Enthalpy

Chemical reactions often have enthalpy change

During reactions, some bonds are broken and some are made. More often than not this will result in energy change (enthalpy change). It is heat energy transferred in a reaction at constant pressure its units are KJ mol^-1.

is used to show that the measurements were made under standard conditions (100kPa/1atm, 298k/25°C) and the elements were in their standard states. (the r shouldn’t be in the image)

Reactions can be either Exothermic or Endothermic

-Exothermic give out energy. Delta H is negative.

-Endothermic absorb energy. Delta H is positive.

Exothermic reactions

Temp often goes up. Oxidation is usually exothermic e.g. the combustion of fuel like methane: CH4 + 202 → CO2 + 2H2O

Energy change is -890KJmol^-1

Endothermic reactions

Temp often falls. The thermal decomposition of Calcium Carbonate: CaCO3 → CaO + CO2

Energy change is +178KJmol^-1

Energy profile diagrams show energy change in reactions

Energy profile diagrams shoe you how enthalpy changes during reactions. The activation energy (E_a), is the minimum amount of energy needed to begin breaking reactant bonds and start a chemical reaction. The less enthalpy a substance has, the more stable it is,

The 4 types of enthalpy change

• Change of reaction

• Change of formation

• Change of combustion

• Change of neutralisation

Standard enthalpy change of reaction

Delta r H, is the enthalpy change when the reaction occurs in the molar quantities shown in the chemical equation, under standard conditions.

Standard enthalpy change of formation

Delta f H, is the enthalpy change when 1 mole of a compound is formed from its elements in their standard states, under standard conditions e.g. 2C + 3H₂ + ½O₂ → C₂H₅OH

Standard enthalpy change of combustion

Delta c H, is the enthalpy change when 1 mole of a substance is completely burned in oxygen, under standard conditions.

Standard enthalpy change of neutralisation

Delta_neutH, is the enthalpy change when an acid and an alkali react together, under standard conditions, to form 1 mole of water.

Reactions are all about making and breaking bonds

You need energy to break bonds, so bond breaking is endothermic (DeltaH is positive).

Energy is released when bonds are formed, so it is exothermic (DeltaH is negative).

Enthalpy change is the overall effect for these two changes.

You need energy to break the attraction between atoms and ions

In ionic bonding, pos & neg ions are attracted. In covalent, pos nuclei attracted to neg charge of shared electrons. The amount of energy required to break the bonds is called bond dissociation enthalpy.

Average bond enthalpies are not exact

It is the average e.g. O-H bonds can have different bond enthalpies.

The energy needed to break one mole of bonds in the gas phase, averaged over many different compounds.

Calculate enthalpy changes using the equation q=mcΔT

q= heat lost or gained (in joules). This is the same as the enthalpy change if the pressure is constant.

m= mass of water (in grams).

c= specific heat capacity of water (4.18J g^-1k^-1).

ΔT= the change is temp of the water in (k)

Can be done in ^oC.

Hess’s law - the total enthalpy is independent of the route taken

Total enthalpy change is the same, no matter which route is taken.

Enthalpy changes can be calculated using average bond enthalpies

Enthalpy change of reaction = Total energy absorbed - Total energy released