Given the following reactions: C4H10 (l) + 13/2 O2(g) à 4 CO2(g) + 5 H2O (g) H°

Question

Given the following reactions:
C4H10 (l) + 13/2 O2(g) à 4 CO2(g) + 5 H2O (g) H° = –2635 kJ/mol Butane (MW = 58 g/mol)

C3H7OH (l) + 5 O2(g) à 3 CO2(g) + 4 H2O (g) H° = –1845 kJ/mol Propanol (MW = 60 g/mol)

C2H4O2 (l) + 3 O2(g) à 2 CO2(g) + 2 H2O (g) H° = –785 kJ/mol Acetic acid (MW = 60 g/mol)

Based on the energies of the individual bonds of reactants and products, what is the principle reason that the butane reaction is so much more exothermic?

Butane has the highest %H by mass

Butane consumes the highest number of moles of O2(g)

Butane is the best fuel

Butane produces the highest number of moles of CO2 (g)

Explanation / Answer

Each CO2 ( O=C=O) has two C=O bonds. Each C=O bond is having the highest bond energy that is about 800 kJ/mol. All the remaining bonds have lower bond energies than C=O. Energies of remaining bonds are as follows

C-H = 413 kJ/mol

C-C = 348 kJ/mol

C-O = 358 kJ/mol

O=O = 495 kJ/mol

All the above bonds are in reactants.

Bonds are in products are as follows

C=O = 800 kJ/mol

O-H = 463 kJ/mol

So the compound which can form more number of CO2 will give more amount of energy.

Among the given compounds butane can form more number of CO2 molecules so butane reaction is more exothermic.

Correct reason is

Butane produces the highest number of moles of CO2.

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