1.8.5 Bond Enthalpy

• 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

• Bond formation is the opposite of bond breaking and so energy is released when bonds are formed
  • It is therefore an exothermic process

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

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

• The relationship between bond breaking and bond making can be shown graphically like this:

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

Exam Tip

• Bond energies are used to find the ΔH_r^ꝋ of a reaction when this cannot be done experimentally
• The process is a step-by-step summation of the bond enthalpies of the all the molecules present finishing with this formula:

Δ_rH^θ = enthalpy change for bonds broken + enthalpy change for bonds formed

• These two worked examples show how to lay out your calculation

Worked Example

Answer:

Step 1: The chemical equation for the Haber process is:

N₂ (g) + 3H₂ (g) ⇌ 2NH₃ (g)

N≡N       3 H-H         6 N-H

Step 2: Set out the calculation as a balance sheet as shown below:

Note! Values for bonds broken are positive (endothermic) and values for bonds formed are negative (exothermic)

Step 3: Calculate the standard enthalpy of reaction

• • Δ_rH^θ = enthalpy change for bonds broken + enthalpy change for bonds formed
  • Δ_rH^θ = (+2253 kJ mol^-1) + (-2346 kJ mol^-1)
  • Δ_rH^θ = -93 kJ mol^-1 

Worked Example

Answer:

Step 1: The enthalpy of combustion is the enthalpy change when one mole of a substance reacts in excess oxygen to produce water and carbon dioxide

The chemical reaction should therefore be simplified such that only one mole of ethyne reacts in excess oxygen:

H-C≡C-H + 2 ½ O=O → H-O-H + 2O=C=O

Step 2: Set out the calculation as a balance sheet as shown below:

• • Δ_rH^θ = enthalpy change for bonds broken + enthalpy change for bonds formed
  • Δ_rH^θ = (+2912 kJ mol^-1) + (- 4142 kJ mol^-1)
  • Δ_rH^θ = -1230 kJ mol^-1

Exam Tip