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Reversible Reactions. A + B <=> C + D In a reversible reaction as soon as some of the products are formed they react together, in the reverse reaction, to form the reactant particles. Example As soon as A + B react forming C + D, some C+ D react together to produce A + B. Equilibrium.

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reversible reactions
Reversible Reactions
  • A + B<=>C + D
  • In a reversible reaction as soon as some of the products are formed they react together, in the reverse reaction, to form the reactant particles.
  • Example
  • As soon as A + B react forming C + D, some C+ D react together to produce A + B.
  • In a reversible reaction the forward and backward reactions occur at the same time.
  • Therefore the reaction mixture will contain some reactant and product particles.
  • When the rate of the forward reaction is equal to the rate of the reverse reaction – we say they are at EQUILLIBRIUM.
  • Dynamic Equilibrium is when the conditions are balanced and the reaction appears to have stopped.
  • We can alter the position of equilibrium by changing:
  • The concentration of reactants or products.
  • Changing the temperature.
  • Changing the pressure ( in gas mixtures only)
le chatelier s principle
Le Chatelier’s Principle
  • If a system is at dynamic equilibrium and is subjected to a change- the system will offset itself to the imposed change.
  • This is only true when a reversible reaction has reached equilibrium.
  • Catalysts will lower the activation energy of the forward and reverse reaction by the same rate.
  • A catalyst increase the rate of the reaction but has no effect on equilibrium position.
  • A+B <=> C + D
  • If we add more A or B we speed up the forward reaction and so more C and D are produced. Equilibrium shifts to RHS
  • If reduce the amount of C and D – then more A and B will react producing more C and D. Equilibrium shifts to RHS
  • If we add add more C or D then the reverse reaction will happen – more A and B will be produced. The same will happen if remove some A or B. In both cases equilibrium shifts to LHS.
  • In a reversible reaction – one will be exothermic and the other will be endothermic.
  • A rise in temperature favours the reaction which absorbs heat – the endothermic reaction.
  • A drop in temperature favours the reaction that releases heat – the exothermic reaction.
  • N2O4 (g)<=> NO2 (g) ΔH = +
  • (clear) (brown)
  • NO2 is formed when most metal nitrates decompose or when you add Cu to HNO3.
  • NO2 is a dark brown gas.
  • The forward reaction is endothermic.
  • If we increase the T, it favours the endothermic reaction and so equilibrium will shift to the RHS. We will see a dark brown gas.
If we decrease the T, it favours the exothermic reaction – the reverse reaction – and so N2O4 will be produced. A colourless gas!
  • Changing the pressure will only affect a gaseous mixture.
  • An increase in P will cause equilibrium position to shift to the side with the least amount of gaseous molecules.
  • 2 SO2 (g) + 1O2 (g) <=> 2 SO3 (g)
  • 3 moles of gas <=> 2 mole of gas
  • If we increase P – the equilibrium will move to the RHS since there are fewer gas molecules.
N2O4 (g) <=> 2 NO2 (g)
  • (clear) (brown)
  • I mole of gas <=> 2 moles of gas
  • If we increase the P – equilibrium will move to the LHS since there are fewer gas molecules. We will see the brown colour vanish.
  • If we decrease the P – equilibrium will shift to RHS – more gas molecules – we will see the brown NO2.
catalysts and equilibrium
Catalysts and Equilibrium
  • A catalyst lowers EA and so speeds up reaction rate.
  • In a reversible reaction it lowers the EA for the forward and reverse reaction by the same amount.
  • Therefore they speed up the rate of both reactions by the same amount.
  • They have no effect on equilibrium position -but a system will reach equilibrium faster.
equilibrium in industry
Equilibrium in Industry
  • The Haber Process
  • Manufacture of NH3
  • N2(g) +3H2(g)<=>2NH3(g) ΔH=-92kJ
  • The forward reaction is exothermic. Therefore a low T will move equilibrium to the RHS. ( If T is too low reaction will be slow)
  • Increasing P will favour equilibrium to shift to the RHS since fewer gas molecules on that side. ( 4moles – 2 moles)
  • Conditions actually used = 200 atmospheres (P), T = 380 – 400 o C. In continuous processor.
  • NH3 is condensed – un reacted N2 and H2 recycled.
acids and bases
Acids and Bases
  • The pH scale is a measure of the concentration of Hydrogen ions.
  • The pH stands for the negative logarithm:
  • pH = - log10 [H+(aq)]

([ ] = concentration)

  • The pH scale is continuous – (below 1 and above 14)
  • An equilibrium exists with water
  • H2O (l)<=> H+(aq) + OH– (aq)
  • The concentration of both H+ and OH- are 10 –7 moles l-1.
  • [H+] = [OH-]= 10 –7 mol/l
  • [H+] [OH-] = 10 –7 x 10 –7
          • = 10 – 14 mol2 l -2
calculating concentration
Calculating concentration
  • [H+] = 10 –14 / [OH-]
  • [OH-] = 10 –14 / [H+]
  • Example
  • Calculate the concentration of OH- ions is a solution contains 0.01 moles of H+
  • [OH-] = 10 –14 / [H+]
  • = 10 –14/ 10 –2 ( 0.01 = 10 –2)
  • = 10 –12 mol/l.
more examples
More examples
  • Calculate the pH of a solution that contains 0.1 moles of OH- ions.
  • [H+] = 10 –14 / 10 –1

= 10 –13 mol/l

pH = - log10 [H+]

= - log 10 –13

= 13

strong weak acids
Strong/Weak Acids
  • A strong acid is one where all the molecules have dissociated (changed into ions)
  • Example
  • HCl(g) + (aq) —> H+ (aq) + Cl- (aq)
  • (molecules) ( ions)
  • Other strong acids – Sulphuric, Nitric, phosphoric.
weak acids
Weak Acids
  • These are acids that have only partially dissociated ( ionised) in water.
  • Example – carboxylic acids, carbonic acid, sulphurous acid.
  • The majority of the particles lie at the molecule side of the equilibrium.
  • CH3COOH (aq) <=> CH3COO- (aq)

(molecules) + H+ (aq) ( ions)

Strong and weak acids differ in:
  • Conductivity, pH and reaction rate.
  • If comparing we must use equimolar solutions I.e. both same mol/1.
strong weak bases
Strong/Weak Bases
  • Strong base – completely dissociated.
  • Example
  • NaOH(s) + (aq) <=> Na+(aq)+OH-(aq)
  • Other examples – alkali metals.
  • Weak bases are partially dissociated.
  • Example
  • NH3(aq) + H2O <=> NH4+ (aq)+ OH-(aq)
affect on equilibrium
Affect on equilibrium
  • If we add Sodium ethanoate to Ethanoic Acid –
  • CH3COOH(aq) <=>CH3COO-(aq) + H+(aq)
  • NaCH3COO(s)+(aq) <=>Na+(aq)+CH3COO- (aq)
  • We have increased the concentration of the ethanoate ions (in the system) – equilibrium will shift to the LHS to offset this. Therefore there will be less H+ ions and so pH will rise.
What happens to equilibrium position if we add NH4Cl to NH4OH?
  • NH4OH(aq) <=> NH4+ (aq) + OH- (aq)
  • NH4Cl (s) => NH4+ (aq) + Cl- (aq)
  • The number of NH4+(aq) ions is increasing on the RHS of the system, equilibrium will shift to the LHS to offset this. The will be fewer OH- (aq) ions and so the pH will decrease.
  • General Rule
  • NH4Cl
  • This is the salt of a weak alkali

( NH4OH) and a strong acid ( HCl).

  • When we add it to water:
  • NH4Cl(s) + (aq) <=> NH4+(aq) + Cl-(aq)
  • H2O (l) <=> H+ (aq) + OH-(aq)
  • The NH4+ ions and the OH- ions in the system react
  • NH4+(aq) + OH-(aq) <=> NH3 (aq) + H2O(l)
  • The concentration of OH- ions in the water equilibrium goes down – the equilibrium shifts to the RHS to offset this – producing more H+ ions and so pH goes down.(acidic!)
  • This is the salt of astrong alkali

( NaOH) and a weak acid (CH3COOH).

  • When we add it to water:
  • NaCH3COO(s) + (aq) <=> CH3COO-(aq) + H+ (aq)
  • H2O (l) <=> H+ (aq) + OH-(aq)
  • The CH3COO(aq) reacts with the H+ (aq) ion.
  • CH3COO-(aq) + H+(aq) <=> CH3COOH(aq)
  • The water equilibrium then moves to RHS to offset this – there are now more OH-(aq) ions and so the pH will increase
  • ( alkaline!)
  • Soaps are formed when we hydrolyse fats and oils using an alkali.
  • They are the salts of weak acids and strong bases – ph of soaps will be slightly alkaline.

CH2 – OCO R CH2 –OH R – COO-Na+


CH - OCO R* <=> CH – OH + R* - COO – Na+


CH2 -OCO R** CH2 – OH R** - COO – Na+

Fat/Oil Glycerol Sodium salts