Reactivity 1.4.2—Change in Gibbs energy, ΔG, relates the energy that can be obtained from a chemical reaction to the change in enthalpy, ΔH, change in entropy, ΔS, and absolute temperature, T. Apply the equation ΔG⦵ = ΔH⦵ − TΔS⦵ to calculate unknown values of these terms.

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 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 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 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.

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

2N₂O (g) → 2N₂ (g) + O₂ (g)

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

2N₂O (g) → 2N₂ (g) + O₂ (g)

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

2N₂O (g) → 2N₂ (g) + O₂ (g)

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

2N₂O (g) → 2N₂ (g) + O₂ (g)

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

2N₂O (g) → 2N₂ (g) + O₂ (g)

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

2N₂O (g) → 2N₂ (g) + O₂ (g)

Dinitrogen monoxide in the stratosphere is converted to nitrogen monoxide, NO (g).

Write two equations to show how NO (g) catalyses the decomposition of ozone.

Dinitrogen monoxide in the stratosphere is converted to nitrogen monoxide, NO (g).

Write two equations to show how NO (g) catalyses the decomposition of ozone.

Dinitrogen monoxide in the stratosphere is converted to nitrogen monoxide, NO (g).

Write two equations to show how NO (g) catalyses the decomposition of ozone.

Dinitrogen monoxide in the stratosphere is converted to nitrogen monoxide, NO (g).

Write two equations to show how NO (g) catalyses the decomposition of ozone.

Dinitrogen monoxide in the stratosphere is converted to nitrogen monoxide, NO (g).

Write two equations to show how NO (g) catalyses the decomposition of ozone.

Dinitrogen monoxide in the stratosphere is converted to nitrogen monoxide, NO (g).

Write two equations to show how NO (g) catalyses the decomposition of ozone.

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

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

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

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

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.

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.

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.

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.

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.)

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.)

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.)

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.)

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.)

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 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 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 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 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 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 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 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 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 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.

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

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

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

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

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

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

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

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