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Heat of combustion

The heating value (or energy value or calorific value) of a substance, usually a fuel or food (see food energy), is the amount of heat released during the combustion of a specified amount of it.

The calorific value is the total energy released as heat when a substance undergoes complete combustion with oxygen under standard conditions. The chemical reaction is typically a hydrocarbon or other organic molecule reacting with oxygen to form carbon dioxide and water and release heat. It may be expressed with the quantities:

  • energy/mole of fuel
  • energy/mass of fuel
  • energy/volume of the fuel

There are two kinds of enthalpy of combustion, called high(er) and low(er) heat(ing) value, depending on how much the products are allowed to cool and whether compounds like H
2O are allowed to condense. The high heat values are conventionally measured with a bomb calorimeter. Low heat values are calculated from high heat value test data. They may also be calculated as the difference between the heat of formation ΔH⦵
f of the products and reactants (though this approach is somewhat artificial since most heats of formation are typically calculated from measured heats of combustion)..[1]

For a fuel of composition CcHhOoNn, the (higher) heat of combustion is 419 kJ/mol × (c + 0.3 h − 0.5 o) usually to a good approximation (±3%),[2][3] though it gives poor results for some compounds such as (gaseous) formaldehyde and carbon monoxide, and can be significantly off if o + n > c, such as for glycerine dinitrate, C3H6O7N2.[4]

By convention, the (higher) heat of combustion is defined to be the heat released for the complete combustion of a compound in its standard state to form stable products in their standard states: hydrogen is converted to water (in its liquid state), carbon is converted to carbon dioxide gas, and nitrogen is converted to nitrogen gas. That is, the heat of combustion, ΔH°comb, is the heat of reaction of the following process:

C
cH
hN
nO
o (std.) + (c + h⁄4 - o⁄2) O
2 (g) → cCO
2 (g) + h⁄2H
2O (l) + n⁄2N
2 (g)

Chlorine and sulfur are not quite standardized; they are usually assumed to convert to hydrogen chloride gas and SO
2 or SO
3 gas, respectively, or to dilute aqueous hydrochloric and sulfuric acids, respectively, when the combustion is conducted in a bomb calorimeter containing some quantity of water.[5][6]

  1. ^ "Effect of structural conduction and heat loss on combustion in micro-channels". Taylor & Francis Online.
  2. ^ Schmidt-Rohr, Klaus (8 December 2015). "Why Combustions Are Always Exothermic, Yielding About 418 kJ per Mole of O 2". Journal of Chemical Education. 92 (12): 2094–2099. Bibcode:2015JChEd..92.2094S. doi:10.1021/acs.jchemed.5b00333.
  3. ^ Dlugogorski, B. Z.; Mawhinney, J. R.; Duc, V. H. (1994). "The Measurement of Heat Release Rates by Oxygen Consumption Calorimetry in Fires Under Suppression". Fire Safety Science 1007: 877.
  4. ^ It gives 545 kJ/mole, whereas the value calculated from heats of formation is around 1561 kJ/mole. For glycerine trinitrate (nitroglycerine) it gives 0, though nitroglycerine does not actually combust.
  5. ^ Kharasch, M.S. (February 1929). "Heats of combustion of organic compounds". Bureau of Standards Journal of Research. 2 (2): 359. doi:10.6028/jres.002.007.
  6. ^ "Theoretical Analysis of Waste Heat Recovery from an Internal Combustion Engine in a Hybrid Vehicle". Jstor.

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