Thermodynamics

Definitions
1. ΔHf
  1. The Standard Molar Enthalpy of Formation, ΔHfƟ is the enthalpy change when one mole of compound is formed from its constituent elements under standard conditions, all reactants and products in their standard states.
  2. H2(g) + 0.5O2(g) → H2O(l) ΔHfƟ = -286 kJ mol-1
2. ΔHa
  1. The Standard Enthalpy of Atomisation, ΔHaƟ is the enthalpy change which accompanies the formation of one mole of gaseous atoms from the element in its standard state under standard conditions.
  2. 0.5Cl2(g) → Cl(g) ΔHaƟ = +121.7 kJ mol-1
3. 1st IE
  1. The First Ionisation Energy (first IE) is the standard enthalpy change when one mole of gaseous atoms is converted into a mole of gaseous ions each with a single positive charge.
  2. Na(g) → Na+(g) + e- First IE = +496 kJ mol-1
4. 2nd IE
  1. The Second Ionisation Energy (second IE) is the standard enthalpy change when one mole of gaseous ions with a single positive charge is converted into a mole of gaseous ions each with a 2+ charge.
  2. O-(g) + e- → O2-(g) Second EA = +798 kJ mol-1
5. 1st EA
  1. The First Electron Affinity (EA), ΔHeaƟ is the standard enthalpy change when a mole of gaseous atoms is converted to a mole of gaseous ions, each with a single negative charge.
  2. O(g) + e- → O-(g) Frist EA = -141.1 kJ mol-1
6. 2nd EA
  1. The Second Electron Affinity (EA), ΔHeaƟ is the enthalpy change when a mole of electrons is added to a mole of gaseous ions each with a single negative charge to form ions each with two negative charges.
  2. O-(g) + e- → O2-(g) Second EA = +798 kJ mol-1
7. ΔHLf
  1. The Lattice Formation Enthalpy, ΔHLfƟ is the standard enthalpy change when one mole of solid ionic compound is formed from its gaseous ions.
  2. NaCl(s) → Na+(g) + Cl-(g) ΔHLfƟ = -788 kJ mol-1
8. ΔHLd
  1. The Enthalpy of Lattice Dissociation, ΔHLdƟ is the standard enthalpy change when one mole of solid ionic compound separates into its gaseous ions.
  2. NaCl(s) → Na+(g) + Cl-(g) ΔHLdƟ = +788 kJ mol-1
9. ΔHhyd
  1. The Enthalpy of Hydration, ΔHhydƟ is the standard enthalpy change when water molecules surround one mole of gaseous ions.
  2. Na+ + aq → Na+(aq) ΔHhydƟ = -406 kJ mol-1
10. ΔHsol
  1. The Enthalpy of Solution, ΔHsolƟ is the standard enthalpy change when one mole of solute dissolves completely in sufficient solvent to form a solution in which the molecules or ions are far enough apart not to interact with each other.
  2. NaCl(s) + aq → Na+(aq) + Cl-(aq) ΔHsolƟ = 19 kJ mol-1
  3. Dissolving an ionic compound in a solvent is the sum of these three processes:
    1. 1. the lattice enthalpy has to be put in
    2. 2. the enthalpy of hydration of the positive ions is given out
    3. 3. the enthalpy of hydration of the negative ions is given out
11. Mean Bond Enthalpy
  1. Mean Bond Enthalpy is the enthalpy change when one mole of gaseous molecules each breaks a covalent bond to form two free radicals, averaged over a range of compounds.
  2. CH4(g) → C(g) + 4H(g) ΔHLdƟ = +1664 kJ mol-1
    1. C-H Bond Enthalpy in Methane
      1. 1664/4= +416 kJ mol^(−1)
  3. Using Mean Bond Enthalpies
    1. Bond BREAKING takes energy IN
    2. Bond FORMING releases energy OUT
    3. ΔH = Total Energy In - Total Energy Out
12. Standard Enthalpy of Bond Dissociation
  1. The Standard Enthalpy of Bond Dissociation is the enthalpy change when a mole of gaseous molecules each breaks a covalent bond to form two free radicals.
13. An Enthalpy Change is a heat change at constant pressure
Born-Haber Cycles
1. When drawing Born-Haber Cycles:
  1. 1. Make a rough scale
    1. 2. Plan it out roughly so you don't go off the paper
      1. 3. Remember to put in the sign of each enthalpy change, and their direction
        +ve up
        -ve down
2. Sodium Chloride
3. Magnesium Chloride
4. Stages of drawing a cycle:
  1. 1st. Elements in their standard states
    1. 2nd. Atomise both the atoms
      1. 3rd. Ionise the positive ion
        4th. Ionise the negative ion
        5th. Add in the enthalpy of formation of the ionic compound
Trends in Lattice Enthalpies
1. Lattice Formation Energies of Simple Ionic Compounds
  2. Lattices of larger ions have smaller formation enthalpies, as larger ions don't approach each other so closely, therefore less energy is released when the lattice is formed.
2. Lattice formation enthalpies for compound with larger charges
  2. The same trend can be seen in this table
  3. Comparing the two tables shows that lattice formation enthalpy increases with the size of the charge on the ions.
Entropy
1. It is the randomness of the system
2. This drives chemical processes, as a system has a tendency towards randomising
3. The general order of randomness from least to most random is: Solids < Liquids < Gases
4. ΔS
5. Quoted for standard conditions of 298K and 100kPa
6. The larger the entropy, the greater the randomness
7. The units of Entropy are J K-1 mol-1
  1. REMEMBER! Must be converted to kJ K-1 mol-1 for use in calculations
8. Entropies increase with temperature.
9. A positive ΔS means that the products are more random than the reactants
10. ΔS = S products - S reactants
Gibbs Free Energy
1. Used to predict the feasibility of a reaction
2. Combination of enthalpy and entropy change
3. ΔG
5. ΔG = ΔH - TΔS
6. T is given in Kelvins
7. ΔS may be in J mol-1 K-1 and not kJ mol-1 K-1
8. ΔG varies with temperature
9. When ΔG = 0, the reaction is said to be just feasible, and is in equilibrium, so both products and reactants are present
10. ΔG does not take rate of reaction into account, so a reaction may be thermodynamically feasible at a specific temperature, but may not happen when carried out practically
  1. It is said to have a large activation energy barrier for the reaction
  2. It is also said to be thermodynamically unstable, but kinetically stable

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