In  an<unk> , the total energy absorbed in bond breaking is less than the total energy released in bond making. In other words, the energy needed for the reaction to occur is less than the total energy provided. As a result of this, the extra energy is released, usually in the form of heat.
When using a calorimeter, the change in heat of the calorimeter is equal to the opposite of the change in heat of the system. This means that when the medium in which the reaction is taking place gains heat, the reaction is<unk> .
The absolute amount of energy in a chemical system is extremely difficult to measure or calculate. The enthalpy change, ΔH, of a chemical reaction is much easier to measure and calculate. A bomb calorimeter is very suitable for measuring the energy change, ΔH, of a combustion reaction. Measured and calculated ΔH values are related to bond energies by:
by definition the enthalpy change has a negative value:
For an<unk> , this gives a negative value for ΔH, since a larger value (the energy released in the reaction) is subtracted from a smaller value (the energy used for the reaction). For example, when hydrogen burns:
Exothermic reaction
Isothermal process
Organic redox reaction
Dielectric relaxation
Free radical reaction
Condensation reaction
Hydrogen-deuterium exchange
Degenerate energy level