3.5.1 Thermodynamics - Enthalpy change (ΔH) definitions
Students should:
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The standard enthalpy of formation
"The enthalpy of formation is the energy change when 1 mole of a substance is formed from its constituent elements in their standard states"
Particular points to note:
- The elements are in their usual states under standard conditions. i.e at 25ºC and 1 atmosphere pressure (100 kPa).
- The enthalpy of formation of an element in its standard state is zero by this definition.
- The energy is given per mole of substance.
- The standard enthalpy of formation is represented by the symbol ΔHf
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Example: Give the equation for the enthalpy of formation of sulfuric acid The enthalpy of formation of sulfuric acid is represented by the following equation:
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ΔHf|
Example: The enthalpy of formation of calcium carbonate is represented by the following equation:
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The sign of the enthalpy of formation of a compound gives us the relative stability of the compound being formed with respect to its constituent elements.
A negative value tells us that it is an exothermic compound (compared to its elements). It DOES NOT tell us that the elements can be made to react directly together to form the compound
Similarly, a positive enthalpy of formation value tells us that the compound is relatively unstable with respect to its elements, it is an endothermic compound.
It DOES NOT mean that it is unstable, only that in comparison with its constituent elements its formation would require energy. Once again there is no implication regarding the possibility of direct formation from the elements.
Ionisation enthalpy
This is defined as the energy required to remove 1 mole of electrons from 1 mole of gaseous particles producing 1 mole of ions. It can be stated as 1st ionisation energy, 2nd ionisation energy etc.
The 1st ionisation energy is the energy required to remove 1 mole of electrons from 1 mole of gaseous atoms to produce 1 mole of singly charged positive ions. It is always endothermic.
Key points to remember are:
- The value is quoted per mole of species
- The particles must be gaseous
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Example: The first ionisation energy of sodium Na(g) |
Enthalpy of atomisation
This is the energy needed to transform an element in its standard state into 1 mole of gaseous atoms. It is always endothermic.
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Example: The enthalpy of atomisation of sodium Na(s) |
Bond dissocation enthalpy
This is energy required to break 1 mole of a specific bond in a specific compound. It is not the same as the bond energy term, which is an average measured over many compounds.
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Example: The first dissociation of methane CH4(g) |
The electron affinity
This is the energy change when 1 mole of electrons are captured by 1 mole of gaseous atoms forming 1 mole of negative ions. Once again the electron affinity may be quoted in terms of 1st electron affinity, 2nd electron affinity etc.
This value is usually negative, but in the case of second electron affinities may be positive, as some processes are exothermic, while others are endothermic.
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Example: The first electron affinity of chlorine Cl(g) |
Lattice enthalpy
The is possible to define from two points of view:
- Forming of 1 mole of a crystal lattice from gaseous ions (exothermic)
- Separating 1 mole of a crystal lattice into gaseous ions at infinite separation (endothermic).
The way you define it is not important providing that you understand the consequences for the enthalpy change in terms of the exo- and endothermic processes.
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Example: The lattice enthalpy of sodium chloride Na+(g) + Cl-(g) Note that in this definition the enthalpy change is exothermic (negative) |
Enthalpy of hydration
This is the energy change when 1 mole of a gaseous species dissolves in an infinite volume of water. It is caused by the solvent water molecules bonding to the dissolving species. It is always an exothermic process as bonds are being formed.
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Example: The hydration enthalpy of the sodium ion Na+(g) + (aq) |
Enthalpy of solution
Definition
The energy change when 1 mole of a substance is dissolved in a solution to infinite dilution.
MgSO4(s) 
Mg2+(aq) + SO42-(aq) |
Clearly, the idea of infinite dilution is not to be taken literally. It just means that a solution is prepared by dissolving the solute to such a dilution as would produce no further energy change.
| Compound | ΔH(solution)/kJ mol-1 | Compound | ΔH(solution)/kJ mol-1 |
| Sodium hydroxide | -44.4 | Ammonium chloride | 14.6 |
| Potassium hydroxide | -57.6 | Ammonium nitrate | 25.7 |
| Sodium chloride | 3.9 | sulfuric acid | -96.2 |
| Sodium bromide | -0.6 | Nitric acid | -33.3 |
| Sodium iodide | -7.5 | Hydrogen fluoride | -61.5 |
| Sodium fluoride | 1.9 | Hydrogen chloride | -74.8 |
| Sodium chloride | 3.9 | Hydrogen bromide | -85.1 |
| Sodium bromide | -0.6 | Hydrogen iodide | -81.7 |
| Sodium iodide | -7.5 | Ethanoic acid | -1.5 |
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Ref: CRC Handbook of chemistry and physics - Edition
44
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Processes involved
At face value, dissolution is a simple process, however it comprises at least two stages.
- 1 The solute is bonded together in some way. These bonds must be broken.
- 2 The separated solute particles must bond to the water (or solvent) molecules.
MgSO4(s)
Mg2+(g) + SO42-(g)
Mg2+(g) + SO42-(g)
+ (aq) Mg2+(aq)
+ SO42-(aq)
If an ionic substance is being dissolved then the individual ions must be dealt with separately in the second stage.