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Enthalpy Of Combustion And Atomisation

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Enthalpy Of Combustion And Atomisation - Lesson Summary

Combustion is a process of burning a chemical compound at high temperatures in presence of excess of oxygen or air.

Combustion reactions are exothermic reactions. The energy released during these reactions is used to drive automobiles, ships, air planes and trains, to operate industrial processes, rocketry, and for numerous other purposes. The energy change accompanying the process of combustion is called the enthalpy of combustion.

It is defined as the amount of heat energy evolved when one mole of substance is completely burnt or oxidised in excess of air or oxygen, and is represented as ΔcH.

Ex:

C(s) + O2(g) → CO2(g); ΔcH = - 393.5 kj mol-1

When 1mole of carbon is completely burnt, it releases 393.5 kilo joules of heat. The negative sign of enthalpy change indicates the exothermic nature of the reaction.

Standard enthalpy of combustion is defined as "the enthalpy change per mole of a substance, when it undergoes combustion, with all the reactants and products being in their standard states at the specified temperature." It is represented by ΔcH not.

The energy released by the combustion of food or fuel is usually compared in terms of their combustion energy per gram, which is known as calorific value.

Calorific value of food or fuel is defined as "the amount of heat released by the complete combustion of one gram of fuel or food." Calorific value is usually expressed in kilo joules per gram or kilo calorie per gram. 

                           Calorific Values of Some Fuels     Fuel     Calorific Value KJ/g     Fuel   Calorific Value KJ/g  Hydrogen          150  Biogas          35-40  Methane          55  Charcoal          33  Butane          50  Ethanol          30  Petrol          50  Coal          25-33  Kerosene          48  Wood          17  Diesel          45  Dung cake          6-8

When one mole, 180 grams of glucose burns, it releases 2840 kilo joules of heat. When one mole of given a substance dissociates into gaseous atoms, the enthalpy change accompanying the process is called the standard enthalpy of atomisation. It is represented by ΔH°.

The conversion of 1mole of hydrogen molecules into hydrogen atoms requires 435.5 kilo joules per mole of energy. Therefore, the enthalpy of atomisation of dihydrogen is 435.5 kilo joules per mole.

The chemical reactions involve the breaking and making of bonds. The breaking of bonds requires energy, while the making of bonds involves the release of energy. The enthalpy changes associated with chemical bonds are expressed in terms of bond dissociation enthalpy and mean bond enthalpy. In case of diatomic molecules, bond dissociation enthalpy is the change in enthalpy when one mole of covalent bonds of a gaseous covalent compound is broken to form products in the gas phase. For all diatomic molecules, bond dissociation enthalpy is the same as atomisation enthalpy.

In case of polyatomic molecules, bond dissociation enthalpy is different for different bonds within the same molecule. Bond enthalpies can be calculated using Hess's law.

The bond enthalpies useful in predicting the enthalpy of a reaction in gas phase.

ΔrHo = ∑ bond Enthalpies reactions - ∑ Bond Enthalpies products
ΔrHo = Standard enthalpy of a reaction

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