Bond Enthalpy (DP IB Chemistry: SL)

• When bonds are broken or made enthalpy changes take place
  • A chemical bond is a force of attraction between two atoms
  • Breaking the bond requires the input of energy it is therefore an endothermic process
• The energy change required to break the bond depends on the atoms that form the bond
  • The energy required to break a particular bond is called the bond dissociation enthalpy
  • This is usually just shortened to bond enthalpy or bond energy

To break bonds energy is required from the surroundings and to make new bonds energy is released from the reaction to the surroundings

• The amount of energy released when a particular bond is formed has the same magnitude as the energy taken in when the bond is broken but has the opposite sign

Overall enthalpy changes

• If more energy is released when new bonds are formed than energy is required to break bonds, the reaction is exothermic
  • The products are more stable than the reactants

Bond enthalpy profiles

Average bond energy

• Bond energies are affected by other atoms in the molecule (the environment)
• Therefore, an average of a number of the same type of bond but in different environments is calculated
• This bond energy is known as the average bond energy and is defined as

'The energy needed to break one mole of bonds in a gaseous molecule averaged over similar compounds'

Average bond enthalpy of C-H in methane

• The average bond enthalpy of C-H is found by taking the bond dissociation enthalpy for the whole molecule and dividing it by the number of C-H bonds
• The first C-H bond is easier to break than the second as the remaining hydrogens are pulled more closely to the carbon
• However, since it is impossible to measure the energy of each C-H bond an average is taken
• This value is also compared with a range of similar compounds to obtain an accepted value for the average bond enthalpy