Bond Enthalpy

What is bond enthalpy?

Calorimetry is a technique used to measure changes in the enthalpy of chemical reactions, but it does not explain why energy is absorbed or released. Models and theories are used to explain energy changes on the basis of bonds breaking and being formed. 

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

A bond is a force of attraction between two atoms. To overcome this force, energy is required, so bond breaking is endothermic and bond making is exothermic. The bonds in question are covalent bonds and the bond enthalpy definition is:

‘The energy required to break 1 mole of bonds in gaseous covalent molecules under standard conditions'

Since the definition refers to bond breaking, the term is sometimes known as bond dissociation enthalpy.

So, for example, applying the definition to the bond dissociation enthalpy of hydrogen bonds, the equation is:

hydrogen bond enthalpy:   H₂ (g) → 2H (g)   ΔH^ϴ  = +436 kJ mol^-1

We can also apply the term bond dissociation to a whole molecule. For example, the bond dissociation enthalpy of water equation is:

H₂O (g) → 2H (g) + O (g)   ΔH^ϴ  = +497.1 kJ mol^-1

Note: The state symbol of water is (g) as by definition the molecule must be gaseous

What is mean bond enthalpy?

Bond enthalpy varies slightly depending on the influence of other factors such as the shape and structure of the molecule, and the presence of other atoms and groups. Values can be found in a bond enthalpy table or bond enthalpy chart.

This means that a bond enthalpy value is an approximation from a range of molecules and therefore can be considered a mean bond enthalpy or average bond enthalpy. It is found by taking a measured enthalpy change for the dissociation of a molecule and dividing by the number of bonds. Here it is for the C-H bond in methane:

Average bond enthalpy of C-H in methane

How to calculate enthalpy change using bond energies?

A typical use of bond enthalpy data is calculating enthalpy change from bond energies. This is done because there are many reactions whose energy changes cannot be measured experimentally.

You can use a bond enthalpy equation or bond enthalpy formula to find the sum of the bond energies broken and bond energies made.

ΔH^ϴ = Σ [BE bonds broken] + Σ [BE bonds formed]

The bonds broken are always positive terms and the bonds formed are always negative terms, so by adding them together you are finding the next enthalpy change.

A variation of the formula that some people use is: 

ΔH^ϴ = Σ [BE bonds broken] - Σ [BE bonds formed]

This formula does not use positive or negative signs in front of the bond enthalpies, so the overall enthalpy change is the difference between the two summations. It doesn’t matter which formula you use as long as it gives you the correct answer!

• For examples of how to calculate bond enthalpy: "Bond Enthalpy Calculations"

What are energy profiles?

Energy profiles are a way to represent reaction enthalpy changes graphically. For example, a general energy profile for an exothermic reaction  looks like this:

The energy level diagram for exothermic reactions

The transition state at the top of the energy hill represents the point at which all the bonds have been broken and so it is the maximum energy state of the molecule. Bonds have to be broken before they can recombine to form products, so this input of energy is the activation energy of the reaction.

A case study of ozone

A knowledge of enthalpy change using bond energy can help scientists understand the energy flows in the atmosphere that take place during the formation and destruction of ozone.

The structure of oxygen and ozone

The oxygen molecule has a double bond between the two atoms, but the ozone molecule has a delocalised pi-bond, which gives it the equivalent of one and a half oxygen-oxygen bonds in terms of bond strength.

An oxygen-oxygen double bond requires UV light of wavelength 242 nm to break

O₂ (g) → O• (g) + O• (g)   ∆H = +ve, UV light, λ < 242 nm

Remember: The wavelength is inversely proportional to energy, so the shorter the wavelength the greater the energy.

The bond in ozone, requires less energy, so a longer wavelength is needed to break it:

O₃ (g) → O• (g) + O₂ (g)   ∆H = +ve, UV light, λ< 330 nm

At the same time, ozone is formed in an exothermic reaction:

O• (g) + O₂ (g) → O₃ (g)   ∆H = –ve

The temperature in the stratosphere is maintained by the balance of ozone formation and ozone depletion in a process known as the Chapman Cycle

Unfortunately, chlorofluorocarbons from aerosols and refrigerants interfere with this cycle and accelerate the depletion of ozone, so the amount of ozone in the stratosphere has decreased over time leading to increasing levels of harmful UV radiation reaching the Earth’s surface.

• More information: "A case study of ozone & bond strength"

What keyword definitions do I need to know for bond enthalpy?

Some keyword definitions you need to know are:

• Bond enthalpy - the energy required to break 1 mole of bonds in gaseous covalent molecules under standard conditions
• Mean bond enthalpy -  the average energy required to break 1 mole of bonds in gaseous covalent molecules under standard conditions
• Dissociation - breaking apart
• Transition state - is the energy needed to reach the transition state
• Activation energy - the minimum amount of energy needed for reactant molecules to have a successful collision and start the reaction

This is a quick summary of some key concepts on bond enthalpy - remember to go through the full set of revision notes, which are tailored to your specification, to make sure you know everything you need for your exams!