(i) bonds broken: 6 (C–C) / 6 × 346 «kJ» ✔

bonds formed: 3 (C=C) / 3 × 614 «kJ» ✔

$\Delta H^\ominus$ = «6 × 346 kJ − 3 × 614 kJ = 2076 kJ − 1842 kJ =» «+»234 «kJ» ✔

(ii) «$\Delta H = \Sigma\Delta H_\mathrm{f}$(products) − $\Sigma\Delta H_\mathrm{f}$(reactants)»

$\Delta H$ = [3(+0.1) + (−105)] − (−352.7) «kJ mol$^{-1}$» ✔

$\Delta H$ = «+»248.0 «kJ mol$^{-1}$» ✔

(iii) the value from enthalpies of formation, (d)(ii), is the more reliable one because the bond enthalpies in section 12 are averages taken over many compounds

OR

because the bond enthalpies do not apply exactly to the bonds in these particular molecules ✔

(iv) positive AND the number of moles/molecules of gas increases «from 1 to 4» ✔

(v) $\Delta H$ and $\Delta S$ are both positive, so $\Delta G$ «= $\Delta H - T\Delta S$» is negative only when $T\Delta S$

the reaction is therefore not spontaneous at low temperature and becomes spontaneous above a certain temperature / raising the temperature makes $\Delta G$ more negative ✔

(i) [3]: M1 six C–C bonds broken (net); M2 three C=C bonds formed; M3 +234 kJ. Award [3] for a correct final answer with working; a full inventory that also breaks and re-forms the C–H bonds is acceptable if it reaches +234 kJ. (ii) [2]: M1 correct expression with but-1-ene (+0.1) and propane (−105) taken from section 13; M2 +248.0 kJ (accept +248). (iii) [1]. (iv) [1]: sign and reason both required. (v) [2]: M1 signs of ΔH and ΔS with the condition ΔG < 0; M2 spontaneous at high temperature.