Thermodynamic Terms (AQA A Level Chemistry)

Defining 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 (ΔH_f^ꝋ) 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 ΔH_f^ꝋ 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 (ΔH_f^ꝋ) for compounds

• For example, the enthalpy of formation sodium chloride is shown as:

Na(s) + ½Cl₂ (g) → NaCl (s)            ΔH_f^ꝋ = -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 (ΔH_ie^ꝋ) 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^–          ΔH_ie^ꝋ = +500 kJ mol^-1

Enthalpy change of atomisation

• The standard enthalpy change of atomisation (ΔH_at^ꝋ) is the enthalpy change when 1 mole of gaseous atoms is formed from its element under standard conditions

• The ΔH_at^ꝋ 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 positive value

• Equations can be written to show the standard enthalpy change of atomisation (ΔH_at^ꝋ) 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)           ΔH_at^ꝋ = +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)    ΔH_at^ꝋ = +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 (ΔH_latt^ꝋ) can be expressed as a formation or dissociation process, so it is sometimes termed lattice formation enthalpy or lattice dissociation enthalpy

• In this course it is the enthalpy change when 1 mole of a ionic crystal is formed from the gaseous ions (at standard conditions)

• The ΔH_latt^ꝋ is therefore exothermic, as when the ions are combined there is an extremely large release of energy

  • Since this is an exothermic process, the enthalpy change will have a large negative value

• The large negative value of ΔH_latt^ꝋ 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 lattice enthalpy value is, the stronger the ionic bonds within the lattice are

• The ΔH_latt^ꝋ of an ionic compound cannot be determined directly by one single experiment

• Multiple experimental values and an energy cycle are used to find the ΔH_latt^ꝋ of ionic compounds

• The lattice energy (ΔH_latt^ꝋ) 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)  ΔH_latt^ꝋ = -776 kJ mol ^-1

Electron Affinity

• The electron affinity (ΔH_ea^ꝋ) of an element is the energy change when one mole of electrons is gained by one mole of gaseous atoms of an element to form one mole of gaseous ions under standard conditions

• For example, the first electron affinity of chlorine is:

Cl (g)+ e^– → Cl^- (g)          ΔH_ea^ꝋ = -364 kJ mol^-1

• The first electron affinity is always exothermic as energy is released when electrons are attracted to the atoms

• However, the second electron affinity of an element can be endothermic as illustrated by oxygen:

O^– (g) + e^– → O^2- (g)          ΔH_ea^ꝋ = +844 kJ mol^-1

• This is because a large force of repulsion must be overcome between the negatively charged ion and the second electron requiring a large input of energy