Thermodynamic Terms
- Thermodynamics literally means heat and movement and is the branch of physical chemistry that deals with heat, energy, temperature and the physical properties of matter
- Energy cycles are special representations of enthalpy changes for ionic compounds using the principles of Hess's Law
- In order to understand how energy cycles work you need a good knowledge of some key enthalpy change definitions
- Enthalpy change (ΔH) refers to the amount of heat energy transferred during a chemical reaction, at a constant pressure
- The definitions you need to know are:
- enthalpy of formation
- ionisation enthalpy
- enthalpy of atomisation
- bond enthalpy
- electron affinity
Enthalpy of formation
- The enthalpy of formation (ΔHfꝋ) is the enthalpy change when 1 mole of a compound is formed from its elements under standard conditions
- Standard conditions in this syllabus are a temperature of 298 K and a pressure of 100 kPa
- The ΔHfꝋ can be endothermic or exothermic as the energy change is the sum of the bonds broken and formed, so the enthalpy change can have positive or negative values
- Equations can be written to show the standard enthalpy change of formation (ΔHfꝋ) for compounds
- For example, the enthalpy of formation sodium chloride is shown as:
Na (s) + ½Cl2 (g) → NaCl (s) ΔHfꝋ = -411 kJ mol -1
- Notice that enthalpy of formation only refers to compounds
- By definition the enthalpy of formation of elements is zero
Ionisation enthalpy
- The ionisation enthalpy (ΔHieꝋ) of an element is the amount of energy required to remove an electron from a gaseous atom of an element to form a gaseous ion under standard conditions
- Ionisation enthalpy is always endothermic as energy is need to overcome the attraction between an electron and the nucleus
- The first ionisation energy is the energy required to remove one mole of electrons from one mole of gaseous atoms of an element to form one mole of gaseous 1+ ions
- E.g. the first ionisation energy of gaseous sodium:
Na (g) → Na+ (g) + e– ΔHieꝋ = +500 kJ mol-1
Enthalpy change of atomisation
- The standard enthalpy change of atomisation (ΔHatꝋ) is the enthalpy change when 1 mole of gaseous atoms is formed from its element under standard conditions
- The ΔHatꝋ is always endothermic as energy is always required to break any bonds between the atoms in the element, to break the element into its gaseous atoms
- Since this is always an endothermic process, the enthalpy change will always have a positivevalue
- Equations can be written to show the standard enthalpy change of atomisation (ΔHatꝋ) for elements
- For example, sodium in its elemental form is a solid
- The standard enthalpy change of atomisation for sodium is the energy required to form 1 mole of gaseous sodium atoms:
Na (s) → Na (g) ΔHatꝋ = +108 kJ mol -1
Bond enthalpy
- The amount of energy required to break one mole of a specific covalent bond in the gas phase is called the bond dissociation energy
- Bond dissociation energy (E) is usually just simplified to bond energy or bond enthalpy
- In symbols, the type of bond broken is written in brackets after E
- Eg. E (H-H) is the bond energy of a mole of single bonds between two hydrogen atoms
- Bond enthalpy is usually treated as a bond breaking process, so it is quoted in data tables as an endothermic energy change with positive values
- For bond forming processes simply put a negative sign in front of the value
- Equations can be written to show the bond enthalpy
- For example, chlorine in its elemental form is a gas
- The bond enthalpy of chlorine is shown as
Cl2 (g) → 2Cl (g) E(Cl-Cl) = +242 kJ mol -1
- Notice this looks very similar to atomisation enthalpy for chlorine
- However, atomisation enthalpy, by definition, produces 1 mole of atoms, whereas bond enthalpy is expressed per mole of bonds
- So the atomisation enthalpy of chlorine would be half the bond enthalpy
½Cl2 (g) → Cl (g) ΔHatꝋ = +121 kJ mol -1
- If the element was a liquid, instead of a gas, then atomisation enthalpy would also include vaporisation enthalpy - a change of state, before the bonds are broken
Lattice energy
- As with bond enthalpy, lattice enthalpy (ΔHlattꝋ) can be expressed as a formation or dissociation process
- As a formation process, it is the enthalpy change when 1 mole of an ionic compound is formed from its gaseous ions (under standard conditions)
- The ΔHlattꝋ is therefore exothermic, as when ions are combined to form an ionic solid lattice there is an extremely large release of energy
- Since this is an exothermic process, the enthalpy change will have a negative value
- Because of the huge release in energy when the gaseous ions combine, the value will be a very large negative value
- The large negative value of ΔHlattꝋ suggests that the ionic compound is much more stable than its gaseous ions
- This is due to the strong electrostatic forces of attraction between the oppositely charged ions in the solid lattice
- Since there are no electrostatic forces of attraction between the ions in the gas phase, the gaseous ions are less stable than the ions in the ionic lattice
- The more exothermic the value is, the stronger the ionic bonds within the lattice are
- The ΔHlattꝋ of an ionic compound cannot be determined directly by one single experiment
- Multiple experimental values and an energy cycle are used to find the ΔHlattꝋ of ionic compounds
- The lattice energy (ΔHlattꝋ) of an ionic compound can be written as an equation
- For example, sodium chloride is an ionic compound formed from sodium (Na+) and chloride (Cl-) ions
- Since the lattice energy is the enthalpy change when 1 mole of sodium chloride is formed from gaseous sodium and chloride ions, the equation for this process is:
Na+(g) + Cl-(g) → NaCl (s) ΔHlattꝋ = -776 kJ mol -1