AH Chem Notes - Unit 2 Physical

• Chemical equilibrium

  Many reactions never go to completion but are in a state of equilibrium with reactants and products always present
• Equilibrium

  A chemical reaction is said to be in equilibrium when the composition of the reactants and products remains constant indefinitely
• Closed system


  In order for equilibrium to be established, the reaction must take place in a closed system
• Dynamic equilibrium

  In a closed system, both forward and backward reactions go at the same rate
• Le Chatelier's principle

  An equilibrium system always changes to reduce the effect of any outside change made on it
• Effects of external changes on equilibrium position

  • Increase reactant concentration - shifts towards products
  • Increase temperature - shifts in endothermic direction
  • Increase pressure - shifts in direction of smaller gas volume
  • Add catalyst - no effect
• Equilibrium constant


  Characterises the composition of the reaction mixture at equilibrium
• Homogeneous equilibria

  • Fe3+(aq) + 6CN-(aq) ⇌ [Fe(CN)6]3-(aq)
  • 2NO2(g) ⇌ N2O4(g)
• Heterogeneous equilibria

  • CH3COOCH3(aq) + H2O(l) ⇌ CH3COOH(aq) + CH3OH(aq)
  • 3Fe(s) + 4H2O(g) ⇌ Fe3O4(s) + 4H2(g)
• Equilibrium constant value

  • For large values, the position of equilibrium lies well to the right (products)
  • For small values, the position of equilibrium lies well to the left (reactants)
• Changing concentration or pressure does not affect the equilibrium constant
• Catalysts have no effect on the position of equilibrium or the equilibrium constant
• Temperature increase

  • Increases equilibrium constant for endothermic reactions
  • Decreases equilibrium constant for exothermic reactions
• Water equilibrium

  Equilibrium between water molecules and hydrogen and hydroxide ions
• In pure water at 25°C, [H3O+] = [OH-] = 1 x 10-7 mol l-1
• Ionic product of water (Kw)

  Kw = [H3O+][OH-] = 1.01 x 10-14 at 25°C
• Kw increases with temperature

  Indicates water dissociation is endothermic
• pH

  • Acidic solutions have pH < 7, [H+] > [OH-]
  • Alkaline/basic solutions have pH > 7, [H+] < [OH-]
  • Neutral solutions have pH = 7, [H+] = [OH-]
• pH + pOH = 14
• pH calculations

  • pH of 0.01 M HCl = 2
  • pH of 0.001 M NaOH = 11
  • pH of 0.2 M HCl = 0.7
  • pH of 0.5 M NaOH = 13.7
  • pH of solution with [H+] = 2.51 x 10-9 M = 8.6
• Strong acid

  Completely dissociated into ions in aqueous solution
• Weak acid

  Only partially dissociated into ions in aqueous solution
• Weak acid

  • Ethanoic acid
• Comparing strong and weak acids

  1. Measure pH, conductivity, reactions with alkali, carbonates and metals

  2. Hydrochloric acid (strong) has lower pH, faster reactions than ethanoic acid (weak)
• Procedure

  1. Count the number of drops used

  2. Add sodium carbonate to the tube containing hydrochloric acid and indicator

  3. Add sodium hydroxide solution (0.4 mol l-1) to ethanoic acid and hydrochloric acid

  4. Place 2 cm depth of ethanoic acid solution (1.0 mol l-1) in a new test-tube

  5. Place 2 cm depth of hydrochloric acid (1.0 mol l-1) in another new tube

  6. Add a small piece of magnesium ribbon to the hydrochloric acid tube

  7. Repeat with the ethanoic acid tube and compare the rate of reaction

  8. Measure the conductivity of 1 mol l-1 hydrochloric acid

  9. Repeat with 1 mol l-1 ethanoic acid
• The pH is neutral
• Hydrochloric acid (strong) has a lower pH than ethanoic acid (weak)
• Hydrochloric acid (strong) has a faster rate of reaction with metals and metal carbonates than ethanoic acid (weak)
• The volumes of bases required for neutralising hydrochloric acid and ethanoic acid were the same
• In weak acids, the position of the equilibrium lies to the left so there is a very low concentration of hydrogen ions compared to the concentration of acid molecules
• The value of the equilibrium constant, K, for the dissociation of weak acids is less than 1
• All carboxylic acids are weak acids
• Dissociation of methanoic and propanoic acids

  Equations to show the dissociation
• Acids such as sulfuric, carbonic and sulfurous acid are known as diprotic acids because they contain two hydrogen atoms per molecule
• Ionisation of diprotic acids takes place in two steps
• Strong base

  Completely ionised in aqueous solution
• Strong bases

  • Sodium hydroxide
  • Oxides and hydroxides of alkali metals
  • Soluble oxides and hydroxides of Group 2 metals
• Weak base

  Only partially ionised in solution
• Weak bases

  • Ammonia
  • Amines
• When salts dissolve in water they become fully ionised