NOTE: Since the enthalpy of vapourisation of water is 40.7 kJmol^-1, 3 x 40.7 kJ would be released when the water condenses, giving an enthalpy of combustion of -1442 -(3 x 40.7) = -1564 kJmol^-1, which is very close to the standard value. See Questions 3, 4 and 8 for examples of comparing standard enthalpy of reaction values with those calculated from bond enthalpies.

ΔH_comb(corrected) = ΔH_comb(via bond enthalpies) + ΔH_vap(cyclohexane) - 6 x ΔH_vap(water)

ΔH_comb(corrected) = -3720 + 30.0 - 6 x 40.7 = -3934.2 kJmol^-1

(d) Calculation (a) gave a value of -3915.6 kJmol^-1 for the standard enthalpy of combustion of cyclohexane (298K/1atm).

(a) Apart from (a) none of the answers can be considered as a standard enthalpy values at 298K/1 atm, since bond enthalpy calculations can only involve gaseous species and the fact that bond energies are based on average/typical values for similar molecules.

(b) The value of -3720 kJmol^-1, ignoring state changes, has an error of 195.6 kJmol^-1 compared to the standard enthalpy value based on standard enthalpy changes calculated in (a). This is ~5% error, hardly insignificant!

(c) When corrections are applied to take into account state changes with respect to 298K/1 atm. the error is reduced to 18.6 kJmol^-1 (~0.5%), and not a bad result from typical/average bond enthalpies in organic molecules

ΔH_comb(corrected) = -932 +51.6 - 2 x 40.7 = -3934.2 kJmol^-1 = -961.8 kJmol^-1

(d) Unlike in Q3, the corrected bond energy enthalpy value does not give a more accurate one, than that based solely on the bond enthalpies of gaseous species. In fact, applying the correction, actually makes the value even further from the standard enthalpy change calculated in (a)

I'm not sure why these bond enthalpy calculations are so 'inaccurate' (~7% and ~11%), but there is some uncertainty in the bond energies eg I could not find C-O and C=O bond energies for a carboxylic acid and this easily make a significant difference.

 C₂H₆(g) + I₂(s) ==> C₂H₅I(l) + HI(g)

ΔH^θ_{reaction,298K} = {ΔH^θ_f,298(iodoethane) + ΔH^θ_f,298(hydrogen iodide)} - (ΔH^θ_f,298(ethane)

ΔH^θ_{reaction,298K} = -39.1 +25.9 +84.7 = +71.5 kJmol^-1

(b)

Endothermic changes

to atomise one mole of iodine = 2 x +107 = +214 kJ (remember atomisation refers to 1 mol atoms)

breaking one C-H bond = +412 kJ

total = +626 kJ

Exothermic changes

one C-I bond formed = -228

one HI bond formed = -299

condensation of 1 mole iodoethane = -32

total = -559

Overall enthalpy change ΔH_reaction = +626 -559 = +67 kJmol^-1

(c) The two values are reasonably close with an error of 4.5kJ in (b) compared to the standard enthalpy of reaction value.

NOTE: The error of ~6% is quite significant BUT if any of the bond energy values are out by a few kJ, then an error of that magnitude readily results.

Don't forget to change the positive signs for bond enthalpies of dissociation into negative values of bond formation.

ΔH_sub(C(s)) + 4 x ΔH_{bond diss(H2(g))}

= (3 x 715) + (4 x 436)

= 2145 +  1744 = +3889 kJ

2 x ΔH_{bond form(C-C(g))} + 8 x ΔH_{bond form(C-H(g))}

(2 x -348) + (8 x -412) =

= -696 - 3296 = -3992 kJ