Thermodynamic Terms (AQA A Level Chemistry): Revision Note

Defining Thermodynamic Terms

• Thermodynamics is the branch of physical chemistry concerned with heat energy, temperature, and energy changes in chemical and physical processes

• Energy cycles, such as Born–Haber cycles, are diagrams used to represent enthalpy changes for ionic compounds using Hess’s Law

• To understand these cycles, you must know the definitions of several key enthalpy changes

• The enthalpy change (ΔH) is the heat energy transferred during a reaction at constant pressure

• Key enthalpy definitions include:

  • Enthalpy of formation

  • Ionisation energy

  • Enthalpy of atomisation

  • Bond enthalpy

  • Electron affinity

Enthalpy of formation

• The standard enthalpy of formation (Δ_f H^ꝋ) is the enthalpy change when one mole of a compound is formed from its elements in their standard states under standard conditions

• In this specification, standard conditions are 298 K and a pressure of 100 kPa

• The value of Δ_f H^ꝋ can be positive (endothermic) or negative (exothermic), depending on the balance between bond breaking and bond formation during the reaction

• An equation must show the formation of exactly one mole of the compound from its elements in their standard states

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

Na(s) + ½Cl₂ (g) → NaCl (s)            Δ_f H^ꝋ = -411 kJ mol ^-1

• The enthalpy of formation applies only to compounds

  • By definition, the standard enthalpy of formation of an element in its standard state is zero

Ionisation energy

• The ionisation energy 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 needed 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 ^ꝋ = +500 kJ mol^-1

Enthalpy of atomisation

• The standard enthalpy of atomisation is the enthalpy change when one mole of gaseous atoms is formed from an element in its standard state under standard conditions

• Atomisation is always endothermic because energy is required to break the bonds holding the atoms together in the element

  • Therefore, ΔH ^ꝋ always has a positive value

• An equation for atomisation must show the formation of one mole of gaseous atoms

• For example, sodium is a solid in its standard state. Its standard enthalpy of atomisation is:

Na (s) → Na (g)           ΔH ^ꝋ = +108 kJ mol ^-1

• This represents the energy required to form one mole of gaseous sodium atoms from solid sodium

Bond enthalpy

• Bond enthalpy (also called bond dissociation enthalpy) is the energy required to break one mole of a specific covalent bond in the gaseous state

• It is usually represented as E(bond), for example:

  • E(H–H) is the bond enthalpy of one mole of H–H bonds

• Bond enthalpy values are always quoted as positive because breaking bonds is an endothermic process

• When considering bond formation, the same value is used but with a negative sign, since bond formation is exothermic

• For example, chlorine exists as Cl₂ (g).

  • The bond enthalpy of chlorine is:

Cl₂ (g) → 2Cl (g)    E(Cl-Cl) = +242 kJ mol ^-1

• This looks similar to the atomisation enthalpy of chlorine, but there is an important difference

• Bond enthalpy refers to breaking one mole of bonds

  • Atomisation enthalpy refers to the formation of one mole of gaseous atoms

  • Therefore, the atomisation enthalpy of chlorine is half the bond enthalpy:

½Cl₂ (g) → Cl (g)    ΔH ^ꝋ = +121 kJ mol ^-1

• If an element is a liquid in its standard state, its atomisation enthalpy would also include the enthalpy of vaporisation before bond breaking occurs

Enthalpy of Lattice Formation and Dissociation

• Lattice enthalpy can be defined either as lattice formation enthalpy or lattice dissociation enthalpy

• Because strong electrostatic attractions form between oppositely charged ions in the solid lattice, a large amount of energy is released

  • Therefore, the enthalpy of lattice formation is exothermic and has a large negative value

  • The large negative value indicates that the ionic lattice is much more stable than the separated gaseous ions

  • In the gas phase, the ions are far apart, and there are no strong electrostatic attractions between them

• The more exothermic the lattice enthalpy, the stronger the ionic bonding in the lattice

• Lattice enthalpy cannot be measured directly in a single experiment

  • Instead, it is determined indirectly using a Born–Haber cycle and other experimentally measured enthalpy changes

• For example, for sodium chloride:

Na⁺(g) + Cl^-(g) →NaCl (s)  ΔH^ꝋ = -776 kJ mol ^-1

• This equation represents the enthalpy change when one mole of solid sodium chloride is formed from its gaseous ions

Electron Affinity

• The electron affinity of an element is the enthalpy change when one mole of electrons is added to one mole of gaseous atoms to form one mole of gaseous negative ions under standard conditions

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

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

• The first electron affinity is usually exothermic because energy is released when an electron is attracted to the positively charged nucleus of a gaseous atom

• However, the second electron affinity is always endothermic

  • For example, for oxygen:

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

• This is because energy must be supplied to overcome the electrostatic repulsion between the negatively charged ion and the incoming electron