Ap chem 프린스턴 화학 summary part 6

Problem 1:

Consider the reaction: 2 H₂(g) + O₂(g) → 2 H₂O(l).

Given that the enthalpy change for this reaction is -571.6 kJ/mol, calculate the enthalpy change for the formation of 1 mole of H₂O(g) under standard conditions. Assume the enthalpy of vaporization of water is 44.0 kJ/mol.

Options:

• (A) -571.6 kJ/mol
• (B) -285.8 kJ/mol
• (C) -241.8 kJ/mol
• (D) -227.6 kJ/mol

Ammonia (NH₃) is formed through the reaction: N₂(g) + 3 H₂(g) → 2 NH₃(g).

If the bond dissociation energies are as follows: N≡N = 946 kJ/mol, H–H = 436 kJ/mol, and N–H = 391 kJ/mol, what is the enthalpy change (ΔH) for the formation of 1 mole of NH₃?

Options:

• (A) -92 kJ/mol
• (B) -45 kJ/mol
• (C) -38 kJ/mol
• (D) -46 kJ/mol

Problem 3:

In which of the following reactions is entropy change (ΔS) positive?

Options:

• (A) H₂(g) + I₂(s) → 2 HI(g)
• (B) CO₂(g) + H₂O(l) → H₂CO₃(aq)
• (C) 2 NO₂(g) → N₂O₄(g)
• (D) C(s) + O₂(g) → CO₂(g)

Problem 4:

Using standard enthalpies of formation, calculate the enthalpy change for the combustion of methane (CH₄) in oxygen to form CO₂(g) and H₂O(l): CH₄(g) + 2 O₂(g) → CO₂(g) + 2 H₂O(l).

Given: ΔH_f° [CH₄(g)] = -75 kJ/mol, ΔH_f° [CO₂(g)] = -393.5 kJ/mol, ΔH_f° [H₂O(l)] = -285.8 kJ/mol.

Options:

• (A) -890 kJ/mol
• (B) -572 kJ/mol
• (C) -350 kJ/mol
• (D) -802 kJ/mol

Problem 5:

Given the bond dissociation energies: C–H = 412 kJ/mol, O=O = 498 kJ/mol, C=O = 799 kJ/mol, and O–H = 463 kJ/mol, calculate the enthalpy change for the combustion of 1 mole of methane (CH₄): CH₄(g) + 2 O₂(g) → CO₂(g) + 2 H₂O(g).

Options:

• (A) -680 kJ/mol
• (B) -810 kJ/mol
• (C) -1640 kJ/mol
• (D) -1200 kJ/mol

Problem 6:

The combustion of ethanol (C₂H₅OH) can be represented by the equation: C₂H₅OH(l) + 3 O₂(g) → 2 CO₂(g) + 3 H₂O(l).

Given: ΔH_f° [C₂H₅OH(l)] = -277.0 kJ/mol, ΔH_f° [CO₂(g)] = -393.5 kJ/mol, ΔH_f° [H₂O(l)] = -285.8 kJ/mol, calculate the enthalpy change for the combustion of 1 mole of ethanol.

Options:

• (A) -1367 kJ/mol
• (B) -1350 kJ/mol
• (C) -1270 kJ/mol
• (D) -1420 kJ/mol

Problem 7:

The hydrogenation of ethene (C₂H₄) can be represented by the equation: C₂H₄(g) + H₂(g) → C₂H₆(g).

If the bond dissociation energies are as follows: C=C = 612 kJ/mol, C–H = 412 kJ/mol, C–C = 348 kJ/mol, and H–H = 436 kJ/mol, calculate the enthalpy change (ΔH) for the reaction.

Options:

• (A) -135 kJ/mol
• (B) -130 kJ/mol
• (C) -120 kJ/mol
• (D) -100 kJ/mol

Problem 8:

Calculate the enthalpy change for the combustion of propane (C₃H₈) using the following average bond enthalpies: C–H = 412 kJ/mol, C–C = 348 kJ/mol, O=O = 498 kJ/mol, C=O = 799 kJ/mol, O–H = 463 kJ/mol.

The balanced equation for the combustion of propane is: C₃H₈(g) + 5 O₂(g) → 3 CO₂(g) + 4 H₂O(g).

Options:

• (A) -2100 kJ/mol
• (B) -2400 kJ/mol
• (C) -2016 kJ/mol
• (D) -2320 kJ/mol

Problem 9:

The formation of sulfur dioxide (SO₂) can be described by the reaction: S(s) + O₂(g) → SO₂(g).

If the standard enthalpies of formation are ΔH_f° [SO₂(g)] = -296.8 kJ/mol, calculate the enthalpy change for the combustion of 2 moles of sulfur.

Options:

• (A) -296.8 kJ/mol
• (B) -593.6 kJ/mol
• (C) -148.4 kJ/mol
• (D) -198.2 kJ/mol

Problem 10:

The decomposition of hydrogen peroxide (H₂O₂) is represented by the equation: 2 H₂O₂(l) → 2 H₂O(l) + O₂(g).

Given the standard enthalpies of formation: ΔH_f° [H₂O₂(l)] = -187.8 kJ/mol, ΔH_f° [H₂O(l)] = -285.8 kJ/mol, calculate the enthalpy change for the decomposition of 2 moles of H₂O₂.

Options:

• (A) -196.0 kJ
• (B) +196.0 kJ
• (C) +98.0 kJ
• (D) -98.0 kJ

Problem 11:

Using the following bond dissociation energies: C–H = 412 kJ/mol, O=O = 498 kJ/mol, C=O (in CO₂) = 799 kJ/mol, and O–H = 463 kJ/mol, calculate the enthalpy change for the combustion of methanol (CH₃OH) in the reaction: 2 CH₃OH(l) + 3 O₂(g) → 2 CO₂(g) + 4 H₂O(g).

Options:

• (A) -726 kJ/mol
• (B) -1452 kJ/mol
• (C) -1500 kJ/mol
• (D) -1800 kJ/mol

Problem 12:

Calculate the enthalpy change for the formation of 1 mole of sulfur trioxide (SO₃) from sulfur dioxide (SO₂) and oxygen: 2 SO₂(g) + O₂(g) → 2 SO₃(g).

Given: ΔH_f° [SO₂(g)] = -296.8 kJ/mol, ΔH_f° [SO₃(g)] = -395.7 kJ/mol.

Options:

• (A) -197.8 kJ/mol
• (B) -198.9 kJ/mol
• (C) -98.9 kJ/mol
• (D) -99.5 kJ/mol

Problem 13:

The standard enthalpy changes for the combustion of carbon and carbon monoxide are shown below:

C(s) + O₂(g) → CO₂(g)      ΔH^⦵ = –394 kJ/mol

CO(g) + 1/2 O₂(g) → CO₂(g)      ΔH^⦵ = –283 kJ/mol

What is the standard enthalpy change, in kJ, for the following reaction?

C(s) + 1/2 O₂(g) → CO(g)

Options:

• (A) -677 kJ/mol
• (B) -111 kJ/mol
• (C) +111 kJ/mol
• (D) +677 kJ/mol

Problem 14:

Consider the equations below:

CH₄(g) + O₂(g) → HCHO(l) + H₂O(l)      ΔH^⦵ = x

HCHO(l) + 1/2 O₂(g) → HCOOH(l)      ΔH^⦵ = y

2 HCOOH(l) → (COOH)₂(s) + H₂O(l)      ΔH^⦵ = z

What is the enthalpy change of the reaction below?

2 CH₄(g) + 3 O₂(g) → (COOH)₂(s) + 3 H₂O(l)

Options:

• (A) x + y + z
• (B) 2x + y + z
• (C) 2x + 2y + z
• (D) 2x + 2y + 2z

Problem 15:

Consider the two reactions involving iron and oxygen:

2 Fe(s) + O₂(g) → 2 FeO(s)      ΔH^⦵ = –544 kJ

4 Fe(s) + 3 O₂(g) → 2 Fe₂O₃(s)      ΔH^⦵ = –1648 kJ

What is the enthalpy change, in kJ, for the following reaction?

4 FeO(s) + O₂(g) → 2 Fe₂O₃(s)

Options:

• (A) -1648 – 2(−544)
• (B) -544 – (−1648)
• (C) -1648 - 544
• (D) -1648 – 2(544)

Problem 16:

Consider the following reactions.

Cu₂O(s) + 1/2 O₂(g) → 2 CuO(s)      ΔH^⦵ = –144 kJ

Cu₂O(s) → Cu(s) + CuO(s)      ΔH^⦵ = +11 kJ

What is the value of ΔH^⦵, in kJ, for the following reaction?

Cu(s) + 1/2 O₂(g) → CuO(s)

Options:

• (A) -144 + 11
• (B) +144 - 11
• (C) -144 - 11
• (D) +144 + 11

Problem 17:

Consider the following reactions.

N₂(g) + O₂(g) → 2 NO(g)      ΔH^⦵ = +180 kJ

2 NO₂(g) → 2 NO(g) + O₂(g)      ΔH^⦵ = +112 kJ

What is the ΔH^⦵ value, in kJ, for the following reaction?

N₂(g) + 2 O₂(g) → 2 NO₂(g)

Options:

• (A) -1 × (+180) - 1 × (+112)
• (B) -1 × (+180) + 1 × (+112)
• (C) 1 × (+180) - 1 × (+112)
• (D) 1 × (+180) + 1 × (+112)

Problem 18:

Consider the following enthalpy of combustion data.

C(s) + O₂(g) → CO₂(g)      ΔH^⦵ = –x kJ/mol

H₂(g) + 1/2 O₂(g) → H₂O(l)      ΔH^⦵ = –y kJ/mol

C₂H₆(g) + 7/2 O₂(g) → 2 CO₂(g) + 3 H₂O(l)      ΔH^⦵ = –z kJ/mol

What is the enthalpy of formation of ethane in kJ/mol?

2 C(s) + 3 H₂(g) → C₂H₆(g)

Options:

• (A) [(-x) + (-y)] - (-z)
• (B) (-z) - [(-x) + (-y)]
• (C) [(-2x) + (-3y)] - (-z)
• (D) (-z) - [(-2x) + (-3y)]

Problem 19:

Two 100 cm³ aqueous solutions, one containing 0.010 mol NaOH and the other 0.010 mol HCl, are at the same temperature.

When the two solutions are mixed, the temperature rises by y °C.

Assume the density of the final solution is 1.00 g cm⁻³.

Specific heat capacity of water = 4.18 J g⁻¹ K⁻¹.

What is the enthalpy change of neutralization in kJ mol⁻¹?

Options:

• (A) (200 × 4.18 × y) / (1000 × 0.020)
• (B) (200 × 4.18 × y) / (1000 × 0.010)
• (C) (100 × 4.18 × y) / (1000 × 0.010)
• (D) (200 × 4.18 × (y + 273)) / (1000 × 0.010)

Problem 20:

Which expression gives the mass, in g, of ethanol required to produce 683.5 kJ of heat upon complete combustion?

(M_r for ethanol = 46.0, ΔH^⦵_c = −1367 kJ mol⁻¹)

Options:

• (A) 683.5 / (1367 × 46.0)
• (B) 1367 / (683.5 × 46.0)
• (C) (683.5 × 46.0) / 1367
• (D) (1367 × 46.0) / 683.5

Problem 21:

5.35 g of solid ammonium chloride, NH₄Cl(s), was added to water to form 25.0 g of solution. The maximum decrease in temperature was 14 K. What is the enthalpy change, in kJ mol⁻¹, for this reaction? (Molar mass of NH₄Cl = 53.5 g mol⁻¹; the specific heat capacity of the solution is 4.18 J g⁻¹ K⁻¹)

Options:

• (A) (25.0 × 4.18 × (14 + 273)) / (0.1 × 1000)
• (B) - (25.0 × 4.18 × 14) / (0.1 × 1000)
• (C) + (25.0 × 4.18 × 14) / (0.1 × 1000)
• (D) (25.0 × 4.18 × 14) / 1000

Problem 22:

When 25.0 cm³ of 0.100 mol dm⁻³ NaOH(aq) is mixed with 25.0 cm³ of 0.100 mol dm⁻³ HCl(aq) at the same temperature, a temperature rise, ΔT, is recorded. What is the expression, in kJ mol⁻¹, for the enthalpy of neutralization? (Assume the density of the mixture = 1.00 g cm⁻³ and its specific heat capacity = 4.18 J g⁻¹ K⁻¹)

Options:

• (A) (25.0 × 4.18 × ΔT) / (50.0 × 0.100)
• (B) (25.0 × 4.18 × ΔT) / (25.0 × 0.100)
• (C) (50.0 × 4.18 × ΔT) / (50.0 × 0.100)
• (D) (50.0 × 4.18 × ΔT) / (25.0 × 0.100)

Problem 23:

When 100 cm³ of 1.0 mol dm⁻³ HCl is mixed with 100 cm³ of 1.0 mol dm⁻³ NaOH, the temperature of the resulting solution increases by 5.0 °C. What will be the temperature change, in °C, when 50 cm³ of these two solutions are mixed?

Options:

• (A) 2.5
• (B) 5.0
• (C) 10
• (D) 20

Problem 24:

When 50.0 cm³ of 1.0 mol dm⁻³ NaOH(aq) is mixed with 50.0 cm³ of 1.0 mol dm⁻³ HCl(aq), the temperature increases by 6.5 °C. Calculate the enthalpy change of neutralization in kJ mol⁻¹. (Assume the density of the solution is 1.00 g cm⁻³ and the specific heat capacity is 4.18 J g⁻¹ K⁻¹)

Options:

• (A) -54.3 kJ mol⁻¹
• (B) -27.1 kJ mol⁻¹
• (C) -104.5 kJ mol⁻¹
• (D) -26.0 kJ mol⁻¹

Problem 25:

When 75.0 cm³ of 1.5 mol dm⁻³ NaOH(aq) is mixed with 75.0 cm³ of 1.5 mol dm⁻³ HCl(aq) at the same temperature, the temperature rise is recorded as 7.2 °C. What is the enthalpy change of neutralization in kJ mol⁻¹? (Assume the density of the mixture = 1.00 g cm⁻³ and its specific heat capacity = 4.18 J g⁻¹ K⁻¹)

Options:

• (A) -40.1 kJ mol⁻¹
• (B) -75.4 kJ mol⁻¹
• (C) -90.3 kJ mol⁻¹
• (D) -45.2 kJ mol⁻¹

Problem 26:

50.0 cm³ of 1.0 mol dm⁻³ H₂SO₄ is mixed with 50.0 cm³ of 1.0 mol dm⁻³ NaOH. The temperature of the solution increases by 5.5 °C. Assume the density of the solution is 1.00 g cm⁻³ and its specific heat capacity is 4.18 J g⁻¹ K⁻¹. Calculate the enthalpy change of neutralization in kJ mol⁻¹.

Options:

• (A) -46.0 kJ mol⁻¹
• (B) -27.5 kJ mol⁻¹
• (C) -55.5 kJ mol⁻¹
• (D) -110.0 kJ mol⁻¹

Problem 27:

The combustion of ethanol, C₂H₅OH, releases 1367 kJ of heat per mole of ethanol burned. How much heat will be released if 46.0 g of ethanol is completely combusted?

Options:

• (A) 1367 kJ
• (B) 2734 kJ
• (C) 683.5 kJ
• (D) 91.0 kJ

Problem 28:

When 50.0 cm³ of 2.0 mol dm⁻³ NaOH(aq) is mixed with 50.0 cm³ of 2.0 mol dm⁻³ HCl(aq), the temperature of the mixture rises by 12.0 °C. Assume the density of the mixture is 1.00 g cm⁻³ and its specific heat capacity is 4.18 J g⁻¹ K⁻¹. What is the enthalpy change of neutralization in kJ mol⁻¹?

Options:

• (A) -50.2 kJ mol⁻¹
• (B) -100.5 kJ mol⁻¹
• (C) -25.1 kJ mol⁻¹
• (D) -75.4 kJ mol⁻¹

Problem 29:

When 0.025 moles of an unknown compound is combusted, it releases 1250 kJ of heat. What is the molar enthalpy change of combustion of the compound in kJ mol⁻¹?

Options:

• (A) -50.0 kJ mol⁻¹
• (B) -500.0 kJ mol⁻¹
• (C) -50,000 kJ mol⁻¹
• (D) -1250.0 kJ mol⁻¹

Problem 30:

100 cm³ of 0.5 mol dm⁻³ NaOH(aq) is mixed with 100 cm³ of 0.5 mol dm⁻³ HCl(aq). The temperature increases by 3.5 °C. Assume the density of the mixture is 1.00 g cm⁻³ and its specific heat capacity is 4.18 J g⁻¹ K⁻¹. Calculate the enthalpy change of neutralization in kJ mol⁻¹.

Options:

• (A) -14.6 kJ mol⁻¹
• (B) -58.5 kJ mol⁻¹
• (C) -7.3 kJ mol⁻¹
• (D) -35.5 kJ mol⁻¹

Problem 31:

For a certain chemical reaction, the enthalpy change (ΔH) is -120 kJ/mol, and the entropy change (ΔS) is -50 J/mol·K. Calculate the Gibbs free energy change (ΔG) at 298 K.

Options:

• (A) -105 kJ/mol
• (B) -135 kJ/mol
• (C) +105 kJ/mol
• (D) +135 kJ/mol

Problem 32:

A reaction has an enthalpy change (ΔH) of +80 kJ/mol and an entropy change (ΔS) of +200 J/mol·K. At what temperature (in K) will the reaction become spontaneous?

Options:

• (A) 200 K
• (B) 400 K
• (C) 600 K
• (D) 800 K

Problem 33:

The reaction below has an enthalpy change (ΔH) of -40 kJ/mol and an entropy change (ΔS) of -100 J/mol·K. Determine if this reaction is spontaneous at 350 K.

Options:

• (A) Yes, because ΔG is negative.
• (B) No, because ΔG is positive.
• (C) It is spontaneous only at lower temperatures.
• (D) It is spontaneous only at higher temperatures.

Problem 34:

A reaction has an enthalpy change (ΔH) of -85 kJ/mol and an entropy change (ΔS) of -200 J/mol·K. At what temperature (in K) does this reaction become non-spontaneous?

Options:

• (A) 200 K
• (B) 425 K
• (C) 575 K
• (D) 850 K

Problem 35:

The decomposition of a compound has an enthalpy change (ΔH) of +150 kJ/mol and an entropy change (ΔS) of +300 J/mol·K. Calculate the Gibbs free energy change (ΔG) at 350 K. Will the reaction be spontaneous?

Options:

• (A) ΔG = +45 kJ/mol, Non-spontaneous
• (B) ΔG = -45 kJ/mol, Spontaneous
• (C) ΔG = +45 kJ/mol, Spontaneous
• (D) ΔG = -45 kJ/mol, Non-spontaneous

Problem 36:

For the reaction: N₂(g) + 3 H₂(g) → 2 NH₃(g), the enthalpy change (ΔH) is -92 kJ/mol and the entropy change (ΔS) is -198 J/mol·K. Determine the Gibbs free energy change (ΔG) at 500 K.

Options:

• (A) -188 kJ/mol
• (B) -2 kJ/mol
• (C) -178 kJ/mol
• (D) +7 kJ/mol