Chemical reactions can either release energy to their surroundings, exothermic, or energy can be transferred to them from the surroundings, endothermic.
Energy-level profile diagrams.
Exothermic reactions are most common, however, an important example of an endothermic reaction is photosynthesis in plants, where the energy supplied is from sunlight.
Law of conservation of energy: Energy cannot be destroyed or created but only transferred from one form to another. The total energy of a system of reacting chemicals and surroundings remains constant.
Enthalpy change is the term used to describe the energy exchange that takes place with the surroundings at a constant pressure and is given the symbol ΔH.
Enthalpy is the total energy content of the reacting materials. It is given the symbol, H.
ΔH = ΔH products - ΔH reactants
The units are kilojoules per mole (kJmol-1)
An exothermic enthalpy change is always given a negative value, as energy is lost to the surroundings.
ΔH = -xkJmol-1
An endothermic enthalpy change is always given a positive value, as the energy is gained by the system from the surroundings.
ΔH = + ykJmol-1.
If we are to compare the enthalpy changes of a various reactions we must use standard conditions, such as known temperatures, pressures, amounts and concentrations of reactants or products.
The standard conditions are:
- A pressure of 100kilopascals (102kPa)
- A temperature of 298K (25oC)
- Reactants and products in physical states, normal for the above conditions.
- A concentration of 1.0mol dm-3 for solutions.
The o sign indicates standard conditions.
The standard enthalpy change of reaction is the enthalpy change when the amounts of reactants shown in the equation for the reaction, react under standard conditions to give the products in their standard states.
The standard enthalpy change of formation is the enthalpy change when one mole of a compound is formed from its elements under standard conditions; both compound and elements are in standard states.
The standard enthalpy change of combustion is the enthalpy change when one mole of an element or compound reacts completely with oxygen under standard conditions.
For a chemical reaction to occur bonds must break before new bonds can be made. When bonds break energy is absorbed (endothermic). When bonds form, energy is released (exothermic).
If the energy absorbed whilst making bonds is greater than the energy transferred to the surroundings as bonds are made, then an endothermic reaction will occur.
Whereas, if the energy released on bond formation is greater than that absorbed through breaking bonds then an exothermic reaction is observed.
Bond energy: This is the amount of energy required to break a covalent bond, it indicates the strength of a bond. Values are always quoted as bond energy per mole, E.
O2(g) → 2O(g)
ΔH = +498kJmol-1
Values are always positive, as they refer to bond breaking which is endothermic. They are average values as a particular bond energy will depend upon the molecule the bond is held within, hence average values are taken.