# Enthalpy

### Enthalpy

#### Lessons

In this lesson, we will learn:

• To understand what is meant by exothermic and endothermic reactions in terms of heat.
• To understand how bond enthalpies are used to find total enthalpy changes in reactions.
• To understand how enthalpy changes are represented in chemistry.
Notes:

• Recall the two ways (mentioned in lessons on Introduction to kinetics) to classify chemical reactions in terms of enthalpy – heat energy of a system:
• Exothermic reactions are chemical reactions which have the overall effect of releasing heat energy to the environment. In other words, the energy put in that was needed to break up the reactant bonds was less than the energy given out when the new bonds in the products formed. This is shown by the energy diagram below:
• Endothermic reactions are chemical reactions which have the overall effect of absorbing heat energy from the environment. In other words, the energy given out when the products formed was less than what was needed to break up the bonds in the reactant substances. This is shown by the energy change diagram below:
• The "Energy" referred to in these diagrams above is potential energy. This is stored energy, a sum of attractive and repulsive forces in the atom or molecule being investigated. This is the reason why "high energy" molecules are associated with instability. Molecules or atoms of "low energy" are associated with being more stable.
• Chemical bonds (e.g. ionic and covalent bonds) provide stability, or a stable 'state' to molecules and the atoms they are made of. Therefore in order to break chemical bonds, extra energy must be supplied until those bonds are overcome. Once the bonds are broken, the constituent atoms are not in their stable energy 'state' anymore; they contain more energy than they used to and so are in a higher energy 'state'. Therefore the process of breaking bonds requires energy to be put in.
• When a chemical bond (e.g. a strong covalent bond) forms, its stabilizing effect allows the two atoms forming the bond to have a lower energy 'state' than their separate, higher energy, uncombined states. This surplus energy from their higher energy states is given off to the environment as heat. Therefore the process of forming bonds releases heat energy to the environment.
• Bond enthalpy or bond energy is the average amount of energy required to break one mole of a chemical bond (one mole of molecules of the bond in question). This is usually taken as an average value from a variety of molecules that contain the specific bond being investigated.
• This is the most convenient measure of a bond's strength – a bond with a high bond enthalpy means a lot of energy is required to break this type of bond (for example the N≡N triple bond or the C≡O bond)./li>
• You can use more than one type of enthalpy data to find the total enthalpy change of a reaction. Enthalpy change of a reaction is given the symbol ΔH ('delta' H). Enthalpy change of a system can be found by the following equation:

$\Delta H = H_{Products} - H_{Reactants}$

• Actual enthalpy values for the particular states of chemical substances are rarely found and used. The enthalpy change of a chemical process is though, and can be worked out; you can use bond enthalpy, the enthalpy of formation, or calorimetry. Examples of calculating these will be done in later lessons.
• The sign of the enthalpy change of a reaction is clear evidence of whether a reaction is exothermic or endothermic:
• A ΔH value that is negative means a reaction is exothermic, because the enthalpy of the products is less than the enthalpy of the reactants. This means that the reaction produced lower energy products from higher energy reactants, and the change in heat was released from the system to the surroundings.
• A ΔH value that is positive means a reaction is endothermic, because the enthalpy of the products is less than the enthalpy of the reactants. This means that the reaction produced higher energy products from lower energy reactants, and the change in heat was absorbed from the surroundings.
• Enthalpy changes for chemical reactions can be presented in a number of ways:
• In the form of an equation, for example:
$A + B + 25kJ \to C$ Also written as
$A + B \to C \Delta H = +25 kJ/mol$
• Using an energy change diagram as shown below. The enthalpy change value is marked by the final difference between reactant(s) and product(s) energy.
• Introduction
Enthalpy changes
a)
Exothermic and Endothermic reactions

b)
Bond enthalpy.

c)
Representing enthalpy changes.

• 1.
Recall how to represent the enthalpy change of a reaction.
The reaction between fluorine gas, F2 and hydrogen gas, H2 is as follows. The enthalpy change of the reaction is -539 kJ/mol:

$\mathrm{H_{2(g)} + F_{2(g)} \to 2HF_{(g)}}$

a)
i) How can the enthalpy change, ΔH, of the reaction be shown in a reaction equation?
ii) Name one way in which the enthalpy of reaction could be measured.

b)
How could this enthalpy change be shown on an energy diagram?