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Reference.org
Heat of combustion
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<p>The heating value (or energy value or calorific value) of a <a href="/facts/Chemical_substance/EWcXU41d">substance</a>, usually a <a href="/facts/Fuel/H8DIRL1j">fuel</a> or <a href="/facts/Food/RgNq6NAK">food</a> (see <a href="/facts/Food_energy/TtuzDgbX">food energy</a>), is the amount of <a href="/facts/Heat/FjDGssCs">heat</a> released during the <a href="/facts/Combustion/ohQoWj8z">combustion</a> of a specified amount of it. </p><p>The <i>calorific value</i> is the total <a href="/facts/Energy/wd8PrQM7">energy</a> released as heat when a substance undergoes complete <a href="/facts/Combustion/ohQoWj8z">combustion</a> with <a href="/facts/Oxygen/acyn6o8r">oxygen</a> under <a href="/facts/Standard_temperature_and_pressure/JhJv9KFJ">standard conditions</a>. The chemical reaction is typically a <a href="/facts/Hydrocarbon/Dcg95Ju5">hydrocarbon</a> or other organic molecule reacting with oxygen to form <a href="/facts/Carbon_dioxide/xGzpppEp">carbon dioxide</a> and <a href="/facts/Water/0lkXnJPK">water</a> and release heat. It may be expressed with the quantities: </p> <ul><li>energy/<a href="/facts/Mole_(unit)/kCMFwCc0">mole</a> of fuel</li> <li>energy/mass of fuel</li> <li>energy/volume of the fuel</li></ul> <p>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 H2O are allowed to condense. The high heat values are conventionally measured with a <a href="/facts/Bomb_calorimeter/Njp46oI8">bomb calorimeter</a>. Low heat values are calculated from high heat value test data. They may also be calculated as the difference between the <a href="/facts/Standard_enthalpy_of_formation/UFgGC7UQ">heat of formation</a> Δ<i>H</i>⦵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).. </p><p>For a fuel of composition C<i>c</i>H<i>h</i>O<i>o</i>N<i>n</i>, the (higher) heat of combustion is 419 kJ/mol × (<i>c</i> + 0.3 <i>h</i> − 0.5 <i>o</i>) usually to a good approximation (±3%), though it gives poor results for some compounds such as (gaseous) <a href="/facts/Formaldehyde/Ck8RLfqk">formaldehyde</a> and <a href="/facts/Carbon_monoxide/CKuLtEMJ">carbon monoxide</a>, and can be significantly off if <i>o</i> + <i>n</i> > <i>c</i>, such as for glycerine dinitrate, C3H6O7N2. </p><p>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, Δ<i>H</i>°comb, is the heat of reaction of the following process: </p> C<i>c</i>H<i>h</i>N<i>n</i>O<i>o</i> (std.) + (<i>c</i> + <i>h</i>⁄4 - <i>o</i>⁄2) O2 (g) → <i>c</i>CO2 (g) + <i>h</i>⁄2H2O (<i>l</i>) + <i>n</i>⁄2N2 (g) <p>Chlorine and sulfur are not quite standardized; they are usually assumed to convert to hydrogen chloride gas and SO2 or SO3 gas, respectively, or to dilute aqueous hydrochloric and <a href="/facts/Sulfuric_acid/Wom6fymJ">sulfuric acids</a>, respectively, when the combustion is conducted in a bomb calorimeter containing some quantity of water. </p>