Reversible reactions
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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

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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.


Equilibrium

Equilibrium

  • 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.


Factors

Factors

  • 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

Catalysts

  • 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.


Concentration

Concentration

  • 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.


Temperature

Temperature

  • 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.


Example

Example

  • 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.


Reversible reactions

  • If we decrease the T, it favours the exothermic reaction – the reverse reaction – and so N2O4 will be produced. A colourless gas!


Pressure

Pressure

  • 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.


Reversible reactions

  • 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)


Water

Water

  • 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)


Reversible reactions

  • 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.


Reversible reactions

  • 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.


Salts

Salts

  • General Rule


Explanation

Explanation!

  • 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!)


Reversible reactions

  • NaCH3COO

  • 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

Soaps

  • 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+

    I I

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

    I I

    CH2 -OCO R** CH2 – OH R** - COO – Na+

    Fat/Oil Glycerol Sodium salts

    Soaps


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