THERMODYNAMICS

• The enthalpy of atomisation is the enthalpy change when 1 mole of gaseous atoms is formed from the element in its standard state
• The enthalpy of atomisation and sublimation are numerically the same for a solid metal.
• The bond dissociation enthalpy is the enthalpy change when 1 mole of a covalent bond is broken into two gaseous atoms or free radicals.
• For diatomic molecules, the bond dissociation enthalpy is equal to 2 times the enthalpy of atomisation.
• The first ionisation enthalpy is the enthalpy change required to remove 1 mole of electrons from 1 mole of gaseous atoms.
• The first electron affinity is exothermic for atoms that normally form negative ions.
• The enthalpy of lattice formation is the enthalpy change when 1 mole of an ionic crystal lattice is formed from its constituent ions in gaseous form.
• Match the enthalpy change with its definition:

  Enthalpy of atomisation ↔️ Formation of gaseous atoms from an element in its standard state

  Bond dissociation enthalpy ↔️ Breaking 1 mole of a covalent bond into gaseous atoms

  Enthalpy of hydration ↔️ Change when gaseous ions become aqueous ions
• The lattice enthalpy cannot be determined directly but can be calculated using a Born-Haber cycle.
• In a Born-Haber cycle calculation, the lattice enthalpy can be calculated using the formula: LattH = fH - (atH + IEH + eaH)LattH.
• The bond energy of a diatomic molecule is equal to 2 times its enthalpy of atomisation.
• The theoretical lattice enthalpy assumes a perfect ionic model where ions are 100% ionic and spherical, and the attractions are purely electrostatic.
• Match the type of lattice enthalpy with its property:

  Born-Haber lattice enthalpy ↔️ Real experimental value

  Theoretical lattice enthalpy ↔️ Assumes a perfect ionic model
• Entropy is a description of the number of ways atoms can share quanta of energy.
• Arrange the states of matter in order of increasing entropy:

  1️⃣ Solid

  2️⃣ Liquid

  3️⃣ Gas
• What is the result of an exothermic reaction in terms of product stability compared to reactants?

  More thermodynamically stable
• Spontaneous reactions always require external influence.

  False
• What does entropy describe at the atomic level?

  Energy sharing arrangements
• High entropy corresponds to a high number of energy arrangement possibilities.
• Substances with more ways of arranging atoms have higher entropy
• Order the following substance types from lowest to highest entropy:

  1️⃣ Elements

  2️⃣ Compounds

  3️⃣ Mixtures
• Which state of matter has the lowest entropy?

  Solid
• Heating a solid increases its entropy due to increased particle vibration.
• The jump in entropy is bigger with boiling than with melting
• What can balanced chemical equations predict about entropy changes?

  Whether ∆S˚ is positive or negative
• A change of state from solid to gas results in a positive entropy change.
• In the reaction NH4Cl (s) → HCl (g) + NH3 (g), the entropy change ∆S˚ is positive
• Match the reaction with its entropy change:

  Na (s) + ½ Cl2 (g) → NaCl (s) ↔️ ∆S˚ = -ve

  NH4Cl (s) → HCl (g) + NH3 (g) ↔️ ∆S˚ = +ve
• What is the minimum standard entropy value for a substance?

  Zero
• Only perfect crystals at absolute zero have zero entropy.
• What is the formula for calculating the standard entropy change ∆S˚?

  ∆S˚ = Σ S˚products - ΣS˚reactants
• The unit of entropy is J K-1 mol-1
• What does Gibbs free energy combine into a single value?

  Enthalpy and entropy
• A negative Gibbs free energy change indicates a spontaneous reaction.
• A reaction with a positive entropy change and a negative enthalpy change will always have a negative ∆G
• What is the formula for calculating Gibbs free energy change ∆G?

  ∆G = ∆H - T∆S
• Temperature in Gibbs free energy calculations must be converted to Kelvin.
• The temperature at which a reaction becomes feasible is calculated when ∆G is equal to zero
• What is the value of ∆G during phase changes like melting or boiling?

  Zero
• A high activation energy can prevent a reaction from occurring even if ∆G is negative.