Module 3

• Group 2


  Alkali earth metals, outer flections
• Group 2 elements


  • Outer electrons s²₂ in 3-subshell
  • Redox reactions are most common, each metal atom oxidised losing valence electrons
  • Group 2 is reducing agent
  • Reactivity increases down group - atomic radius increase, more shielding, less attraction to nucleus, easier to remove outer electrons
  • Locations form more easily
• Reactions of Group 2 elements

  1. 2M(s) + O₂(g) → 2MO(s)

  2. 2M(s) + 2H₂O(l) → 2M(OH)₂(s) + H₂(g)

  3. M(s) + 2HCl(aq) → MCl₂(aq) + H₂(g)

  4. M(s) + H₂SO₄(aq) → MSO₄(aq) + H₂(g)
• Solubility of Group 2 hydroxides

  Increases down the group
• Ionisation energies of Group 2

  Decrease down the group - attraction between nucleus and outer electrons decrease, easier to lose
• Group 2 oxides with water

  Release OH- to form alkaline solution
• Solubility of Group 2 hydroxides

  Increases down the group, more soluble when solution saturated form precipitate
• Group 2 compounds


  • Calcium oxide + water → Calcium hydroxide
  • Magnesium hydroxide - partially soluble, used in suspension to neutralise excess stomach acid and treat constipation
• Halogens


  • Boiling point increases down the group, simple molecular structures with weak London forces
  • Reactivity increases up the group, more attraction to nucleus, easier to gain electron
  • Oxidising power decreases down the group, atomic radius increase, more shielding, less nuclear attraction to attract electron
• Halogen displacement reaction


  More reactive halogen displaces less reactive from an aqueous solution of its halide
• Appearance of halogens in solution and organic solvent

  • Cl₂ - pale green, Br₂ - yellow, I₂ - brown in solution

  Cl₂ - colourless, Br₂ - yellow, I₂ - purple in organic solvent
• Disproportionation reaction - same species is oxidised and reduced, e.g. chlorine with cold, dilute aqueous alkali
• Uses of chlorine

  Makes water drinkable, sterilises water, bleaches, universal indicator
• Reaction of chlorine with cold, dilute sodium hydroxide

  Cl₂(aq) + 2NaOH(aq) → NaCl(aq) + NaClO(aq) + H₂O(l)
• Tests for halide ions, carbonates, sulfates, ammonium ions

  Halide ions - dissolve in nitric acid, add silver nitrate solution

  Carbonates - add dilute HCl, then sodium carbonate, effervescence

  Sulfates - acidify sample with dilute HCl, add barium chloride, white precipitate

  Ammonium ions - react with warm NaOH, release ammonia gas
• Periodicity

  Row of elements = group, column of elements = period

  Properties change gradually across a period, similar properties in same group
• Ionisation energy


  Energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous ions

  Increases across a period, decreases down a group
• Metallic structure

  Atoms highly packed in lattice with delocalised electrons, metal ions repel and are neatly arranged
• Giant covalent structures


  Bonds continue indefinitely, lattice forms, covalent bonding exists between all adjacent atoms

  Diamond - tetrahedral, Graphite - hexagonal layers with delocalised electrons between
• Enthalpy change


  Total chemical energy inside a substance, exothermic - products have less energy than reactants, endothermic - products have more energy than reactants
• Activation energy


  Minimum amount of energy needed for reactant molecules to successfully collide and start reaction
• Standard enthalpy changes


  Enthalpy change under standard conditions - 100 kPa, 298 K

  Standard enthalpy of reaction, formation, combustion, neutralisation
• Calorimetry


  Measuring enthalpy changes, q = m x c x ΔT
• Collision


  Particles need to collide with correct orientation and enough energy to successfully react
• ΔH


  Overall energy given in / given out to surroundings
• Enthalpy changes


  • Standard enthalpy of reaction
  • Enthalpy of combustion
  • Enthalpy of formation
  • Enthalpy of neutralisation
• Standard conditions


  Pressure - 100kPa, Temperature - 298K (25°C)
• Standard enthalpy of reaction

  ΔH when reactants in the equation react to give products under standard conditions (exothermic/endothermic)
• Enthalpy of combustion

  ΔH when one mole of a substance is burnt in excess oxygen under standard conditions (exothermic)
• Enthalpy of formation


  ΔH when one mole of a compound is formed from its elements under standard conditions
• Enthalpy of neutralisation


  ΔH when one mole of water is formed by reacting an acid and alkali under standard conditions
• Calorimetry


  1. Measure ΔH, use polystyrene cup/vacuum flask

  2. q = m x c x ΔT (heat transferred, mass, specific heat capacity, temperature change)
• Specific heat capacity


  Energy needed to increase the temperature of 1g of a substance by 1°C
• Calculating ΔH/n (moles)

  Use q = m x c x ΔT, consider water as largely water so use mass of water
• Bond enthalpy


  Energy required to break one mole of a specific covalent bond in the gas phase
• Average bond energy


  Average energy of the same type of bond but in different environments, found by entropy cycles
• Calculating ΔHr


  ΔHr = Σ(bond enthalpies of bonds broken) - Σ(bond enthalpies of bonds formed)
• Using ΔHf (arrows up)

  ΔHr = Σ(ΔHf of products) - Σ(ΔHf of reactants)
• Using ΔHc (arrows down)

  ΔHr = Σ(ΔHc of reactants) - Σ(ΔHc of products)
• Activation energy (Ea)


  Energy needed for particles to react, larger for endothermic reactions