IB Chemistry SL 1.2 Energy Cycles Questions

Practise IB Chemistry SL energy cycles by calculating enthalpy from bond energies, changing Hess equations, and interpreting formation, combustion and solution pathways.

Exam points

Calculate ΔH from mean bond enthalpies by counting bonds broken and formed, then using Σbroken − Σformed.
Explain why bond-enthalpy estimates differ from compound-specific ΔH values, including gaseous-state limits.
Calculate percentage uncertainty in ΔH from supplied uncertainties for each mean bond enthalpy.
Reverse or scale thermochemical equations so unwanted species cancel and the target reaction remains.
Sum signed Hess-cycle steps to calculate target ΔH with consistent signs and units.
Interpret an enthalpy cycle and apply formation, combustion or solution data consistently.

Covered syllabus points

1.2.1—Bond energies 1.2.2—Hess's law

[Maximum number: 4]

Alkenes, such as A (shown below), are important intermediates in the petrochemical industry because they undergo addition reactions to produce a wide variety of products, such as the conversion shown below.

In the gas phase, A reacts with hydrogen to form D.

Determine the enthalpy change, in $\mathrm{kJ} \mathrm{mol}^{-1}$ , for the reaction of A with hydrogen, using Table 10 of the Data Booklet, and state whether the reaction is exothermic or endothermic.

bonds broken: $(\mathrm{E}(\mathrm{C}=\mathrm{C})+\mathrm{E}(\mathrm{H}-\mathrm{H})=612+436=) 1048\left(\mathrm{~kJ} \mathrm{~mol}^{-1}\right)$ ;
Accept (6956+436=) 7392 if all bonds in alkene broken.
bonds formed: $\mathrm{E}(\mathrm{C}-\mathrm{C})+2 \times \mathrm{E}(\mathrm{C}-\mathrm{H})=347+(2 \times 413)=1173\left(\mathrm{~kJ} \mathrm{~mol}^{-1}\right)$ ;

Marking guidance:

Accept 7517 if all the bonds in the product are summed.
$\Delta H=1048-1173 / 7392-7517=-125\left(\mathrm{~kJ} \mathrm{~mol}^{-1}\right) ;$
Award [3] for correct final answer.
Award [2] for +125 .
exothermic;

Apply ECF if sign of $\Delta H$ incorrect.
Do not award a mark for "exothermic" if $\Delta H$ given as positive.

[Maximum number: 1]

The students repeated the experiment using 6.16 g of solid hydrated magnesium sulfate, $\mathrm{MgSO}_{4} \cdot 7 \mathrm{H}_{2} \mathrm{O}(\mathrm{s})$ , and $50.0 \mathrm{~cm}^{3}$ of water. They found the enthalpy change, $\Delta H_{2}$ , to be $+18 \mathrm{~kJ} \mathrm{~mol}^{-1}$ .

The enthalpy of hydration of solid anhydrous magnesium sulfate is difficult to determine experimentally, but can be determined using the diagram below.

\mathrm{MgSO}_{4} \cdot 7 \mathrm{H}_{2} \mathrm{O}(\mathrm{~s}) \xrightarrow[\text { water }]{\Delta H_{2}} \mathrm{Mg}^{2+}(\mathrm{aq})+\mathrm{SO}_{4}^{2-}(\mathrm{aq})

Determine the enthalpy change, $\Delta H$ , in $\mathrm{kJ} \mathrm{mol}^{-1}$ , for the hydration of solid anhydrous magnesium sulfate, $\mathrm{MgSO}_{4}$ .

$\quad \Delta H\left(=\Delta H_{1}-\Delta H_{2}\right)=-99\left(\mathrm{~kJ} \mathrm{~mol}^{-1}\right)$ ;

Marking guidance:

Award [1] if -86 is used giving an answer of $-104\left(\mathrm{~kJ} \mathrm{~mol}^{-1}\right)$ .

All question bank results loaded

Covered syllabus points

Question 1

Question 2