Rates and Equilibria

Rates and Equilibria

1. Buffer Action

   • A substance which resists the change in pH when small amounts of acid or alkali are added.
   • Acidic Buffers
     • Made from a weak acid and its salt (E.g: Ethanoic Acid and Sodium Ethanoate)
   • Alkali Buffers
     • Made from a weak base and its salt. (E.g: Ammonium and Ammonium Chloride)
   • Adding Acid (H⁺)
     • H⁺ reacts with ethanoate to give more ethanoic acid. Excess H⁺ is removed, [H⁺] conc. approx. the same, hence little change in pH.
       • E.g: CH₃COOH ⇋ CH₃COO⁻ + H⁺
   • Adding Alkali (⁻OH)
     • ⁻OH reacts with H⁺, removing it. Ethanoic acid dissociates producing more H⁺. Little change in the overall [H⁺], little change in pH.
   • NOTE: Adding a large amount of acid and alkali will cause the ethanoic acid/ethanoate to be used up, the equilibrium can no longer shift and the extra H⁺/⁻OH cannot be buffered - The pH will begin to change.
   • pH of Buffers.

2. Rates

   • The Rate Equation
     • Tells us how the rate is affected by the conc. of reactants!
     • Rate = k[A]a[B]b
       • [A] = Conc. of A
       • [B] = Conc. of B
       • a/b = The orders of the reactions with respect to reactants A and B
         • a states the order of reaction with respect to A
         • b states the order of reaction with respect to B
       • k = Rate Constant
     • Units: moldm-3s-1
     • Reaction Orders
       • If [A] changes and the rate stays the same, the order of reaction WRT A is 0.
         • Meaning: If the [A] doubles/triples etc. the rate will stay the same!
       • If the rate is proportional to [A], than the order of reaction WRT A is 1
         • Meaning: If the [A] doubles the rate will double. If the [A] triples, so will the rate.
       • If the rate is proportional to [A]^2, than the order of reaction WRT A is 2.
         • Meaning: If the [A] doubles, the rate will be 2^2 (4x faster.) If [A] triples, the rate will be 3^2 (9x faster.)
     • Rate Constant
       • k = Rate Constant
       • The bigger 'k' is, the faster the reaction
       • The rate constant is the same for a certain temperature, but if the temp, is increased, so is 'k'
       • When temp. of a reaction is increased, the rate of reaction increases. The conc. of reactants/orders haven't increased. So it must be an increase in the value of 'k' which has increased the rate of reaction.
   • The Initial Rates Method
     • Finding the Initial Rate.
       • Once the reaction is complete, plot a graph showing Conc. of reactants against Time. With this graph, find the initial rate by drawing a tangent at T=0. Connect the tangent to form a triangle. Label the horizontal line x and the vertical line y.
       • Initial rate = Change in y/Change in x
   • The Rate Determining Step
     • In a reaction with a series of steps, each step can have a different rate. The overall rate is decided by the step with the slowest rate. THE RATE DETERMINING STEP
     • The reactants which appear in the Rate Equation, it must be in the RDS. if it isn't in the Rate Equation, it cannot affect the RDS.
     • IT DOESN'T HAVE TO BE THE FIRST STEP.

3. Acids and Bases

   • pH
     • pH = -Log[H⁺]
     • 2 D.P!!
     • Conc. H⁺ = 10⁻^pH
   • Strong Acid = An Acid Which Fully Dissociates!
     • E.g: HCl → H⁺ + Cl⁻
   • Weak Acid = An Acid Which Only Partially Dissociates
     • E.g: HCOOH ⇋ H⁺ + HCOO⁻
   • Acid Dissociation Constant
     • Ka
     • Ka = [H⁺][A⁻]/[HA]
     • Units: moldm⁻³
     • THIS ONLY APPLIES IF ACID IS DISSOLVED IN WATER!!!!
       • Assumption 1 → [H⁺] = [A⁻]
       • Assumption 2 → [HA] = [HA]
       • Therefore:
         • Ka = [H⁺][A⁻]/[HA]
           • Ka = [H⁺][H⁺]/[HA]
             • Ka = [H⁺]²/[HA]
               • [H⁺] = √Ka[HA]
   • Ionic Product of Water
     • H₂O ⇋ H⁺ + ⁻OH
     • Kw
     • Kw = [H⁺][⁻OH]
     • In Pure Water
       • [H⁺] = [⁻OH]
         • Kw = [H⁺][⁻OH]
           • Kw = [H⁺]²
             • [H⁺] = √Kw
     • In Alkaline Solution
       • Kw = [H⁺][⁻OH]
         • [H⁺] = Kw/[⁻OH]

4. Equilibria

   • Dynamic Equilibrium
     • When there is no net movement of atoms in either direction in the equation
     • Equilibrium is reached when the rate of the forwards reaction = the rate of the backwards reaction. Reactions are still occurring, but the concentration remains constant.
   • Le Chatelier's Principle
     • The position of equilibrium moves as to minimise any change imposed upon it
     • E.g: CH₄+H₂O ⇋ CO + 3H₂
       • Increasing conc. of reactants = Product conc. increases
       • Decreasing conc. of reactants = Product conc. decreases.
     • PRESSURE
       • N₂ + 3H₂ ⇋ 2NH₃ 4 Moles 2 Moles
       • Increasing the pressure will shift the equilibrium to the side with less moles (E.g: Favouring Ammonia Production)
       • Decreasing the pressure will shift the equilibrium to the side with more moles (E.g: Reducing Ammonia Production)
     • TEMPERATURE
       • 2NO₂ ⇋ N₂O₄ (Exothermic)
       • Increasing temperature shifts the equilibrium in the endothermic direction
       • Decreasing temperature shifts the equilibrium in the exothermic direction
   • Equilibrium Constant
     • A constant which tells us where the position of equilibrium lies for a particular reaction
     • Kc>1 = Products Favoured Kc<1 = Reactants Favoured Kc=1 Neither Products Nor Reactants Favoured.
     • Kc = [Products]ⁿ/[Reactants]ⁿ¹
     • E.g: N₂ + 3H₂ ⇋ 2NH₃
       • Kc = [NH₃]²/[N₂][H₂]³

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