Introducing Reversible Reactions

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Reversibility and Dynamic Equilibria:

Example #1:

Blue copper (II) sulphate crystals heated gently turn to white powder

Heating --> crystals lose water of crystallisation --> form white anhydrous copper (II) sulphate

If water added to white solid --> turns blue and gets warm

Therefore, reaction reversed.

CuSO4 . 5H2O(s) --> CuSO4(s) +5H2O(l)

CuSO4(s) +5H2O(l) --> CuSO4 . 5H2O(s)

Example #2:

Heating ammonium chloride, white crystals disappear from bottom of tube -- reappear further up

Splits into ammonia and hydrogen chloride gases (both colourless)

Combine further up the tube (cooler) --> reaction reverses when conditions changed from hot to cool

NH4Cl(s) --> NH3(g) + HCl(g)

NH3(g) + HCl(g) --> NH4Cl(s)

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Reversible Reactions Under ‘Closed’ Conditions:

Closed - no substances added to reaction mixture, none escape. However, heat can be given off or absorbed.

Start with high concentration of blue squares

Beginning, rate of reaction of blue turning yellow higher

Rate of change falls as blue used up

Yellow can also turn blue since reversible reaction

Beginning, rate of change to blue zero

As yellow increases, rate of change to blue higher

At equilibrium, rate of both reactions equal - blue to yellow and yellow to blue at same rate

Dynamic equilibrium:

Dynamic - reactions still continuing

Equilibrium - total amount of substances constant

Only set up if substance closed

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Writing Equations for Reactions in Dynamic Equilibrium:

A + B ⇌ C + D

Forward reaction: between A and B

Back reaction: between C and D

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Manipulating Reversible Reactions:

To produce substances, e.g. C, as efficiently as possible, reversible reactions, where it turned back into A and B wouldn't be helpful.

To produce it, alter position of equilibrium.

Position of equilibrium: proportions of various things in equilibrium mixture.

E.G. if high proportion of C and D, position of equilibrium lies towards C and D or to the right.

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Le Chatelier’s Principle:

If a dynamic equilibrium is disturbed by changing the conditions, the reaction moves to counteract the change.

Conditions which could be changed:

Increase/decrease concentrations of substances present

Change pressure

Change temperature

Add catalyst (has no effect on position of equilibrium)

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Adding and Removing Substances:

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Changing Pressure:

Only for gas reactions and where total number of molecules on both sides different.

A(g) + 2B(g) ⇌ C(g) + D(g)

Pressure - molecules hitting walls of container

Increase pressure -- reaction responds by reducing it -- producing fewer molecules to hit walls of container -- creating more C and D

Increasing pressure makes reaction go in direction producing fewer molecules

Same number of molecules -- no effect on position of equilibrium

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Changing Temperature:

A + 2B ⇌ C + D ΔH = -1000 kJ/mol

Forward reaction exothermic, so back reaction endothermic by same amount

Decreasing temperature -- reaction responds to increase temperature -- produce more C and D

C and D produce more heat due to exothermic change -- warms reaction mixture up

Increasing temperature -- reaction will get rid of extra heat -- absorb in endothermic change -- back reaction favoured

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Adding Catalyst:

Speeds up reactions

Reversible change - speeds up forward and back reactions by same proportion

Net effect - no change in position of equilibrium but increases rate at which equilibrium reached

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Example Illustrating Le Chatelier’s Principle:

NO2 - dark brown poisonous gas

Joins up in pairs (dimerise) -- forms molecules of dinitrogen tetroxide N204 - colourless

2NO2(g) ⇌ N2O4 (g) ΔH = -57 kJ/mol

Pressure -

Increase -- equilibrium moves to reduce -- produce fewer molecules -- more N2O4 produced

Decrease -- equilibrium moves to increase -- produce more molecules - more 2NO2 produced

Temperature -

Forward reaction exothermic, so back reaction endothermic

Decrease -- equilibrium moves to produce more heat -- produces more N2O4 -- colour fades

Increase - equilibrium moves to produce less heat -- produces more 2NO2 -- colour darkens