Topic 15: Energetics/thermochemistry

Which equation represents the lattice enthalpy of magnesium sulfide?

A. MgS (s) → Mg (g) + S (g)

B. MgS (s) → Mg⁺ (g) + S^– (g)

C. MgS (s) → Mg²⁺ (g) + S^2– (g)

D. MgS (s) → Mg (s) + S (s)

What is the enthalpy of solution of MgF₂(s) in kJ mol⁻¹?

Lattice enthalpy of MgF₂(s) = 2926 kJ mol⁻¹

Hydration enthalpy of Mg²⁺(g) = −1963 kJ mol⁻¹

Hydration enthalpy of F⁻(g) = −504 kJ mol⁻¹

A.     2926 − 1963 + 2(−504)

B.     2926 − 1963 − 504

C.     −2926 − (−1963) − (−504)

D.     −2926 − (−1963) − 2(−504)

Which value represents the lattice enthalpy, in kJ mol⁻¹, of strontium chloride, SrCl₂?

A.     – (–829) + 164 + 243 + 550 + 1064 – (–698)

B.     –829 + 164 + 243 + 550 + 1064 – 698

C.     – (–829) + 164 + 243 + 550 + 1064 – 698

D.     –829 + 164 + 243 + 550 + 1064 – (–698)

Which change is exothermic?

A.  $\frac{1}{2}$Cl₂ (g) → Cl (g)

B.   K (g) → K⁺ (g) + e⁻

C.   KCl (s) → K⁺ (g) + Cl⁻ (g)

D.   Cl (g) + e⁻ → Cl⁻ (g)

Calculate values for the following changes using section 8 of the data booklet.

ΔH_atomisation (Na) = 107 kJ mol⁻¹
ΔH_atomisation (O) = 249 kJ mol⁻¹

$\frac{1}{2}$O₂(g) → O^2- (g):

Na (s) → Na⁺ (g):

The standard enthalpy of formation of sodium oxide is −414 kJ mol⁻¹. Determine the lattice enthalpy of sodium oxide, in kJ mol⁻¹, using section 8 of the data booklet and your answers to (d)(i).

(If you did not get answers to (d)(i), use +850 kJ mol⁻¹ and +600 kJ mol⁻¹ respectively, but these are not the correct answers.)

Justify why K₂O has a lower lattice enthalpy (absolute value) than Na₂O.

Which equation represents the standard enthalpy of atomization of bromine, Br₂?

A. $\frac{1}{2}$Br₂ (l) → Br (g)

B. Br₂ (l) → 2Br (g)

C. Br₂ (l) → 2Br (l)

D. $\frac{1}{2}$Br₂ (l) → Br (l)

Which equation represents lattice enthalpy?

A. NaCl (g) → Na⁺ (g) + Cl⁻ (g)

B. NaCl (s) → Na⁺ (g) + Cl⁻ (g)

C. NaCl (s) → Na⁺ (aq) + Cl⁻ (aq)

D. NaCl (s) → Na⁺ (s) + Cl⁻ (s)

Determine the enthalpy of solution of copper(II) chloride, using data from sections 18 and 20 of the data booklet.

The enthalpy of hydration of the copper(II) ion is −2161 kJ mol⁻¹.

What is the order of increasing (more exothermic) enthalpy of hydration?

Xⁿ⁺ (g) → Xⁿ⁺ (aq)

A.  Ca²⁺, Mg²⁺, K⁺, Na⁺

B.  Na⁺, K⁺, Mg²⁺, Ca²⁺

C.  K⁺, Na⁺, Ca²⁺, Mg²⁺

D.  Mg²⁺, Ca²⁺, Na⁺, K⁺

Which combination gives the standard hydration enthalpy of ${\mathrm{Na}}^{+} \left(\mathrm{g}\right)$?

A.  $4+359+790$

B.  $4+359-790$

C.  $-4-359+790$

D.  $4-359+790$

Which substance has the highest lattice enthalpy?

A.  $\text{KCl}$

B.  ${\text{CaCl}}_2$

C.  $\text{KF}$

D.  ${\text{CaF}}_{\text{2}}$

Which represents electron affinity?

A.  Al²⁺(g) → Al³⁺(g) + e⁻

B.  C (g) + e⁻ → C⁻(g)

C.  Cl₂(g) → 2Cl (g)

D.  S (s) → S⁺(g) + e⁻

Deduce the change in enthalpy, ΔH, in kJ, when 56.00 g of ethanol is burned. Use section 13 in the data booklet.

Consider the Born–Haber cycle for the formation of sodium oxide:

What is the lattice enthalpy, in kJ mol⁻¹, of sodium oxide?

A.  414 + 2(108) + 249 + 2(496) − 141 + 790

B.  414 + 2(108) + 249 + 2(496) + 141 + 790

C.  −414 + 2(108) + 249 + 2(496) − 141 + 790

D.  −414 − 2(108) − 249 − 2(496) + 141 − 790

Which equation represents hydration enthalpy?

A.  Na⁺(g) → Na⁺(aq)

B.  Na⁺(aq) → Na⁺(g)

C.  NaCl (s) → NaCl (aq)

D.  NaCl (aq) → NaCl (s)

Determine the standard enthalpy of reaction ($\text{Δ}{H}_{\text{r}}^{⦵}$), in kJ mol⁻¹, for the oxidation of SO₂ to SO₃.

Nitrogen (IV) oxide exists in equilibrium with dinitrogen tetroxide, N₂O₄(g), which is colourless.

2NO₂ (g) ⇌ N₂O₄ (g)

Nitrogen (IV) oxide exists in equilibrium with dinitrogen tetroxide, N₂O₄(g), which is colourless.

2NO₂ (g) ⇌ N₂O₄ (g)

What is ΔH, in kJ, for the reaction N₂H₄ (l) + H₂ (g) → 2NH₃ (g)? 

A.  −18

B.  18

C.  −83

D.  −148

What is the enthalpy of combustion of propan-1-ol, in kJ mol⁻¹, according to the following calorimetry data?

A.  $\frac{-0.015\times 4.2\times 24}{0.075}$

B.  $\frac{-75\times 4.2\times 24}{0.015}$

C.  $\frac{-0.015\times 4.2\times 24}{75}$

D.  $\frac{-75\times 4.2\times 24}{0.015\times 1000}$

Calculate values for the following changes using section 8 of the data booklet.

ΔH_atomisation (Na) = 107 kJ mol⁻¹
ΔH_atomisation (O) = 249 kJ mol⁻¹

$\frac{1}{2}$O₂(g) → O^2- (g):

Na (s) → Na⁺ (g):

The standard enthalpy of formation of sodium oxide is −414 kJ mol⁻¹. Determine the lattice enthalpy of sodium oxide, in kJ mol⁻¹, using section 8 of the data booklet and your answers to (d)(i).

(If you did not get answers to (d)(i), use +850 kJ mol⁻¹ and +600 kJ mol⁻¹ respectively, but these are not the correct answers.)

Justify why K₂O has a lower lattice enthalpy (absolute value) than Na₂O.

Calculate values for the following changes using section 8 of the data booklet.

ΔH_atomisation (Na) = 107 kJ mol⁻¹
ΔH_atomisation (O) = 249 kJ mol⁻¹

$\frac{1}{2}$O₂(g) → O^2- (g):

Na (s) → Na⁺ (g):

The standard enthalpy of formation of sodium oxide is −414 kJ mol⁻¹. Determine the lattice enthalpy of sodium oxide, in kJ mol⁻¹, using section 8 of the data booklet and your answers to (d)(i).

(If you did not get answers to (d)(i), use +850 kJ mol⁻¹ and +600 kJ mol⁻¹ respectively, but these are not the correct answers.)

Justify why K₂O has a lower lattice enthalpy (absolute value) than Na₂O.

Which equation represents the standard enthalpy of atomization of bromine, Br₂?

A. $\frac{1}{2}$Br₂ (l) → Br (g)

B. Br₂ (l) → 2Br (g)

C. Br₂ (l) → 2Br (l)

D. $\frac{1}{2}$Br₂ (l) → Br (l)

Which equation represents lattice enthalpy?

A. NaCl (g) → Na⁺ (g) + Cl⁻ (g)

B. NaCl (s) → Na⁺ (g) + Cl⁻ (g)

C. NaCl (s) → Na⁺ (aq) + Cl⁻ (aq)

D. NaCl (s) → Na⁺ (s) + Cl⁻ (s)

Determine the enthalpy of solution of copper(II) chloride, using data from sections 18 and 20 of the data booklet.

The enthalpy of hydration of the copper(II) ion is −2161 kJ mol⁻¹.

Determine the enthalpy of solution of copper(II) chloride, using data from sections 18 and 20 of the data booklet.

The enthalpy of hydration of the copper(II) ion is −2161 kJ mol⁻¹.

What is the order of increasing (more exothermic) enthalpy of hydration?

Xⁿ⁺ (g) → Xⁿ⁺ (aq)

A.  Ca²⁺, Mg²⁺, K⁺, Na⁺

B.  Na⁺, K⁺, Mg²⁺, Ca²⁺

C.  K⁺, Na⁺, Ca²⁺, Mg²⁺

D.  Mg²⁺, Ca²⁺, Na⁺, K⁺

Which combination gives the standard hydration enthalpy of ${\mathrm{Na}}^{+} \left(\mathrm{g}\right)$?

A.  $4+359+790$

B.  $4+359-790$

C.  $-4-359+790$

D.  $4-359+790$

Which substance has the highest lattice enthalpy?

A.  $\text{KCl}$

B.  ${\text{CaCl}}_2$

C.  $\text{KF}$

D.  ${\text{CaF}}_{\text{2}}$

Which represents electron affinity?

A.  Al²⁺(g) → Al³⁺(g) + e⁻

B.  C (g) + e⁻ → C⁻(g)

C.  Cl₂(g) → 2Cl (g)

D.  S (s) → S⁺(g) + e⁻

Deduce the change in enthalpy, ΔH, in kJ, when 56.00 g of ethanol is burned. Use section 13 in the data booklet.

Deduce the change in enthalpy, ΔH, in kJ, when 56.00 g of ethanol is burned. Use section 13 in the data booklet.

Consider the Born–Haber cycle for the formation of sodium oxide:

What is the lattice enthalpy, in kJ mol⁻¹, of sodium oxide?

A.  414 + 2(108) + 249 + 2(496) − 141 + 790

B.  414 + 2(108) + 249 + 2(496) + 141 + 790

C.  −414 + 2(108) + 249 + 2(496) − 141 + 790

D.  −414 − 2(108) − 249 − 2(496) + 141 − 790

Which equation represents hydration enthalpy?

A.  Na⁺(g) → Na⁺(aq)

B.  Na⁺(aq) → Na⁺(g)

C.  NaCl (s) → NaCl (aq)

D.  NaCl (aq) → NaCl (s)

Which equation represents the lattice enthalpy of magnesium sulfide?

A. MgS (s) → Mg (g) + S (g)

B. MgS (s) → Mg⁺ (g) + S^– (g)

C. MgS (s) → Mg²⁺ (g) + S^2– (g)

D. MgS (s) → Mg (s) + S (s)

What is the enthalpy of solution of MgF₂(s) in kJ mol⁻¹?

Lattice enthalpy of MgF₂(s) = 2926 kJ mol⁻¹

Hydration enthalpy of Mg²⁺(g) = −1963 kJ mol⁻¹

Hydration enthalpy of F⁻(g) = −504 kJ mol⁻¹

A.     2926 − 1963 + 2(−504)

B.     2926 − 1963 − 504

C.     −2926 − (−1963) − (−504)

D.     −2926 − (−1963) − 2(−504)

Which value represents the lattice enthalpy, in kJ mol⁻¹, of strontium chloride, SrCl₂?

A.     – (–829) + 164 + 243 + 550 + 1064 – (–698)

B.     –829 + 164 + 243 + 550 + 1064 – 698

C.     – (–829) + 164 + 243 + 550 + 1064 – 698

D.     –829 + 164 + 243 + 550 + 1064 – (–698)

Determine the standard enthalpy of reaction ($\text{Δ}{H}_{\text{r}}^{⦵}$), in kJ mol⁻¹, for the oxidation of SO₂ to SO₃.

Determine the standard enthalpy of reaction ($\text{Δ}{H}_{\text{r}}^{⦵}$), in kJ mol⁻¹, for the oxidation of SO₂ to SO₃.

Nitrogen (IV) oxide exists in equilibrium with dinitrogen tetroxide, N₂O₄(g), which is colourless.

2NO₂ (g) ⇌ N₂O₄ (g)

Nitrogen (IV) oxide exists in equilibrium with dinitrogen tetroxide, N₂O₄(g), which is colourless.

2NO₂ (g) ⇌ N₂O₄ (g)

Nitrogen (IV) oxide exists in equilibrium with dinitrogen tetroxide, N₂O₄(g), which is colourless.

2NO₂ (g) ⇌ N₂O₄ (g)

Nitrogen (IV) oxide exists in equilibrium with dinitrogen tetroxide, N₂O₄(g), which is colourless.

2NO₂ (g) ⇌ N₂O₄ (g)

What is ΔH, in kJ, for the reaction N₂H₄ (l) + H₂ (g) → 2NH₃ (g)? 

A.  −18

B.  18

C.  −83

D.  −148

What is the enthalpy of combustion of propan-1-ol, in kJ mol⁻¹, according to the following calorimetry data?

A.  $\frac{-0.015\times 4.2\times 24}{0.075}$

B.  $\frac{-75\times 4.2\times 24}{0.015}$

C.  $\frac{-0.015\times 4.2\times 24}{75}$

D.  $\frac{-75\times 4.2\times 24}{0.015\times 1000}$

Which change is exothermic?

A.  $\frac{1}{2}$Cl₂ (g) → Cl (g)

B.   K (g) → K⁺ (g) + e⁻

C.   KCl (s) → K⁺ (g) + Cl⁻ (g)

D.   Cl (g) + e⁻ → Cl⁻ (g)

What is the standard enthalpy of formation, in kJ mol^–1, of IF (g)?

IF₇ (g) + I₂ (s) → IF₅ (g) + 2IF (g)        ΔH${}^{\theta }$ = –89 kJ

ΔH${}_f^{\theta }$ (IF₇) = –941 kJ mol^–1

ΔH${}_f^{\theta }$ (IF₅) = –840 kJ mol^–1

A. –190

B. –95

C. +6

D. +95

The combustion of glucose is exothermic and occurs according to the following equation:

C₆H₁₂O₆ (s) + 6O₂ (g) → 6CO₂ (g) + 6H₂O (g)

Which is correct for this reaction?

Calculate the standard entropy change for this reaction using the following data.

The standard free energy change, ΔG^θ, for the above reaction is –103 kJ mol^–1 at 298 K.

Suggest why ΔG^θ has a large negative value considering the sign of ΔH^θ in part (a).

Calculate the standard entropy change, ΔS^Θ, in J K⁻¹, for the reaction in (ii) using section 12 of the data booklet.

Determine, showing your working, the spontaneity of the reaction in (ii) at 25 °C.

Which system has the most negative entropy change, ΔS, for the forward reaction?

A.     N₂(g) + 3H₂(g) $\rightleftharpoons$ 2NH₃(g)

B.     CaCO₃(s) → CaO(s) + CO₂(g)

C.     2S₂O₃²⁻(aq) + I₂(aq) → S₄O₆²⁻(aq) + 2I^–(aq)

D.     H₂O(l) → H₂O(g)

The table lists standard entropy, S^Θ, values.

Calculate the standard entropy change for the reaction, ΔS^Θ, in J K⁻¹.

CH₄(g) + H₂O(g) → 3H₂(g) + CO(g)

Calculate the standard free energy change, ΔG^Θ, in kJ, for the reaction at 298 K using your answer to (b)(ii).

Determine the temperature, in K, above which the reaction becomes spontaneous.

What are the signs of ΔH^Θ and ΔS^Θ for the reaction, which is spontaneous at low temperature and non-spontaneous at very high temperature?

ΔG^Θ = ΔH^Θ − TΔS^Θ

SO₃ (g) + CaO (s) → CaSO₄ (s)

Predict, giving your reasons, whether the forward reaction is endothermic or exothermic. Use your answers to (b) and (c).

Determine the standard entropy change, in J K−1, for the decomposition of dinitrogen monoxide.

2N₂O (g) → 2N₂ (g) + O₂ (g)

Dinitrogen monoxide has a positive standard enthalpy of formation, ΔH_f^θ.

Deduce, giving reasons, whether altering the temperature would change the spontaneity of the decomposition reaction.

Which is correct for the reaction H₂O (g) → H₂O (l) ?

A. Enthalpy increases and entropy increases.

B. Enthalpy decreases and entropy increases.

C. Enthalpy increases and entropy decreases.

D. Enthalpy decreases and entropy decreases.

Which change has the greatest increase in entropy?

A. CO₂ (s) → CO₂ (g)

B. CO₂ (g) → CO₂ (l)

C. CO₂ (g) → CO₂ (s)

D. CO₂ (l) → CO₂ (s)

Calculate the standard Gibbs free energy change, $\mathrm{\Delta }{G}^{\text{θ}}$, in kJ mol⁻¹, for the first dissociation of citric acid at 298 K, using section 1 of the data booklet.

Comment on the spontaneity of the reaction at 298 K.

Predict, giving a reason, the effect of increasing temperature on the stability of copper(I) chloride solution.

Which reaction has the greatest increase in entropy of the system?

A. HCl (g) + NH₃ (g) → NH₄Cl (s)

B. (NH₄)₂Cr₂O₇ (s) → Cr₂O₃ (s) + N₂ (g) + 4H₂O (g)

C. CaCO₃ (s) → CaO (s) + CO₂ (g)

D. I₂ (g) → I₂ (s)

Which reaction becomes more spontaneous as temperature increases?

A.  ${\mathrm{CaCO}}_3 \left(\mathrm{s}\right)\to \mathrm{CaO} \left(\mathrm{s}\right)+{\mathrm{CO}}_2 \left(\mathrm{g}\right)$

B.  ${\mathrm{N}}_2 \left(\mathrm{g}\right)+3{\mathrm{H}}_2 \left(\mathrm{g}\right)\rightleftharpoons 2{\mathrm{NH}}_3 \left(\mathrm{g}\right)$

C.  $3{\mathrm{CO}}_2 \left(\mathrm{g}\right)+4{\mathrm{H}}_2\mathrm{O} \left(\mathrm{g}\right)\to {\mathrm{C}}_3{\mathrm{H}}_8 \left(\mathrm{g}\right)+5{\mathrm{O}}_2 \left(\mathrm{g}\right)$

D.  ${\mathrm{SO}}_2 \left(\mathrm{g}\right)+{\mathrm{H}}_2{\mathrm{O}}_2 \left(\mathrm{l}\right)\to {\mathrm{H}}_2{\mathrm{SO}}_4 \left(\mathrm{l}\right)$

Calculate the standard Gibbs free energy change, $∆G^{⦵}$, in $\mathbf{kJ}\mathbf{ }{\mathbf{mol}}^{\mathbf{–}\mathbf{1}}$, for the reaction (A to B) at $298 \mathrm{K}$. Use sections 1 and 2 of the data booklet.

Predict, giving a reason, whether the entropy change, $∆S^{⦵}$, for this reaction is negative or positive.

Calculate $∆S^{⦵}$ for the reaction in $\mathrm{J} {\mathrm{K}}^{-1}$, using section 12 of the data booklet.

The standard molar entropy for oxygen gas is $205 \mathrm{J} {\mathrm{K}}^{-1} {\mathrm{mol}}^{-1}$.

Calculate the standard Gibbs free energy change, $∆{\mathrm{G}}^{⦵}$, in $\mathbf{kJ}$, for the reaction at 5 °C, using your answers to (b) and (d). Use section 1 of the data booklet.

(If you did not obtain an answer to (b) or (d) use values of $-1952 \mathrm{kJ}$ and $+113 \mathrm{J} {\mathrm{K}}^{-1}$ respectively, although these are not the correct answers.)

The table shows the variation of standard Gibbs energy with temperature for a reversible reaction.

$∆G^{⦵}=∆H^{⦵}\mathit{-}T∆S^{⦵}$

$∆G^{⦵}=\mathit{-}RT \ln K$

What can be concluded about the reaction?

A.  Equilibrium shifts left as temperature increases.

B.  The forward reaction is more spontaneous below 300 K.

C.  Entropy is higher in the products than in the reactants.

D.  K_c decreases as temperature increases.

Which change results in the largest negative value of ΔS?

A.  C₂H₅OH (l) + SOCl₂(l) → C₂H₅Cl (l) + SO₂(g) + HCl (g)

B.  CaCO₃(s) → CaO (s) + CO₂(g)

C.  H₂O (l) → H₂O (s)

D.  NH₃(g) + HCl (g) → NH₄Cl (s)

Calculate a value for the entropy change, ΔS^⦵, in J K^–1 mol^–1 at 298 K. Use your answers to (e)(i) and section 1 of the data booklet.

If you did not get answers to (e)(i) use –1 kJ, but this is not the correct answer.

Calculate the change in entropy, ΔS, in J K⁻¹, for the decomposition of calcium carbonate.

Determine the temperature, in K, at which the decomposition of calcium carbonate becomes spontaneous, using b(i), b(ii) and section 1 of the data booklet.

(If you do not have answers for b(i) and b(ii), use ΔH = 190 kJ and ΔS = 180 J K⁻¹, but these are not the correct answers.)

Calculate the entropy change, ΔS, in J K⁻¹mol⁻¹, for this reaction.

Chemistry 2e, Chpt. 21 Nuclear Chemistry, Appendix G: Standard Thermodynamic Properties for Selected Substances https://openstax.org/books/chemistry-2e/pages/g-standard-thermodynamic-properties-for- selectedsubstances# page_667adccf-f900-4d86-a13d-409c014086ea © 1999-2021, Rice University. Except where otherwise noted, textbooks on this site are licensed under a Creative Commons Attribution 4.0 International License. (CC BY 4.0) https://creativecommons.org/licenses/by/4.0/.

Calculate the Gibbs free energy change (ΔG), in kJ mol⁻¹, for this reaction at 25 °C. Use section 1 of the data booklet.

If you did not obtain an answer in c(i) or c(ii) use −87.6 kJ mol⁻¹ and −150.5 J mol⁻¹K⁻¹ respectively, but these are not the correct answers.

In which reaction does entropy decrease?

A.  NaCl (s) → NaCl (aq)

B.  Zn (s) + H₂SO₄(aq) → ZnSO₄(aq) + H₂(g)

C.  NH₃(g) + HCl (g) → NH₄Cl (s)

D.  CuCO₃(s) → CuO (s) + CO₂(g)

Calculate the entropy change for the Haber–Bosch process, in J mol^–1K^–1 at 298 K. Use your answer to (b)(i) and section 1 of the data booklet.

Calculate the entropy change of reaction, ΔS^⦵, in J K⁻¹ mol⁻¹.

Predict, giving a reason, how the value of the ΔS^⦵_reaction would be affected if ${\text{I}}_2$ (g) were used as a reactant.

Calculate the Gibbs free energy change, ΔG^⦵, in kJ mol⁻¹, for the reaction at 298 K. Use section 1 of the data booklet.

Which alkane has the lowest standard entropy, S^⦵?

A.  CH₄ (g)

B.  C₂H₆ (g)

C.  C₃H₈ (g)

D.  C₄H₁₀ (g)

Calculate the standard entropy change, ΔS^⦵, for the dissolution of ammonium nitrate.

S^⦵NH₄NO₃ (s) = 151.1 J mol⁻¹ K⁻¹

S^⦵NH₄NO₃(aq) = 259.8 J mol⁻¹ K⁻¹

Calculate the standard Gibbs free energy change, ΔG^⦵, in kJ mol⁻¹, for the dissolution of ammonium nitrate at 298 K. Use sections 1 and 19 of the data booklet as well as your answer for question part (d)(v).

If you did not obtain an answer in (d)(v), use 102.3 J mol⁻¹ K⁻¹, although this is not the correct answer.

Calculate the value of the equilibrium constant for the dissolution of ammonium nitrate at 298 K using the answer to question part (d)(vi) and section 1 of the data booklet.

NH₄NO₃(s) $\rightleftharpoons$ NH₄NO₃(aq)

If you did not obtain an answer in (d)(vi), use −7.84 kJ/mol, although this is not the correct answer.

Nitrogen (IV) oxide exists in equilibrium with dinitrogen tetroxide, N₂O₄(g), which is colourless.

2NO₂ (g) ⇌ N₂O₄ (g)

Nitrogen (IV) oxide exists in equilibrium with dinitrogen tetroxide, N₂O₄(g), which is colourless.

2NO₂ (g) ⇌ N₂O₄ (g)

What is the enthalpy of combustion of propan-1-ol, in kJ mol⁻¹, according to the following calorimetry data?

A.  $\frac{-0.015\times 4.2\times 24}{0.075}$

B.  $\frac{-75\times 4.2\times 24}{0.015}$

C.  $\frac{-0.015\times 4.2\times 24}{75}$

D.  $\frac{-75\times 4.2\times 24}{0.015\times 1000}$

Determine the standard entropy change, in J K−1, for the decomposition of dinitrogen monoxide.

2N₂O (g) → 2N₂ (g) + O₂ (g)

Dinitrogen monoxide has a positive standard enthalpy of formation, ΔH_f^θ.

Deduce, giving reasons, whether altering the temperature would change the spontaneity of the decomposition reaction.

Determine the standard entropy change, in J K−1, for the decomposition of dinitrogen monoxide.

2N₂O (g) → 2N₂ (g) + O₂ (g)

Dinitrogen monoxide has a positive standard enthalpy of formation, ΔH_f^θ.

Deduce, giving reasons, whether altering the temperature would change the spontaneity of the decomposition reaction.

Which is correct for the reaction H₂O (g) → H₂O (l) ?

A. Enthalpy increases and entropy increases.

B. Enthalpy decreases and entropy increases.

C. Enthalpy increases and entropy decreases.

D. Enthalpy decreases and entropy decreases.

Which change has the greatest increase in entropy?

A. CO₂ (s) → CO₂ (g)

B. CO₂ (g) → CO₂ (l)

C. CO₂ (g) → CO₂ (s)

D. CO₂ (l) → CO₂ (s)

Calculate the standard Gibbs free energy change, $\mathrm{\Delta }{G}^{\text{θ}}$, in kJ mol⁻¹, for the first dissociation of citric acid at 298 K, using section 1 of the data booklet.

Comment on the spontaneity of the reaction at 298 K.

Calculate the standard Gibbs free energy change, $\mathrm{\Delta }{G}^{\text{θ}}$, in kJ mol⁻¹, for the first dissociation of citric acid at 298 K, using section 1 of the data booklet.

Comment on the spontaneity of the reaction at 298 K.

Predict, giving a reason, the effect of increasing temperature on the stability of copper(I) chloride solution.

Predict, giving a reason, the effect of increasing temperature on the stability of copper(I) chloride solution.

Which reaction has the greatest increase in entropy of the system?

A. HCl (g) + NH₃ (g) → NH₄Cl (s)

B. (NH₄)₂Cr₂O₇ (s) → Cr₂O₃ (s) + N₂ (g) + 4H₂O (g)

C. CaCO₃ (s) → CaO (s) + CO₂ (g)

D. I₂ (g) → I₂ (s)

Which reaction becomes more spontaneous as temperature increases?

A.  ${\mathrm{CaCO}}_3 \left(\mathrm{s}\right)\to \mathrm{CaO} \left(\mathrm{s}\right)+{\mathrm{CO}}_2 \left(\mathrm{g}\right)$

B.  ${\mathrm{N}}_2 \left(\mathrm{g}\right)+3{\mathrm{H}}_2 \left(\mathrm{g}\right)\rightleftharpoons 2{\mathrm{NH}}_3 \left(\mathrm{g}\right)$

C.  $3{\mathrm{CO}}_2 \left(\mathrm{g}\right)+4{\mathrm{H}}_2\mathrm{O} \left(\mathrm{g}\right)\to {\mathrm{C}}_3{\mathrm{H}}_8 \left(\mathrm{g}\right)+5{\mathrm{O}}_2 \left(\mathrm{g}\right)$

D.  ${\mathrm{SO}}_2 \left(\mathrm{g}\right)+{\mathrm{H}}_2{\mathrm{O}}_2 \left(\mathrm{l}\right)\to {\mathrm{H}}_2{\mathrm{SO}}_4 \left(\mathrm{l}\right)$

Calculate the standard Gibbs free energy change, $∆G^{⦵}$, in $\mathbf{kJ}\mathbf{ }{\mathbf{mol}}^{\mathbf{–}\mathbf{1}}$, for the reaction (A to B) at $298 \mathrm{K}$. Use sections 1 and 2 of the data booklet.

Calculate the standard Gibbs free energy change, $∆G^{⦵}$, in $\mathbf{kJ}\mathbf{ }{\mathbf{mol}}^{\mathbf{–}\mathbf{1}}$, for the reaction (A to B) at $298 \mathrm{K}$. Use sections 1 and 2 of the data booklet.

Predict, giving a reason, whether the entropy change, $∆S^{⦵}$, for this reaction is negative or positive.

Calculate $∆S^{⦵}$ for the reaction in $\mathrm{J} {\mathrm{K}}^{-1}$, using section 12 of the data booklet.

The standard molar entropy for oxygen gas is $205 \mathrm{J} {\mathrm{K}}^{-1} {\mathrm{mol}}^{-1}$.

Calculate the standard Gibbs free energy change, $∆{\mathrm{G}}^{⦵}$, in $\mathbf{kJ}$, for the reaction at 5 °C, using your answers to (b) and (d). Use section 1 of the data booklet.

(If you did not obtain an answer to (b) or (d) use values of $-1952 \mathrm{kJ}$ and $+113 \mathrm{J} {\mathrm{K}}^{-1}$ respectively, although these are not the correct answers.)

Predict, giving a reason, whether the entropy change, $∆S^{⦵}$, for this reaction is negative or positive.

Calculate $∆S^{⦵}$ for the reaction in $\mathrm{J} {\mathrm{K}}^{-1}$, using section 12 of the data booklet.

The standard molar entropy for oxygen gas is $205 \mathrm{J} {\mathrm{K}}^{-1} {\mathrm{mol}}^{-1}$.

Calculate the standard Gibbs free energy change, $∆{\mathrm{G}}^{⦵}$, in $\mathbf{kJ}$, for the reaction at 5 °C, using your answers to (b) and (d). Use section 1 of the data booklet.

(If you did not obtain an answer to (b) or (d) use values of $-1952 \mathrm{kJ}$ and $+113 \mathrm{J} {\mathrm{K}}^{-1}$ respectively, although these are not the correct answers.)

The table shows the variation of standard Gibbs energy with temperature for a reversible reaction.

$∆G^{⦵}=∆H^{⦵}\mathit{-}T∆S^{⦵}$

$∆G^{⦵}=\mathit{-}RT \ln K$

What can be concluded about the reaction?

A.  Equilibrium shifts left as temperature increases.

B.  The forward reaction is more spontaneous below 300 K.

C.  Entropy is higher in the products than in the reactants.

D.  K_c decreases as temperature increases.

Which change results in the largest negative value of ΔS?

A.  C₂H₅OH (l) + SOCl₂(l) → C₂H₅Cl (l) + SO₂(g) + HCl (g)

B.  CaCO₃(s) → CaO (s) + CO₂(g)

C.  H₂O (l) → H₂O (s)

D.  NH₃(g) + HCl (g) → NH₄Cl (s)

Calculate a value for the entropy change, ΔS^⦵, in J K^–1 mol^–1 at 298 K. Use your answers to (e)(i) and section 1 of the data booklet.

If you did not get answers to (e)(i) use –1 kJ, but this is not the correct answer.

Calculate a value for the entropy change, ΔS^⦵, in J K^–1 mol^–1 at 298 K. Use your answers to (e)(i) and section 1 of the data booklet.

If you did not get answers to (e)(i) use –1 kJ, but this is not the correct answer.

Calculate the change in entropy, ΔS, in J K⁻¹, for the decomposition of calcium carbonate.

Determine the temperature, in K, at which the decomposition of calcium carbonate becomes spontaneous, using b(i), b(ii) and section 1 of the data booklet.

(If you do not have answers for b(i) and b(ii), use ΔH = 190 kJ and ΔS = 180 J K⁻¹, but these are not the correct answers.)

Calculate the change in entropy, ΔS, in J K⁻¹, for the decomposition of calcium carbonate.

Determine the temperature, in K, at which the decomposition of calcium carbonate becomes spontaneous, using b(i), b(ii) and section 1 of the data booklet.

(If you do not have answers for b(i) and b(ii), use ΔH = 190 kJ and ΔS = 180 J K⁻¹, but these are not the correct answers.)

Calculate the entropy change, ΔS, in J K⁻¹mol⁻¹, for this reaction.

Chemistry 2e, Chpt. 21 Nuclear Chemistry, Appendix G: Standard Thermodynamic Properties for Selected Substances https://openstax.org/books/chemistry-2e/pages/g-standard-thermodynamic-properties-for- selectedsubstances# page_667adccf-f900-4d86-a13d-409c014086ea © 1999-2021, Rice University. Except where otherwise noted, textbooks on this site are licensed under a Creative Commons Attribution 4.0 International License. (CC BY 4.0) https://creativecommons.org/licenses/by/4.0/.

Calculate the Gibbs free energy change (ΔG), in kJ mol⁻¹, for this reaction at 25 °C. Use section 1 of the data booklet.

If you did not obtain an answer in c(i) or c(ii) use −87.6 kJ mol⁻¹ and −150.5 J mol⁻¹K⁻¹ respectively, but these are not the correct answers.

Calculate the entropy change, ΔS, in J K⁻¹mol⁻¹, for this reaction.

Chemistry 2e, Chpt. 21 Nuclear Chemistry, Appendix G: Standard Thermodynamic Properties for Selected Substances https://openstax.org/books/chemistry-2e/pages/g-standard-thermodynamic-properties-for- selectedsubstances# page_667adccf-f900-4d86-a13d-409c014086ea © 1999-2021, Rice University. Except where otherwise noted, textbooks on this site are licensed under a Creative Commons Attribution 4.0 International License. (CC BY 4.0) https://creativecommons.org/licenses/by/4.0/.

Calculate the Gibbs free energy change (ΔG), in kJ mol⁻¹, for this reaction at 25 °C. Use section 1 of the data booklet.

If you did not obtain an answer in c(i) or c(ii) use −87.6 kJ mol⁻¹ and −150.5 J mol⁻¹K⁻¹ respectively, but these are not the correct answers.

In which reaction does entropy decrease?

A.  NaCl (s) → NaCl (aq)

B.  Zn (s) + H₂SO₄(aq) → ZnSO₄(aq) + H₂(g)

C.  NH₃(g) + HCl (g) → NH₄Cl (s)

D.  CuCO₃(s) → CuO (s) + CO₂(g)

Calculate the entropy change for the Haber–Bosch process, in J mol^–1K^–1 at 298 K. Use your answer to (b)(i) and section 1 of the data booklet.

Calculate the entropy change for the Haber–Bosch process, in J mol^–1K^–1 at 298 K. Use your answer to (b)(i) and section 1 of the data booklet.

Calculate the entropy change of reaction, ΔS^⦵, in J K⁻¹ mol⁻¹.

Predict, giving a reason, how the value of the ΔS^⦵_reaction would be affected if ${\text{I}}_2$ (g) were used as a reactant.

Calculate the Gibbs free energy change, ΔG^⦵, in kJ mol⁻¹, for the reaction at 298 K. Use section 1 of the data booklet.

Calculate the entropy change of reaction, ΔS^⦵, in J K⁻¹ mol⁻¹.

Predict, giving a reason, how the value of the ΔS^⦵_reaction would be affected if ${\text{I}}_2$ (g) were used as a reactant.

Calculate the Gibbs free energy change, ΔG^⦵, in kJ mol⁻¹, for the reaction at 298 K. Use section 1 of the data booklet.

Which alkane has the lowest standard entropy, S^⦵?

A.  CH₄ (g)

B.  C₂H₆ (g)

C.  C₃H₈ (g)

D.  C₄H₁₀ (g)

Calculate the standard entropy change, ΔS^⦵, for the dissolution of ammonium nitrate.

S^⦵NH₄NO₃ (s) = 151.1 J mol⁻¹ K⁻¹

S^⦵NH₄NO₃(aq) = 259.8 J mol⁻¹ K⁻¹

Calculate the standard Gibbs free energy change, ΔG^⦵, in kJ mol⁻¹, for the dissolution of ammonium nitrate at 298 K. Use sections 1 and 19 of the data booklet as well as your answer for question part (d)(v).

If you did not obtain an answer in (d)(v), use 102.3 J mol⁻¹ K⁻¹, although this is not the correct answer.

Calculate the value of the equilibrium constant for the dissolution of ammonium nitrate at 298 K using the answer to question part (d)(vi) and section 1 of the data booklet.

NH₄NO₃(s) $\rightleftharpoons$ NH₄NO₃(aq)

If you did not obtain an answer in (d)(vi), use −7.84 kJ/mol, although this is not the correct answer.

Calculate the standard entropy change, ΔS^⦵, for the dissolution of ammonium nitrate.

S^⦵NH₄NO₃ (s) = 151.1 J mol⁻¹ K⁻¹

S^⦵NH₄NO₃(aq) = 259.8 J mol⁻¹ K⁻¹

Calculate the standard Gibbs free energy change, ΔG^⦵, in kJ mol⁻¹, for the dissolution of ammonium nitrate at 298 K. Use sections 1 and 19 of the data booklet as well as your answer for question part (d)(v).

If you did not obtain an answer in (d)(v), use 102.3 J mol⁻¹ K⁻¹, although this is not the correct answer.

Calculate the value of the equilibrium constant for the dissolution of ammonium nitrate at 298 K using the answer to question part (d)(vi) and section 1 of the data booklet.

NH₄NO₃(s) $\rightleftharpoons$ NH₄NO₃(aq)

If you did not obtain an answer in (d)(vi), use −7.84 kJ/mol, although this is not the correct answer.

What is the standard enthalpy of formation, in kJ mol^–1, of IF (g)?

IF₇ (g) + I₂ (s) → IF₅ (g) + 2IF (g)        ΔH${}^{\theta }$ = –89 kJ

ΔH${}_f^{\theta }$ (IF₇) = –941 kJ mol^–1

ΔH${}_f^{\theta }$ (IF₅) = –840 kJ mol^–1

A. –190

B. –95

C. +6

D. +95

The combustion of glucose is exothermic and occurs according to the following equation:

C₆H₁₂O₆ (s) + 6O₂ (g) → 6CO₂ (g) + 6H₂O (g)

Which is correct for this reaction?

Calculate the standard entropy change for this reaction using the following data.

The standard free energy change, ΔG^θ, for the above reaction is –103 kJ mol^–1 at 298 K.

Suggest why ΔG^θ has a large negative value considering the sign of ΔH^θ in part (a).

Calculate the standard entropy change for this reaction using the following data.

The standard free energy change, ΔG^θ, for the above reaction is –103 kJ mol^–1 at 298 K.

Suggest why ΔG^θ has a large negative value considering the sign of ΔH^θ in part (a).

Calculate the standard entropy change, ΔS^Θ, in J K⁻¹, for the reaction in (ii) using section 12 of the data booklet.

Determine, showing your working, the spontaneity of the reaction in (ii) at 25 °C.

Calculate the standard entropy change, ΔS^Θ, in J K⁻¹, for the reaction in (ii) using section 12 of the data booklet.

Determine, showing your working, the spontaneity of the reaction in (ii) at 25 °C.

Which system has the most negative entropy change, ΔS, for the forward reaction?

A.     N₂(g) + 3H₂(g) $\rightleftharpoons$ 2NH₃(g)

B.     CaCO₃(s) → CaO(s) + CO₂(g)

C.     2S₂O₃²⁻(aq) + I₂(aq) → S₄O₆²⁻(aq) + 2I^–(aq)

D.     H₂O(l) → H₂O(g)

The table lists standard entropy, S^Θ, values.

Calculate the standard entropy change for the reaction, ΔS^Θ, in J K⁻¹.

CH₄(g) + H₂O(g) → 3H₂(g) + CO(g)

Calculate the standard free energy change, ΔG^Θ, in kJ, for the reaction at 298 K using your answer to (b)(ii).

Determine the temperature, in K, above which the reaction becomes spontaneous.

The table lists standard entropy, S^Θ, values.

Calculate the standard entropy change for the reaction, ΔS^Θ, in J K⁻¹.

CH₄(g) + H₂O(g) → 3H₂(g) + CO(g)

Calculate the standard free energy change, ΔG^Θ, in kJ, for the reaction at 298 K using your answer to (b)(ii).

Determine the temperature, in K, above which the reaction becomes spontaneous.

Nitrogen (IV) oxide exists in equilibrium with dinitrogen tetroxide, N₂O₄(g), which is colourless.

2NO₂ (g) ⇌ N₂O₄ (g)

Nitrogen (IV) oxide exists in equilibrium with dinitrogen tetroxide, N₂O₄(g), which is colourless.

2NO₂ (g) ⇌ N₂O₄ (g)

Nitrogen (IV) oxide exists in equilibrium with dinitrogen tetroxide, N₂O₄(g), which is colourless.

2NO₂ (g) ⇌ N₂O₄ (g)

Nitrogen (IV) oxide exists in equilibrium with dinitrogen tetroxide, N₂O₄(g), which is colourless.

2NO₂ (g) ⇌ N₂O₄ (g)

What is the enthalpy of combustion of propan-1-ol, in kJ mol⁻¹, according to the following calorimetry data?

A.  $\frac{-0.015\times 4.2\times 24}{0.075}$

B.  $\frac{-75\times 4.2\times 24}{0.015}$

C.  $\frac{-0.015\times 4.2\times 24}{75}$

D.  $\frac{-75\times 4.2\times 24}{0.015\times 1000}$

What are the signs of ΔH^Θ and ΔS^Θ for the reaction, which is spontaneous at low temperature and non-spontaneous at very high temperature?

ΔG^Θ = ΔH^Θ − TΔS^Θ

SO₃ (g) + CaO (s) → CaSO₄ (s)

Predict, giving your reasons, whether the forward reaction is endothermic or exothermic. Use your answers to (b) and (c).

Predict, giving your reasons, whether the forward reaction is endothermic or exothermic. Use your answers to (b) and (c).