Paper 3

• Metallic lattice


  Giant metallic lattice of metal cations and delocalized electrons
• Ionic lattice


  Giant ionic lattice of oppositely charged ions
• Covalent lattice


  Lattice of simple molecules with weak intermolecular forces
• Giant covalent lattice


  Lattice of covalently bonded atoms or SiO2
• Melting point


  • Metallic - High
  • Ionic - High
  • Simple covalent - Low
  • Giant covalent - High
• Electrical conductivity in solid state

  • Metallic - Yes
  • Ionic - No
  • Simple covalent - No
  • Giant covalent - No (unless graphite or graphene)
• Electrical conductivity in liquid/aqueous

  • Metallic - Yes
  • Ionic - Yes
  • Simple covalent - No
  • Giant covalent - No (unless graphite or graphene)
• Induced dipole (London) forces


  Exist between all molecules, stronger as number of electrons increases
• Permanent dipole-dipole forces


  Exist between polar molecules
• Hydrogen bonds


  Special type of permanent dipole-dipole, exist between polar molecules with H bonded to F, O, or N
• Relative strength of intermolecular forces

  • Hydrogen bonds - Strongest
  • Permanent dipole-dipole
  • Induced dipole (London)
• Determining molecular shape and bond angles


  1. Count electron regions (including lone pairs)

  2. Determine number of bonding regions and lone pairs

  3. Consider repulsion between regions

  4. Apply rule of thumb for lone pair effects

  5. Determine associated shape and angle
• Periodic trends


  • Atomic radius decreases across a period
  • First ionization energy generally increases across a period
  • Melting points reach a maximum at group 4
  • Simple covalent molecules after group 4
• Preparing a standard solution

  1. Calculate moles needed

  2. Calculate mass of chemical needed

  3. Weigh out chemical and dissolve in small amount of water

  4. Transfer to volumetric flask and make up to mark with distilled water

  5. Shake to mix
• Redox titration procedure


  1. Dissolve solid/solution in water and make up to volume

  2. Take 25 cm3 aliquot and place in conical flask

  3. Titrate with standard solution of known concentration

  4. Use mole ratio in redox equation to calculate moles in flask

  5. Relate moles in flask to original dissolved amount
• Redox cell setup


  One beaker with platinum electrode in I3+/I2+ solution, other beaker with zinc rod in Zn2+ solution, connected by salt bridge
• Increasing concentration of reactant


  Equilibrium shifts right, solution appears darker
• Decreasing concentration of reactant


  Equilibrium shifts left, solution appears lighter
• Lowering temperature


  Favors exothermic forward reaction
• Equilibrium constant (KP)


  Ratio of partial pressures of products to reactants, higher value means equilibrium further to right
• Buffer formula


  pH = pKa + log([salt]/[acid])
• Blood buffer system


  Carbonic acid and bicarbonate ions, maintains pH of 7.4
• Calorimetry calculation


  1. Use q = mcΔT to calculate energy transferred

  2. For exothermic, ΔH = -q/n

  3. For endothermic, ΔH = q/n
• Enthalpy change of neutralization for strong acid-base is around -57 kJ/mol, less for weak acid-base
• If you're losing heat to the surroundings you may not be carrying out in standard conditions
• The heat capacity of the container is not included in the calculation
• If it gets quite hot then the contents of the calorimeter might evaporate
• Enthalpy change of neutralization

  The enthalpy change for the reaction between an aqueous acid and alkali to form a mole of water
• The enthalpy change of neutralization for all strong acid-alkali combinations is around -57 kJ per mole
• This is because they are fully ionized and give the same ionic equation
• If you have a weak acid-alkali combination, it will be less exothermic than -57 kJ per mole
• This is because they are only partially ionized, so some of the energy of the reaction is used to ionize the acid
• Enthalpy change of solution

  The enthalpy change when one mole of a substance dissolves in water
• Determining the enthalpy change of solution by calorimetry

  1. Know the mass of the solid being dissolved

  2. Know the mass of the solution formed

  3. Measure the temperature change of the solution

  4. Calculate the energy transferred to or absorbed by the solution using q = mcΔT

  5. Convert the Q value to kJ

  6. Calculate the moles of substance dissolved using mass/Mr

  7. Enthalpy change of solution = Q/moles dissolved (include a sign, negative for exothermic)
• Born-Haber Cycle


  Involves lattice enthalpies and hydration enthalpies of the two ions, and includes the enthalpy change of the solution process
• Calculating enthalpy changes using Hess's Law

  1. If given ΔHf values, draw a formation cycle

  2. If given ΔHc values, draw a combustion cycle

  3. Use the summation formula: ΔHR = ΣΔHp - ΣΔHr

  4. Order of steps: CRAP for ΔHc, PAR for ΔHf

  5. Downward arrows for exothermic, upward for endothermic

  6. Include state symbols and electrons

  7. Beware bond enthalpy vs enthalpy of atomization
• The bond enthalpy is for breaking one mole of bonds, while the enthalpy of atomization is for half that process
• If there's a 2 in front in the cycle, double the ΔH value
• Be careful with double negatives when rearranging the Hess's Law equation
• Factors affecting lattice enthalpy

  Ionic charge and ionic radius (not atomic radius)