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Chapter 5: Q5.3-74E (page 273)

Both propane and butane are used as gaseous fuels. Which compound produces more heat per gram when burned?

Short Answer

Expert verified

The amount of heat released is more in case of propane compared to butane during the combustion reaction.

Step by step solution

01

Combustion reaction of propane

We have to determine who is going to produce more heat per gram when burned among propane and butane. We will follow the below steps to find that out.

First, we will calculate the amount of heat produced during the combustion of propane.

\(\begin{array}{l}{\rm{The formation enthalpy of }}{{\rm{H}}_{\rm{2}}}{\rm{O(g) is - 241}}{\rm{.82 kJ/mol}}{\rm{.}}\\{\rm{The formation enthalpy of }}{{\rm{O}}_{\rm{2}}}{\rm{(g) is 0 kJ/mol}}{\rm{.}}\\{\rm{The formation enthalpy of C}}{{\rm{O}}_{\rm{2}}}{\rm{(g) is - 393}}{\rm{.51 kJ/mol}}{\rm{.}}\\{\rm{The formation enthalpy of }}{{\rm{C}}_{\rm{3}}}{{\rm{H}}_{\rm{8}}}{\rm{(g) is - 104 kJ/mol}}{\rm{.}}\end{array}\)

\(\begin{array}{l}{\bf{The balanced chemical equation for combustion of propane,}}\\{{\bf{C}}_{\bf{3}}}{{\bf{H}}_{\bf{8}}}{\bf{(g) + 5}}{{\bf{O}}_{\bf{2}}}{\bf{(g) }} \to {\bf{ 3C}}{{\bf{O}}_{\bf{2}}}{\bf{(g) + 4}}{{\bf{H}}_{\bf{2}}}{\bf{O(g)}}\\\\{\rm{Change in enthalpy of the reaction, }}\\{\rm{\Delta }}{{\rm{{\rm H}}}_{{\rm{reaction}}}}{\rm{ = }}\sum {{\rm{\Delta }}{{\rm{{\rm H}}}_{{\rm{products}}}}} {\rm{ - }}\sum {{\rm{\Delta }}{{\rm{{\rm H}}}_{{\rm{reactants}}}}} \\{\rm{\Delta }}{{\rm{{\rm H}}}_{{\rm{reaction}}}}{\rm{ = }}\left( {{\rm{3 \times \Delta }}{{\rm{{\rm H}}}_{{\rm{C}}{{\rm{O}}_{\rm{2}}}{\rm{(g)}}}}{\rm{ + 4 \times \Delta }}{{\rm{{\rm H}}}_{{{\rm{H}}_{\rm{2}}}{\rm{O(g)}}}}} \right){\rm{ - }}\left( {{\rm{\Delta }}{{\rm{{\rm H}}}_{{{\rm{C}}_{\rm{3}}}{{\rm{H}}_{\rm{8}}}{\rm{(g)}}}}{\rm{ + 5 \times \Delta }}{{\rm{{\rm H}}}_{{{\rm{O}}_{\rm{2}}}{\rm{(g) }}}}} \right)\\{\rm{\Delta }}{{\rm{{\rm H}}}_{{\rm{reaction}}}}{\rm{ = }}\left( {{\rm{3 \times - 393}}{\rm{.51 + 4 \times - 241}}{\rm{.82}}} \right){\rm{ - }}\left( {{\rm{ - 104 + 0}}} \right){\rm{ kJ}}\\{\rm{\Delta }}{{\rm{{\rm H}}}_{{\rm{reaction}}}}{\rm{ = - 2147}}{\rm{.81 + 104 kJ}}\\{\rm{\Delta }}{{\rm{{\rm H}}}_{{\rm{reaction}}}}{\rm{ = - 2043}}{\rm{.81 kJ}}\end{array}\)

02

Combustion reaction of butane

Now, we will calculate the amount of heat produced during the combustion of butane.

\(\begin{array}{l}{\rm{The formation enthalpy of }}{{\rm{H}}_{\rm{2}}}{\rm{O(g) is - 241}}{\rm{.82 kJ/mol}}{\rm{.}}\\{\rm{The formation enthalpy of }}{{\rm{O}}_{\rm{2}}}{\rm{(g) is 0 kJ/mol}}{\rm{.}}\\{\rm{The formation enthalpy of C}}{{\rm{O}}_{\rm{2}}}{\rm{(g) is - 393}}{\rm{.51 kJ/mol}}{\rm{.}}\\{\rm{The formation enthalpy of }}{{\rm{C}}_{\rm{4}}}{{\rm{H}}_{{\rm{10}}}}{\rm{(g) is - 126 kJ/mol}}{\rm{.}}\end{array}\)

\(\begin{array}{l}{\bf{The balanced chemical equation for combustion of propane,}}\\{{\bf{C}}_{\bf{4}}}{{\bf{H}}_{{\bf{10}}}}{\bf{(g) + }}\frac{{{\bf{13}}}}{{\bf{2}}}{{\bf{O}}_{\bf{2}}}{\bf{(g) }} \to {\bf{ 4C}}{{\bf{O}}_{\bf{2}}}{\bf{(g) + 5}}{{\bf{H}}_{\bf{2}}}{\bf{O(g)}}\\\\{\rm{Change in enthalpy of the reaction, }}\\{\rm{\Delta }}{{\rm{{\rm H}}}_{{\rm{reaction}}}}{\rm{ = }}\sum {{\rm{\Delta }}{{\rm{{\rm H}}}_{{\rm{products}}}}} {\rm{ - }}\sum {{\rm{\Delta }}{{\rm{{\rm H}}}_{{\rm{reactants}}}}} \\{\rm{\Delta }}{{\rm{{\rm H}}}_{{\rm{reaction}}}}{\rm{ = }}\left( {{\rm{4 \times \Delta }}{{\rm{{\rm H}}}_{{\rm{C}}{{\rm{O}}_{\rm{2}}}{\rm{(g)}}}}{\rm{ + 5 \times \Delta }}{{\rm{{\rm H}}}_{{{\rm{H}}_{\rm{2}}}{\rm{O(g)}}}}} \right){\rm{ - }}\left( {{\rm{\Delta }}{{\rm{{\rm H}}}_{{{\rm{C}}_{\rm{4}}}{{\rm{H}}_{{\rm{10}}}}{\rm{(g)}}}}{\rm{ + }}\frac{{{\rm{13}}}}{{\rm{2}}}{\rm{ \times \Delta }}{{\rm{{\rm H}}}_{{{\rm{O}}_{\rm{2}}}{\rm{(g) }}}}} \right)\\{\rm{\Delta }}{{\rm{{\rm H}}}_{{\rm{reaction}}}}{\rm{ = }}\left( {{\rm{4 \times - 393}}{\rm{.51 + 5 \times - 241}}{\rm{.82}}} \right){\rm{ - }}\left( {{\rm{ - 126 + 0}}} \right){\rm{ kJ}}\\{\rm{\Delta }}{{\rm{{\rm H}}}_{{\rm{reaction}}}}{\rm{ = - 2783}}{\rm{.14 + 126 kJ}}\\{\rm{\Delta }}{{\rm{{\rm H}}}_{{\rm{reaction}}}}{\rm{ = - 2657}}{\rm{.14 kJ}}\end{array}\)

03

Observation

After observing the combustion reaction, we can say that the amount of heat released during the combustion of one mole of propane is 2043.81 kJ, and the amount of heat released during the combustion of one mole of butane is 2657.14 kJ.

Hence, the amount of heat released during the combustion of propane per gram will be:

\(1{\rm{ g }}{{\rm{C}}_3}{{\rm{H}}_8} \times \frac{{1{\rm{ mole of }}{{\rm{C}}_3}{{\rm{H}}_8}{\rm{ }}}}{{44.084{\rm{ g }}{{\rm{C}}_3}{{\rm{H}}_8}}} \times \frac{{ - 2043.81{\rm{ kJ}}}}{{1{\rm{ mole }}{{\rm{C}}_3}{{\rm{H}}_8}}} = - 46.35{\rm{ kJ/gram}}\)

Hence, the amount of heat released during the combustion of butane per gram will be:

\(1{\rm{ g }}{{\rm{C}}_4}{{\rm{H}}_{10}} \times \frac{{1{\rm{ mole of }}{{\rm{C}}_4}{{\rm{H}}_{10}}{\rm{ }}}}{{{\rm{58}}{\rm{.12 g }}{{\rm{C}}_4}{{\rm{H}}_{10}}}} \times \frac{{ - 2657.14{\rm{ kJ}}}}{{1{\rm{ mole }}{{\rm{C}}_4}{{\rm{H}}_{10}}}} = - 45.61{\rm{ kJ/gram}}\)

Hence, the amount of heat released is more in case of propane compared to butane during the combustion reaction.

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Most popular questions from this chapter

Question: How much heat, in joules and in calories, must be added to a 75.0โ€“g iron block with a specific heat of 0.449 J/g ยฐC to increase its temperature from 25 ยฐC to its melting temperature of 1535 ยฐC?

Ethanol, \({{\bf{C}}_{\bf{2}}}{{\bf{H}}_{\bf{5}}}{\bf{OH}}\), is used as a fuel for motor vehicles, particularly in Brazil.

(a) Write the balanced equation for the combustion of ethanol to CO2(g) and H2O(g), and, using the data in Appendix G, calculate the enthalpy of combustion of 1 mole of ethanol.

(b) The density of ethanol is 0.7893 g/ml. Calculate the enthalpy of combustion of exactly 1 L of ethanol.

(c) Assuming that an automobileโ€™s mileage is directly proportional to the heat of combustion of the fuel, calculate how much farther an automobile could be expected to travel on 1 L of gasoline than on 1 L of ethanol. Assume that gasoline has the heat of combustion and the density of nโ€“octane, \({{\bf{C}}_{\bf{8}}}{{\bf{H}}_{{\bf{18}}}}\) (ฮ”Hf=

- 208.4 kJ/mol; density = 0.7025 g/mL).

How many kilojoules of heat will be released when exactly 1 mole of manganese, Mn, is burned to form Mn3O4(s) at standard state conditions?

Which is the least expensive source of energy in kilojoules per dollar; a box of breakfast cereal that weigh 32 ounces and costs \(4.23, or a litre of isooctane (density, 0.6919g/ml) that costs \)0.45? Compare the nutritional value of the cereal with the heat produced by combustion of the isooctane under standard conditions. A 1.0 -ounce serving of the cereal provides 130 calories.

Aluminum chloride can be formed from its elements:

(i)\({\bf{2Al(s) + 3C}}{{\bf{l}}_{\bf{2}}}{\bf{(g)}} \to {\bf{2AlC}}{{\bf{l}}_{\bf{3}}}{\bf{(s) \Delta H^\circ = ?}}\)

Use the reactions here to determine the ฮ”Hยฐ for reaction(i):

\(\begin{array}{*{20}{l}}{\left( {{\bf{ii}}} \right){\rm{ }}{\bf{HCl(g)}} \to {\bf{HCl(aq) \Delta H^\circ (ii) = - 74}}{\bf{.8 kJ}}}\\{\left( {{\bf{iii}}} \right){\rm{ }}{{\bf{H}}_{\bf{2}}}{\bf{(g) + C}}{{\bf{l}}_{\bf{2}}}{\bf{(g)}} \to {\bf{2HCl(g) \Delta H^\circ (iii) = - 185 kJ}}}\\{\left( {{\bf{iv}}} \right){\rm{ }}{\bf{AlC}}{{\bf{l}}_{\bf{3}}}{\bf{(aq)}} \to {\bf{AlC}}{{\bf{l}}_{\bf{3}}}{\bf{(s) \Delta H^\circ (iv) = + 323 kJ}}}\\{\left( {\bf{v}} \right){\rm{ }}{\bf{2Al(s) + 6HCl(aq)}} \to {\bf{2AlC}}{{\bf{l}}_{\bf{3}}}{\bf{(aq) + 3}}{{\bf{H}}_{\bf{2}}}{\bf{(g) \Delta H^\circ (v) = - 1049 kJ}}}\end{array}\)
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