(g) → CO(g) ΔHo298 = -110.5 kJ•mol-1
C(s) + O2(g) → CO2(g) ΔHo298 = -393.5 kJ•mol-1
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b. Determine ΔSo for the reaction above using the table
|Substance |CO(g) |CO2(g) |O2(g) |
|So (J/mol•K) |197.7 |213.7 |205.1 |
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c. Determine ΔGo for the above reaction at 298 K.
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4. The dissolving of AgNO3(s) in water is represented by the equation: AgNO3(s) → Ag+(aq) + NO3-(aq)
a. Is ΔG positive, negative, or zero? Justify your answer.
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b. Is ΔS positive, negative, or zero? Justify your answer.
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c. The solution cools when AgNO3(s) is dissolved. Is ΔH for the dissolving positive, negative or zero? Justify your answer.
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