Pearson Edexcel IAL Chemistry Topic 6: Energetics Question Bank

Practise energetics by interpreting enthalpy definitions, diagrams, calorimetry data, Hess cycles and bond enthalpy tables.

Exam points

calculate ΔH from mass, specific heat capacity, temperature change and moles
draw and interpret enthalpy level diagrams with correct arrows, species, signs and units
apply Hess cycles or bond enthalpy tables to find reaction enthalpy changes

Covered syllabus points

6.1—The enthalpy change, ∆H, is the heat energy change measured at constant pressure and that standard conditions are 100 kPa 6.2—That, by convention, exothermic reactions have a negative enthalpy change and endothermic reactions have a positive enthalpy 6.3—Construct and interpret enthalpy level diagrams, showing exothermic and endothermic enthalpy changes 6.4—The definition of standard enthalpy change of: i reaction, ∆rH ii formation, ∆fH iii combustion, ∆cH iv neutralisation 6.5—Experimental data to calculate: i energy transferred in a reaction recalling and using the expression: energy transferred 6.6—Hess’s Law 6.7—CORE PRACTICAL 2 Determination of the enthalpy change of a reaction using Hess’s Law 6.8—Evaluate the results obtained from experiments and comment on sources of error and uncertainty and any assumptions made 6.9—The terms ‘bond enthalpy’ and ‘mean bond enthalpy’6.10—Mean bond enthalpies from enthalpy changes of reaction 6.11—Bond enthalpy data gives some indication about which bond will break first in a reaction, how easy or difficult it is

[Maximum number: 15]

Fuels

Fuels burn in oxygen to release a lot of energy.
Many hydrocarbons and alcohols are used as fuels. During complete combustion, they produce carbon dioxide and water.

Petrol contains 2,2,4-trimethylpentane, an isomer of octane, that promotes smooth combustion.

Figure for Question Topic 6: Energetics question 1 — Edexcel A-Level Chemistry AS

2,2,4-trimethylpentane

Alcohols, such as methanol and ethanol, can be used as fuels either on their own or as additives in petrol.

(a)

The standard enthalpy change of combustion, $\Delta_{\mathrm{c}} H^{\ominus}$ , of 2,2,4-trimethylpentane is $-5461 \mathrm{~kJ} \mathrm{~mol}^{-1}$ .

[ 6 ]

(i)

State the two standard conditions for this enthalpy change.

[ 1 ]

(ii)

Draw a labelled enthalpy level diagram for the complete combustion of 2,2,4-trimethylpentane.

Figure for Question (ii) — Edexcel A-Level Chemistry AS

[ 2 ]

(iii)

Calculate the heat energy released during the complete combustion of $1 \mathrm{dm}^{3}$ of 2,2,4-trimethylpentane.
[Density of 2,2,4-trimethylpentane $=0.692 \mathrm{~g} \mathrm{~cm}^{-3}$ ]

[ 3 ]

calculation of energy given out by 1 g

- calculation of energy given out by $1 \mathrm{~cm}^{3}$

- calculation of energy given out by $1 \mathrm{dm}^{3}$

Example of calculation:

Method 1 enthalpy change $/ \mathrm{g}=\frac{5461}{114}=47.904(\mathrm{~kJ})$

enthalpy change $/ \mathrm{cm}^{3}=47.904 \times 0.692=33.149(\mathrm{~kJ})$

TE on M1

enthalpy change $/ \mathrm{dm}^{3}=33.149 \times 1000$ $=33149 / 33.149 \times 10^{3}(\mathrm{~kJ})$

TE on M2

Method 2

mass of 2,2,4-trimethylpentane in $1 \mathrm{dm}^{3}$ $=0.692 \times 1000=692(\mathrm{~g})(1)$

mol in $1 \mathrm{dm}^{3}=\frac{692}{114}=6.0702(\mathrm{~mol})(1)$

TE on M1

enthalpy change $/ \mathrm{dm}^{3}=6.0702 \times 5461$ $=33149 / 33.149 \times 10^{3}(\mathrm{~kJ})(1)$

TE on M2

Allow alternative methods

Correct answer with some working scores (3)

Ignore SF except 1 SF

Ignore minus sign

Ignore units, even if incorrect

(b)

In an experiment to determine the enthalpy change of combustion of ethanol, $\mathrm{C}_{2} \mathrm{H}_{5} \mathrm{OH}$ , a student used the apparatus shown.

Figure for Question (b) — Edexcel A-Level Chemistry AS

Results:
Mass of water $=100.0 \mathrm{~g}$
Mass of ethanol used $=0.305 \mathrm{~g}$
Temperature rise of water $=13.2^{\circ} \mathrm{C}$

[ 6 ]

(i)

Calculate the enthalpy change of combustion of ethanol.

Give your answer to an appropriate number of significant figures, and include a sign and units.
[Specific heat capacity of water $=4.18 \mathrm{~J} \mathrm{~g}^{-1} \mathrm{C}^{-1}$ ]

[ 4 ]

(ii)

The student looked in a data book and found the actual value for the standard enthalpy change of combustion of ethanol was more exothermic than the experimental value obtained.

Give two reasons for the difference between the data book value and the experimental value, other than referring to standard conditions.

[ 2 ]

(c)

The enthalpy changes for the conversion of four compounds in the gas phase into their constituent atoms are shown.

\begin{aligned} \mathrm{H}_{2} \mathrm{O}(\mathrm{~g}) & \rightarrow 2 \mathrm{H}(\mathrm{~g})+\mathrm{O}(\mathrm{~g}) & \Delta_{\mathrm{r}} H & =+928 \mathrm{~kJ} \mathrm{~mol}^{-1} \\ \mathrm{CH}_{4}(\mathrm{~g}) & \rightarrow \mathrm{C}(\mathrm{~g})+4 \mathrm{H}(\mathrm{~g}) & \Delta_{\mathrm{r}} H & =+1740 \mathrm{~kJ} \mathrm{~mol}^{-1} \\ \mathrm{CH}_{3} \mathrm{OH}(\mathrm{~g}) & \rightarrow \mathrm{C}(\mathrm{~g})+4 \mathrm{H}(\mathrm{~g})+\mathrm{O}(\mathrm{~g}) & \Delta_{\mathrm{r}} H & =+2105 \mathrm{~kJ} \mathrm{~mol}^{-1} \\ \mathrm{C}_{2} \mathrm{H}_{5} \mathrm{OH}(\mathrm{~g}) & \rightarrow 2 \mathrm{C}(\mathrm{~g})+6 \mathrm{H}(\mathrm{~g})+\mathrm{O}(\mathrm{~g}) & \Delta_{\mathrm{r}} H & =+3322 \mathrm{~kJ} \mathrm{~mol}^{-1} \end{aligned}

Calculate the bond enthalpy of the C-C bond, in $\mathrm{kJ} \mathrm{mol}^{-1}$ .
You must show your working.

[ 3 ]

calculation of bond energies of O-H and C-H

- calculation of bond energy of C-O

- calculation of bond energy of C-C

Example of calculation:
Method 1 bond energy $\mathrm{O}-\mathrm{H}=928 / 2=(+) 464\left(\mathrm{~kJ} \mathrm{~mol}^{-1}\right)$
and
bond energy C-H $=1740 / 4=(+) 435\left(\mathrm{~kJ} \mathrm{~mol}^{-1}\right)$
bond energy $\mathrm{C}-\mathrm{O}=2105-(3 \times 435)-464$ $=(+) 336\left(\mathrm{~kJ} \mathrm{~mol}^{-1}\right)$
TE on M1
bond energy $\mathrm{C}-\mathrm{C}=3322-(5 \times 435)-464-336$ $=(+) 347\left(\mathrm{~kJ} \mathrm{~mol}^{-1}\right)$
Method 2
$3322-2105=1217=\mathrm{C}-\mathrm{C}+2 \times \mathrm{C}-\mathrm{H}(1)$
bond energy $2 \times \mathrm{C}-\mathrm{H}=1740 / 2=(+) 870\left(\mathrm{~kJ} \mathrm{~mol}^{-1}\right)(1)$ $\mathrm{C}-\mathrm{C}=1217-870=(+) 347\left(\mathrm{~kJ} \mathrm{~mol}^{-1}\right)(1)$
M3 TE on M1 and M2 in both methods
Correct answer with some working scores

[Maximum number: 9]

This question is about sodium hydroxide.

(a)

State what is meant by standard enthalpy change of neutralisation, $\Delta_{\text {neut }} H^{\ominus}$ .

[ 2 ]

An answer that makes reference to the following points:

- heat energy released under standard conditions (1)

- (when) 1 mol of water is produced (by the reaction of (1) acid with alkali)

Allow enthalpy change under standard conditions
Allow for standard conditions $1 \mathrm{~atm} / 1(.01) \times 10^{5} \mathrm{~Pa}$
and a stated temperature / $298 \mathrm{~K} / 25^{\circ} \mathrm{C}$
Ignore standard states
Do not award required

(2)

(b)

A student carried out an investigation to determine the enthalpy change of neutralisation of aqueous sodium hydroxide by hydrochloric acid.

Method
- separate $25.0 \mathrm{~cm}^{3}$ samples of $0.80 \mathrm{~mol} \mathrm{dm}^{-3}$ sodium hydroxide and $0.80 \mathrm{~mol} \mathrm{dm}^{-3}$ hydrochloric acid were left to reach room temperature
- after two minutes, the solutions were mixed in a copper calorimeter and the temperature was noted at 30 s intervals.

[ 7 ]

(i)

Use the graph shown to determine the maximum temperature change, $\Delta T$ , in this experiment. You must show your working on the graph.

Figure for Question (i) — Edexcel A-Level Chemistry AS

[ 2 ]

An answer that makes reference to the following points:

- two lines of best fit drawn

(1)

- value $\pm 0.2$

(1)

)

Cooling may be shown as straight line or smooth
curve

(2)

(ii)

Calculate the enthalpy change of neutralisation using your answers to (a) and (b)(i). Give a sign and units with your answer.

Assume: no energy is used to heat the container.
the specific heat capacity of the solution $=4.2 \mathrm{~J}^{\circ} \mathrm{C}^{-1} \mathrm{~g}^{-1}$ . the densities of the solutions of NaOH and HCl are $1.0 \mathrm{~g} \mathrm{~cm}^{-3}$ .

[ 3 ]

An answer that makes reference to the following points:

- energy transferred to solutions (1)

- moles of water formed (1)

- enthalpy change of neutralisation with negative (1) sign and units

Example of calculation:
$\begin{aligned} 0.05 \times 4.2 \times 4.4=0.924(\mathrm{~kJ}) 50 \times 4.2 \times 4.4=924(\mathrm{~J}) \end{aligned}$ $(25 \div 1000) \times 0.8=0.02(\mathrm{~mol})$ $0.924 \div 0.02=-46.2 \mathrm{~kJ} \mathrm{~mol}^{-1} /-46,200 \mathrm{~J} \mathrm{~mol}^{-1}$
T E on b(i) and throughout b(ii)
Ignore SF except 1 SF
Comment: Range based on range on graph - $44.10 \mathrm{~kJ} \mathrm{~mol}^{-1}$
to $-48.30 \mathrm{~kJ} \mathrm{~mol}^{-1}$ .

(3)

(iii)

Explain how, if at all, the enthalpy change of neutralisation obtained in (b)(ii) would differ if the heat capacity of the calorimeter was included in the calculation.

[ 2 ]

All question bank results loaded

Covered syllabus points

Topic 6: Energetics question 1

Question (a)

Question (i)

Question (ii)

Question (iii)

Question (b)

Question (i)

Question (ii)

Question (c)

Topic 6: Energetics question 2

Question (a)

Question (b)

Question (i)

Question (ii)

Question (iii)