CHEMICAL ENERGETICS TUTORIAL

  
1 (a) (i) Describe and explain the variation in the thermal stabilities of the carbonates of the Group 2

elements.
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(ii) Suggest and
  explain a reason why sodium carbonate is more stable to heat than
magnesium carbonate.
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(b) Sodium hydrogencarbonate, NaHCO3, and potassium hydrogencarbonate, KHCO3, decompose
on heating to produce gases and the solid metal carbonate.

(i) Write an equation for the decomposition of KHCO3.
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(ii) Predict which of NaHCO3 or KHCO3 will decompose at the lower temperature. Explain
your answer.
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(c) (i) Use the data in the table below, and relevant data from the Data Booklet, to calculate the
lattice energy, , of potassium oxide, K2O(s).

energy change value / kJ mol–1

enthalpy change of atomisation of potassium, K(s) +89
electron affinity of O(g) –141
electron affinity of O–(g) +798
enthalpy change of formation of potassium oxide, K2O(s) –361

= .............................. kJ mol–1 [3]

(ii) State whether the lattice energy of Na2O would be more negative, less negative or the
same as that of K2O. Give reasons for your answer.

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[Total: 10]

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1 (a) Calcium has atomic number 20.

Complete the electronic structures for a

calcium atom, 1s22s22p6.................................

calcium ion in the +2 oxidation state. 1s22s22p6.................................

[1]

(b) Calcium nitrate, Ca(NO3)2, is used in fertilisers and can be prepared by an acid-base reaction.

Write an equation for the preparation of calcium nitrate by an acid-base reaction.

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(c) (i) When anhydrous calcium nitrate is heated strongly, it decomposes to leave a white solid.

Identify this white solid and suggest another observation for this reaction.

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(ii) The ease of thermal decomposition of the Group II nitrates decreases down the group.

Explain this trend.

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(d) (i) What is meant by the term standard enthalpy change of hydration, ?

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(ii) Use the following data to calculate the lattice energy, , of calcium nitrate, Ca(NO3)2(s).
You may find it helpful to construct an energy cycle.

enthalpy change value

(Ca2+(g)) –1650 kJ mol–1

(NO3–(g)) –314 kJ mol–1

enthalpy change of solution for Ca(NO3)2(s) –19 kJ mol–1

Ca(NO3)2(s) = ........................... kJ mol–1 [3]

(e) The standard enthalpy change of hydration for Ba2+, (Ba2+(g)), is –1305 kJ mol–1.

Suggest an explanation for why the of the Ba2+ ion is less exothermic than the of
the Ca2+ ion.

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[Total: 12]

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2 Most car air bags contain a capsule of sodium azide, NaN3. In a crash, the NaN3 decomposes into
its elements.

(a) Write an equation for the decomposition of NaN3.

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(b) Complete the ‘dot-and-cross’ diagram for the azide ion, N3–.

Use the following key for the electrons.

 electrons from central nitrogen atom

 electrons from the other two nitrogen atoms
□ added electron(s) responsible for the overall negative charge

N N N

[3]

(c) Lattice energies are always negative showing that they represent exothermic changes.

(i) Explain what is meant by the term lattice energy.

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(ii) Explain why lattice energy represents an exothermic change.

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(iii) Use the following data and any relevant data from the Data Booklet to calculate the
standard enthalpy change of formation, , of NaN3(s).
Include a sign in your answer. Show all your working.

lattice energy, , of NaN3(s) –732 kJ mol–1

standard enthalpy change of atomisation, , of Na(g) +107 kJ mol–1
standard enthalpy change, H o, for 1 12 N2(g) + e–  N3–(g) +142 kJ mol–1

of NaN3(s) = ............................. kJ mol–1 [3]

(iv) The lattice energy, , of RbN3(s) is – 636 kJ mol–1.

Suggest why the lattice energy of NaN3(s), –732 kJ mol–1, is more exothermic than that of
RbN3(s).

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[Total: 11]

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4 The table shows some standard entropy data.

standard entropy, S o

substance
/ J K–1 mol–1
PbO2(s) 77
PbO(s) 69
O2(g) 205

Lead(IV) oxide, PbO2, decomposes to lead(II) oxide, PbO, and oxygen when heated.

2PbO2(s) 2PbO(s) + O2(g) ΔH o = +118 kJ mol–1

(a) Use the data to calculate the value of ΔS o for this reaction.

 ΔS o = .............................. J K–1 mol–1 [2]

(b) Use the value of ΔH o and your answer to (a) to calculate the temperature at which this reaction
becomes feasible.

 T = .............................. K [3]

(c) Solid lead(II) oxide can be made by heating lead metal in air.

Predict the sign of the standard entropy change of this reaction. Explain your answer.

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4 (a) Calcium nitride, Ca3N2 , reacts readily with water to form a white precipitate suspended in an
alkaline solution. The oxidation number of nitrogen does not change during the reaction.

Construct an equation for the reaction of Ca3N2 with water.

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(b) The enthalpy changes of solution,  , of the hydroxides of the Group 2 elements become
less endothermic down the group.

State and explain the trend in the solubilities of the Group 2 hydroxides.

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(c) Complete the energy cycle to show the enthalpy changes that occur in the transformations
between aqueous ions, gaseous ions and an ionic solid.

On your diagram label each enthalpy change with its appropriate symbol; lattice energy,  ,
enthalpy change of hydration,  , or enthalpy change of solution,  .

Complete the three arrows showing the correct direction of each enthalpy change.

....................
aqueous ions gaseous ions

.................... ....................

ionic solid

[3]
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3
(a) The energy cycle shown can be used, along with suitable data, to calculate the enthalpy
change of hydration of Ca2+(g).
Each arrow indicates a transformation, W, X, Y and Z. Each transformation consists of one or
more steps.

Ca2+(g) + 2Cl –(g)

W
Ca(s) + Cl 2(g)

X Z
CaCl 2(s)

Y
CaCl 2(aq)

The following data and data from the Data Booklet should be used.

electron affinity of Cl (g) = –349 kJ mol–1

enthalpy change of atomisation of Ca(s) = +193 kJ mol–1
enthalpy change of formation of CaCl 2(s) = –795 kJ mol–1
enthalpy change of solution of CaCl 2(s) = –83 kJ mol–1
enthalpy change of hydration of Cl –(g) = –364 kJ mol–1

(i) Calculate the value of the enthalpy change corresponding to transformation W.

Show your working.

 enthalpy change W = .............................. kJ mol–1 [2]

(ii) 
Use your answer to (a)(i) and other data to calculate the value of the enthalpy change
corresponding to transformation Z.

 enthalpy change Z = .............................. kJ mol–1 [2]

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(iii) Use your answer to (a)(ii) to calculate the enthalpy change of hydration of Ca2+(g).

 enthalpy change of hydration of Ca2+(g) = .............................. kJ mol–1 [2]

(iv) Write an expression, in terms of W, X, Y and/or Z, to show how the enthalpy changes of
two of the transformations can be used to calculate the lattice energy of CaCl 2(s).

lattice energy of CaCl 2(s) = ........................................................................................... [1]

(v) State whether the lattice energy of CaCl 2(s) is more or less exothermic than the lattice
energy of MgF2(s).

Explain your answer.

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(b) The sulfates of the Group 2 elements vary in solubility down Group 2.

(i) Give the names of two solutions that could be mixed to form barium sulfate.

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(ii) State and explain how the solubilities of the sulfates of the Group 2 elements vary down
Group 2.

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[Total: 13]

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5 (a) The arrangement of the anions around a cation is called the geometry of the cation; e.g. in
[CuCl 4]2– the geometry of copper is tetrahedral and the co-ordination number of copper is 4.

The geometry of a cation in an ionic compound can be predicted from the ratio of the ionic radii
of the cation and anion involved.

cation radius geometry

anion radius of cation

0.155 – 0.225 trigonal planar

0.225 – 0.414 tetrahedral
0.414 – 0.732 octahedral

Use data from the Data Booklet to predict the geometry of, and hence the co-ordination number
of, the cation for

● sodium chloride, NaCl,

geometry of Na+ = ............................. co-ordination number of Na+ = .............................

● magnesium chloride, MgCl 2.

geometry of Mg2+ = ............................. co-ordination number of Mg2+ = .............................

[2]

(b) Magnesium(I) chloride, MgCl, is an unstable compound and readily decomposes as shown.

2MgCl (s) Mg(s) + MgCl 2(s)

Use the following data to calculate the enthalpy change of this reaction.

MgCl (s) = –106 kJ mol–1

MgCl 2(s) = – 642 kJ mol–1

enthalpy change = ............................. kJ mol–1 [1]

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(c) (i) The equation for which ΔH is the lattice energy for MgCl is shown.

Mg+(g) + Cl –(g) MgCl (s)

Use the equation, the following data, and relevant data from the Data Booklet to calculate
a value for the lattice energy of MgCl. You might find it helpful to construct an energy cycle.

electron affinity of Cl (g) = –349 kJ mol–1

enthalpy change of atomisation of Mg(s) = +147 kJ mol–1
enthalpy change of formation of MgCl (s) = –106 kJ mol–1

lattice energy MgCl = ............................. kJ mol–1 [3]

(ii) Suggest how the lattice energies of MgCl 2 and NaCl will compare to that of MgCl.
Explain your answers.

MgCl 2 and MgCl .................................................................................................................

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NaCl and MgCl ....................................................................................................................

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[3]

(d) Define the term electron affinity.

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[Total: 11]
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3 The spontaneity (feasibility) of a chemical reaction depends on the standard Gibbs free energy
change, ∆G o. This is related to the standard enthalpy and entropy changes by the equation shown.

∆G o = ∆H o – T∆S o

(a) State and explain whether the following processes will lead to an increase or decrease in
entropy.

(i) the reaction of magnesium with hydrochloric acid

entropy change ...................................................................................................................

explanation .........................................................................................................................
[1]
(ii) solid potassium chloride dissolving in water

entropy change ...................................................................................................................

explanation .........................................................................................................................
[1]
(iii) steam condensing to water

entropy change ...................................................................................................................

explanation .........................................................................................................................
[1]

(b) Magnesium carbonate can be decomposed.

MgCO3(s) → MgO(s) + CO2(g) ∆H o = +117 kJ mol–1

Standard entropies are shown in the table.

substance MgCO3(s) MgO(s) CO2(g)

S o / J mol–1 K–1 +65.7 +26.9 +214

(i) Calculate ∆G o for this reaction at 298 K.

Include a relevant sign and give your answer to three significant figures.

∆G o = ............................. kJ mol–1 [3]

(ii) Explain, with reference to ∆G o, why this reaction becomes more feasible at higher
temperatures.

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(c) On heating, sodium hydrogencarbonate decomposes into sodium carbonate as shown.

2NaHCO3(s) → Na2CO3(s) + CO2(g) + H2O(g) ∆H o = +130 kJ mol–1

∆S o = +316 J mol–1 K–1

Calculate the minimum temperature at which this reaction becomes spontaneous (feasible).
Show your working.

temperature = ............................. K [2]

(d) The solubility of Group 2 sulfates decreases down the Group.

Explain this trend.

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[Total: 11]

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2 (a) Describe and explain how the solubilities of the sulfates of the Group II elements vary For
down the group. Examiner’s
Use
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(b) The following table lists some enthalpy changes for magnesium and strontium
compounds.

enthalpy change value for magnesium value for strontium

/ kJ mol–1 / kJ mol–1
lattice enthalpy of M (OH)2 –2993 –2467

enthalpy change of hydration of M 2+(g) –1890 –1414

enthalpy change of hydration of OH–(g) –550 –550

(i) Use the above data to calculate values of ΔH solution

o
for Mg(OH)2 and for Sr(OH)2.

Mg(OH)2 ...................................................................................................................

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ΔH solution
o
= ..................................... kJ mol–1

Sr(OH)2 ....................................................................................................................

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ΔH solution
o
= ..................................... kJ mol–1

(ii) Use your results in (i) to suggest whether Sr(OH)2 is more or less soluble in water
than is Mg(OH)2. State any assumptions you make.

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(iii) Suggest whether Sr(OH)2 would be more or less soluble in hot water than in cold.
Explain your reasoning.

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[5]
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(c) Calcium hydroxide, Ca(OH)2, is slightly soluble in water. For

Examiner’s
(i) Write an expression for Ksp for calcium hydroxide, and state its units. Use

Ksp = units ............................................

(ii) 25.0 cm3 of a saturated solution of Ca(OH)2 required 21.0 cm3 of 0.0500 mol dm–3
HCl for complete neutralisation.

Calculate the [OH–(aq)] and the [Ca2+(aq)] in the saturated solution, and hence
calculate a value for Ksp.

[OH–(aq)] = .................................

[Ca2+(aq)] = ................................

Ksp = .........................................................................................

(iii) How would the solubility of Ca(OH)2 in 0.1 mol dm–3 NaOH compare with that in
water?
Explain your answer.

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[6]

[Total: 14]

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8 (a) Describe and explain the trend in the solubility of the hydroxides down Group 2.

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(b) Calcium reacts vigorously with HCl (aq) producing H2(g).

Ca(s) + 2HCl (aq) → CaCl 2(aq) + H2(g)

(i) How would you expect the enthalpy change for this reaction to compare with the enthalpy
change for the reaction where HNO3(aq) is used in place of HCl but all other conditions
are the same?
Explain your answer.

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(ii) The ionic equation for this reaction is shown.

Ca(s) + 2H+(aq) → Ca2+(aq) + H2(g) ∆H o = x kJ mol–1

Construct a fully labelled Hess’ Law cycle to connect each side of this equation to the
relevant gas phase ions.

Use your cycle, the following data, and data from the Data Booklet, to calculate a value for x.

standard enthalpy of atomisation of Ca(s), (Ca) +178 kJ mol–1

standard enthalpy of hydration of Ca2+(g), (Ca2+) –1576 kJ mol–1

standard enthalpy of hydration of H+(g), (H+) –1090 kJ mol–1

x = ............................ kJ mol–1 [4]

(c) The standard enthalpy change for the reaction between Ca(s) and CH3CO2H(aq) is less
negative than x by 2 kJ mol–1.

Suggest an explanation for this.

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[Total: 10]

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(e) (i) Explain what is meant by the term enthalpy change of hydration.

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(ii) S
uggest why the enthalpy change of hydration of Br –(g) is more exothermic than that of
I–(g).

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[Total: 9]
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6 (a) Complete the table by placing one tick () in each row to indicate the sign of each type of
energy change under standard conditions.

always always either negative

energy change
positive negative or positive

bond energy

enthalpy change of formation

[1]

(b) Explain what is meant by the term enthalpy change of atomisation.

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(c) The overall reaction for the atomisation of liquid bromine molecules, Br2(l), is shown.

Br2(l) 2Br(g)

This happens via a two-step process.

● Construct a labelled energy cycle to represent this atomisation process, including state
symbols.
● Use your cycle and relevant data from the Data Booklet to calculate the enthalpy change
of vaporisation of Br2(l), .
The enthalpy change of atomisation of bromine, ∆Hat, = +112 kJ mol–1.

= .............................. kJ mol–1 [3]

(d) Suggest how the of iodine, I2(l), would compare to that of bromine, Br2(l). Explain your
answer.

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(e) (i) Explain what is meant by the term enthalpy change of hydration.

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(ii) S
uggest why the enthalpy change of hydration of Br –(g) is more exothermic than that of
I–(g).

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[Total: 9]
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2 Silicon is the second most abundant element by mass in the Earth’s crust.

(a) In industry, silicon is extracted from SiO2 by reaction with carbon at over 2000 °C.

reaction 1 SiO2(s) + 2C(s) → Si(l) + 2CO(g)

(i) Explain why the entropy change, ∆S, of reaction 1 is positive.

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(ii) Reaction 1 is highly endothermic.

 uggest the effect of an increase in temperature on the feasibility of this reaction.

S
Explain your answer.

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(b) S
 ilicon is purified by first heating it in a stream of HCl (g) to form SiHCl 3. The SiHCl 3 formed is
then distilled to remove other impurities.

reaction 2 Si(s) + 3HCl (g) → SiHCl 3(g) + H2(g)

(i) Table 2.1 shows some standard entropy data.

Table 2.1

compound standard entropy, S o / J K–1 mol–1

Si(s) 19
HCl (g) 187
SiHCl 3(g) 314
H2(g) 131

Use the data in Table 2.1 to calculate ∆S o for reaction 2.

 ∆S o = .............................. J K–1 mol–1 [2]

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(ii) Reaction 3 is the reverse of reaction 2 and is used to obtain pure silicon.

reaction 3 SiHCl 3(g) + H2(g) → Si(s) + 3HCl (g) ∆H = +219.3 kJ mol–1

Use this information and your answer to (b)(i) to calculate the temperature, in K, at which
reaction 3 becomes feasible.
Show your working.

[If you were unable to answer (b)(i), you should use ∆S o = –150 J K–1 mol–1 for reaction 2.
This is not the correct answer to (b)(i).]

 temperature = .............................. K [2]

(c) S ilicon can also be produced by electrolysis of SiO2 dissolved in molten CaCl 2.
The relevant half-equation for the cathode is shown.

SiO2 + 4e– → Si + 2O2–

 alculate the time, in seconds, required to produce 1.00 g of Si by this electrolysis, using a
C
current of 6.00 A.
Assume no other substances are produced at the cathode.

 time = .............................. s [2]

 [Total: 9]

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