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A-Level CAIE Chemistry A223.1 Lattice energy and Born-Haber cyclesQuestion Bank

[Maximum number: 2]

Potassium iodide, KI, is used as a reagent in both inorganic and organic chemistry.

(a)

Table 1.1 gives some data about the halide ions, Cl−,Br−\mathrm{Cl}^{-}, \mathrm{Br}^{-}and I−\mathrm{I}^{-}, and their potassium salts.

Table 1.1

Table 1.1

[ 2 ]
(i)

The ionic radius of Pb2+\mathrm{Pb}^{2+} is 0.120 nm compared to 0.133 nm for K+\mathrm{K}^{+}.

Suggest how the ΔHlatt ⊖\Delta H_{\text {latt }}^{\ominus} of PbI2( s)\mathrm{PbI}_{2}(\mathrm{~s}) differs from ΔHlatt ⊖\Delta H_{\text {latt }}^{\ominus} of KI(s).
Explain your answer.

[ 2 ]
(a)
(i)

Write an equation to represent the lattice energy of MgO .

[ 3 ]
(b)

Use the following data, together with appropriate data from the Data Booklet, to calculate a value of ΔHf⊖(MgO)\Delta H_{\mathrm{f}}^{\ominus}(\mathrm{MgO}).

Table
ΔHf⊖(MgO)=\Delta H_{\mathrm{f}}^{\ominus}(\mathrm{MgO})=
[ 3 ]
(a)
(i)

The equation for the formation of a gaseous sulfate ion is shown.

S( s)+2O2( g)+2e−→SO42−(g)ΔH=ΔHf⊖ of SO42−(g)\mathrm{S}(\mathrm{~s})+2 \mathrm{O}_{2}(\mathrm{~g})+2 \mathrm{e}^{-} \rightarrow \mathrm{SO}_{4}{ }^{2-}(\mathrm{g}) \quad \Delta H=\Delta H_{\mathrm{f}}^{\ominus} \text { of } \mathrm{SO}_{4}{ }^{2-}(\mathrm{g})

Calculate the standard enthalpy change of formation, ΔHf⊖\Delta H_{\mathrm{f}}^{\ominus}, of SO42−(g)\mathrm{SO}_{4}{ }^{2-}(\mathrm{g}). It may be helpful to draw a labelled energy cycle. Use relevant data from Table 1.1 in your calculations.

Table 1.1

Table 1.1

ΔHf⊖\Delta H_{\mathrm{f}}^{\ominus} of SO42−(g)=\mathrm{SO}_{4}{ }^{2-}(\mathrm{g})=kJmol−1\mathrm{kJ} \mathrm{mol}^{-1}

[ 3 ]
(ii)

Suggest how the lattice energy of BaSO4( s)\mathrm{BaSO}_{4}(\mathrm{~s}) differs from the lattice energy of Cs2SO4( s)\mathrm{Cs}_{2} \mathrm{SO}_{4}(\mathrm{~s}). Explain your answer.

[ 2 ]
[Maximum number: 4]

Potassium chloride, KCl , and magnesium chloride, MgCl2\mathrm{MgCl}_{2}, are both ionic solids.

Table 1.1

Table 1.1

(a)

Explain the reasons why the lattice energy of MgCl2\mathrm{MgCl}_{2} is more exothermic than the lattice energy of KCl.

[ 2 ]
(b)

Define the following terms.

[ 2 ]
(i)

enthalpy change of atomisation

[ 1 ]
(ii)

first electron affinity

[ 1 ]
[Maximum number: 4]

Calcium chloride, CaCl2\mathrm{CaCl}_{2}, is an ionic solid.
The values of some energy changes are shown in Table 1.1.

Table 1.1

Table 1.1

(a)

Define lattice energy.

[ 1 ]
(b)

Use the data in Table 1.1 to calculate the standard enthalpy change of formation, ΔHf⊖\Delta H_{\mathrm{f}}^{\ominus}, of calcium chloride. It may be helpful to draw an energy cycle. Show all your working.

ΔHf⊖(CaCl2( s))=……………………………...kJ mol−1\Delta H_{\mathrm{f}}^{\ominus}\left(\mathrm{CaCl}_{2}(\mathrm{~s})\right)=\ldots \ldots \ldots \ldots \ldots \ldots \ldots \ldots \ldots \ldots \ldots . . . \mathrm{kJ} \mathrm{~mol}^{-1}
[ 2 ]
(c)

Three possible values for the first electron affinity of bromine are shown in Table 1.2. One of them is correct.

Place a tick by the correct value. Explain your choice.

Table 1.2

Table 1.2

explanation

[ 1 ]
[Maximum number: 4]

Potassium chloride, KCl , and magnesium chloride, MgCl2\mathrm{MgCl}_{2}, are both ionic solids.

Table 1.1

Table 1.1

(a)

Explain the reasons why the lattice energy of MgCl2\mathrm{MgCl}_{2} is more exothermic than the lattice energy of KCl.

[ 2 ]
(b)

Define the following terms.

[ 2 ]
(i)

enthalpy change of atomisation

[ 1 ]
(ii)

first electron affinity

[ 1 ]
(a)

The most common zinc mineral contains zinc(II) sulfide, ZnS .

[ 7 ]
(i)

Complete Fig. 1.2 to show the Born-Haber diagram for the ionic solid ZnS .

Include state symbols of relevant species.

Fig. 1.2

Fig. 1.2

[ 3 ]
(ii)

Describe the trend in the first electron affinity of the Group 16 elements S to Te .

Explain your answer.

[ 2 ]
(iii)

Explain why the lattice energy, ΔHlatt \Delta H_{\text {latt }}, of ZnO is more exothermic than that of ZnS .

[ 2 ]
[Maximum number: 6]

Sodium oxide, Na2O\mathrm{Na}_{2} \mathrm{O}, is a white crystalline solid with a high melting point.

(a)

Use the data below, and other suitable data from the Data Booklet, to calculate the lattice energy of sodium oxide, ΔHlatt ⊖Na2O(s)\Delta H_{\text {latt }}^{\ominus} \mathrm{Na}_{2} \mathrm{O}(\mathrm{s}).

Table
ΔHlatt ⊖Na2O( s)=kJ mol−1\begin{aligned} & \Delta H_{\text {latt }}^{\ominus} \mathrm{Na}_{2} \mathrm{O}(\mathrm{~s})= \\ & \mathrm{kJ} \mathrm{~mol}^{-1} \end{aligned}
[ 4 ]
(b)

State how ΔHlatt ⊖Na2 S( s)\Delta H_{\text {latt }}^{\ominus} \mathrm{Na}_{2} \mathrm{~S}(\mathrm{~s}) differs from ΔHlatt ⊖Na2O(s)\Delta H_{\text {latt }}^{\ominus} \mathrm{Na}_{2} \mathrm{O}(\mathrm{s}).

Indicate this by placing a tick (✓)(\checkmark) in the appropriate box in the table.

Table

Explain your answer.

[ 2 ]
[Maximum number: 5]

paper chromatography,

(a)

Strontium chloride, SrCl2\mathrm{SrCl}_{2}, can be used to produce a red colour in fireworks.

[ 5 ]
(i)

Use the following data, together with relevant data from the Data Booklet, to calculate a value for the lattice energy of strontium chloride. You may find it helpful to construct a Born-Haber cycle.

Table

lattice energy = kJmol−1\mathrm{kJ} \mathrm{mol}^{-1}

[ 5 ]
(a)
(i)

Define lattice energy.

[ 2 ]
(ii)

The lattice energy of the Group 2 carbonates, ΔHlatt θ(MCO3)\Delta H_{\text {latt }}^{\theta}\left(\mathrm{MCO}_{3}\right), becomes less exothermic down the group.

The lattice energy of the Group 2 oxides, ΔHlatt θ(MO)\Delta H_{\text {latt }}^{\theta}(\mathrm{MO}), also becomes less exothermic down the group.
ΔHlatt θ(MCO3)\Delta H_{\text {latt }}^{\theta}\left(\mathrm{MCO}_{3}\right) and ΔHlatt θ(MO)\Delta H_{\text {latt }}^{\theta}(\mathrm{MO}) change by different amounts going down the group.
Suggest how the standard enthalpy change of the decomposition reaction for Group 2 carbonates changes down the group.

Explain your reasoning in terms of the relative sizes of the anions and the relative changes in lattice energy down the group.

[ 2 ]
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