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International A LevelChemistryEdexcelExam Questions4: Rates, Equilibria and Further Organic Chemistry4.2 Entropy

Entropy (Edexcel International A Level Chemistry)

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4.2 Entropy

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4.2 Entropy

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1a
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The standard enthalpy change of solution for ammonium nitrate, NH4NO3 , is +25.7 kJ mol–1.

Calculate the value for the standard entropy change in the surroundings, ∆Sθsurroundings, when ammonium nitrate dissolves in water at 298 K. 
Include a sign and units with your answer.

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1b3 marks

Explain what can be deduced from your answer in (a) about the sign and the value of the standard entropy change in the system, ∆Sθsystem, when NH4NO3 dissolves.

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1a
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The equation for the formation of ammonia in the Haber Process is shown

1⁄2N2 (g) +1 1⁄2H2 (g) ⇌ NH3 (g)

At 298 K the standard entropy change of the system, increment S subscript system superscript ⦵ = –98 JK–1mol–1.

Calculate the standard entropy of one mole of ammonia.
Use the value of increment S subscript system superscript ⦵ and the data in the table.

Substance Standard molar entropy, S to the power of ⦵/JK–1mol–1
N2 192
H2 131

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1b2 marks

The value of the total entropy change, ΔStotal , varies with temperature.

Data for the value of ΔStotal at different temperatures but at standard pressure of 100 kPa are given for this reaction.

Temperature / K 1/T / K–1 ΔStotal / kJ K–1mol–1
250 4.00×10–3 8.27×10–2
375 2.67×10–3 2.25×10–2
500 2.00×10–3 –0.764×10–2
625 1.60×10–3 –2.57×10–2
750 1.33×10–3 –3.77×10–2


Plot a graph of ΔStotal against 1/T on the grid.
Include a line of best fit.

q22b

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1c
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The relationship between ΔStotal and 1/T can be found by combining the two equations:

ΔStotal = ΔSsurroundings + ΔSsystem

and ΔSsurroundings = –ΔH/T
to give

ΔStotal = –ΔH/T + ΔSsystem

i)

Determine the gradient of the line plotted in (b), including units in your answer.

(1)

ii)
Identify the thermodynamic quantity that can be obtained from this gradient.

(1)

iii)
Determine the temperature at which the reaction ceases to be thermodynamically feasible at a pressure of 100kPa.

(1)

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1d
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The industrial synthesis of ammonia

1⁄2N2 (g) +1 1⁄2H2 (g) ⇌ NH3 (g)

is carried out at pressures of about 20000kPa and temperatures between 700 K and 750 K. These temperatures are higher than the answer to (c)(iii).

i)

State the relationship between the total entropy, ΔStotal , and the equilibrium constant, K.

(1)

ii)
Calculate the value of the equilibrium constant K at 750 K.
[ΔStotal at 750K = –37.7JK–1 mol–1]

(2)

iii)

Explain why ΔStotal decreases with an increase in temperature.

(3)

iv)

State how the Haber Process is made economically feasible at 750 K even though the total entropy change is negative.

(1)

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1e6 marks

Ammonia from the Haber Process reacts with acids.
With phosphoric acid, H3PO4, a number of products are formed in solution.
One of these is the fertiliser diammonium hydrogenphosphate.

i)

Write an equation for the production of this fertiliser.
State symbols are not required.

(2)

ii)

Write an ionic equation to show that ammonium ions are acidic in aqueous solution.
State symbols are not required.

(1)

iii)

A solution containing both ammonia and ammonium ions acts as a buffer.
Explain, using a relevant ionic equation, the effect of adding a small amount
of acid to this buffer.

(3)

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