5.2.2 Enthalpy and Entropy

• What is lattice enthalpy?


  • Enthalpy change when 1 mole of an ionic compound is formed from its gaseous ions under standard conditions

  Na+ (g) + Cl- (g) --> NaCl (s)
• What is the standard enthalpy change of atomisation?


  • Enthalpy change when 1 mole of gaseous atoms forms from the element in its standard state

  1/2 Cl2 (g) --> Cl (g)

  Na (s) --> Na (g)
• What is first electron affinity?


  • Enthalpy change when 1 electron added to each atom in 1 mole of gaseous atoms to form 1 mole of gaseous 1- ions

  Cl (g) + e- --> Cl- (g)
• What is second electron affinity?


  • Enthalpy change when 1 electron is added to each ion in 1 mole of gaseous 1- ions to form 1 mole of gaseous 2- ions
• What is the standard enthalpy change of solution?


  • The enthalpy change when one mole of a solute dissolves in a solvent

  NaCl (s) + aq --> Na+ (aq) + Cl- (aq)
• What is the standard enthalpy change of hydration?


  • The enthalpy change when one mole of aqueous ions are formed from one mole of gaseous ions

  Na+ (g) + aq --> Na+ (aq)

  Cl- (g) + aq --> Cl- (aq)
• How does the charge of the ion affect lattice enthalpy?

  • The greater the charge on the ion the greater the attraction between oppositely charged ions
  • Therefore lattice enthalpy is more exothermic
• How does the charge of the ion affect the enthalpy of hydration?

  • The greater the charge on the ion the greater the attraction between ions and water molecules
  • Therefore enthalpy of hydration is more exothermic
• How does ionic radius affect the lattice enthalpy / enthalpy of hydration?

  • The smaller the ionic radius the more attraction between oppositely charged ions because more shells = weaker attraction
  • Therefore lattice enthalpy / enthalpy of hydration is more exothermic
• What is entropy?


  • A measure of how dispersed or spread out a system's energy is
  • It is often used as a measure of the disorder or randomness of a system
• Rank solids, liquids and gases in terms of entropy

  1. Gases
  2. Liquids
  3. Solids
• What is the equation for entropy change?

  ΔS = Σ S products - Σ S reactants

  Units - J mol-1 K-1
• What is the feasibility of a reaction?

  • Chemical reactions can only happen if the products of a reaction have a lower overall energy than the reactants
  • Δ G must be negative for a reaction to be feasible (spontaneous)
• What is free energy? ΔG

  Free energy change Δ G is made up of 2 types of energy:

  • Enthalpy change ΔH (KJ mol-1) - heat transfer between system and the surroundings
  • Entropy change at the temperature of the system TΔS - how energy is dispersed within the system itself
• Label the Gibbs equation:

  • Entropy change needs to be converted into KJ by dividing by 1000

  A) KJ mol-1

  B) KJ mol-1

  C) K

  D) J mol-1 K-1

  
• Rearrange the Gibbs equation to find temperature reactions are feasible at:

  • REMEMBER TO DIVIDE Δ S BY 1000 BEFORE DOING CALCULATION

  A) H

  B) S

  
• Suggest why some reactions do not take place despite being thermodynamically feasible e.g. decomposition of hydrogen peroxide

  • Activation energy is too high
  • Non standard conditions
• Feasibility depending on signs:

  A) Negative

  B) All temperatures

  C) Positive

  D) low

  E) low

  F) magnitude

  G) greater

  H) low

  I) high

  J) high

  K) greater

  L) high

  
• What is enthalpy change of neutralisation?

  • Enthalpy change when one mole of water is formed from the reaction between an acid and a base under standard conditions