Chapter 6: EquilibriaLong Questions with Answers

The example reaction

If a reversible reaction has attained dynamic equilibrium, it will remain in this state for infinite time unless it is somehow disturbed. We can disturb it on purpose to push it the way we want.

(a) Forward direction

(b) Backward direction

What a catalyst does

A catalyst increases both rates equally: it does not drive the reaction either way; it only makes equilibrium come sooner.

At the start

If a reversible reaction is started by mixing the reactants, the reaction moves in the forward direction only. The concentration of reactants is greatest, so the forward rate is at its highest; there are no products yet, so the backward rate is zero.

As the reaction goes on

Reactants are used up, so the forward rate falls. When enough concentrations of the products are built up, they react to give back the reactants: the backward rate rises from zero.

At equilibrium

The reaction keeps on going in both directions until the rate of the forward reaction becomes equal to the rate of the backward reaction. The number of reactant molecules which disappear in the forward reaction becomes equal to the number formed in the reverse reaction; the same is true for product molecules. From then on the rates stay equal and the concentrations constant.

Dynamic, not stopped

It appears as if nothing is going on, but the reaction has not ceased; it keeps on going in both directions, so this state is called dynamic equilibrium.

Catalyst

A catalyst is a substance that increases the rate of a reaction without itself undergoing any permanent chemical change. It provides a path of lower activation energy.

Effect on the two rates

At dynamic equilibrium the rates of both forward and backward reactions are equal. A catalyst increases both the rates of forward and back reactions of a reversible reaction. It lowers the hump on the energy profile, and the hump is the same hill crossed from either side, so both directions speed up by the same factor.

Effect on time

So if a reversible reaction is carried out in the presence of a catalyst, it will decrease the time taken by the reaction to attain the state of equilibrium. With an iron catalyst at about 400 °C, ammonia formation reaches equilibrium within minutes.

No effect on the position

Why it normally does not

A reversible reaction never goes to completion: the products keep reacting to give back the reactants, and at any time all species are present.

The method

However, it may be forced to go to completion if one or all of the products are withdrawn from the reaction mixture as soon as they are formed, or, being a gas, a product escapes into the atmosphere. With no product to react back, only the forward reaction continues.

Example 1: heating limestone

In an open container, carbon dioxide escapes into the atmosphere as soon as it is formed and the reaction is forced to go to completion. In a closed container, CO₂ reacts with CaO to give back the reactants.

Example 2: Haber process

The ability of ammonia to be converted into its liquid form easily is used to drive the reaction to completion. In this way practically 100 % conversion of N₂ and H₂ to NH₃ is achieved.

Example 3: ethyl acetate

Removing water from C₂H₅OH + CH₃COOH ⇌ CH₃COOC₂H₅ + H₂O as soon as it forms gives a better yield of the ester.

Rule

Every reversible reaction is exothermic in one direction and endothermic, by the same amount, in the other. Raising the temperature favours the endothermic direction; lowering it favours the exothermic direction.

Example 1: ammonia (exothermic forward)

The formation of ammonia is exothermic in the forward direction and endothermic in the reverse direction. The total energy of N₂ and H₂ is higher than that of ammonia. Increasing the temperature at equilibrium pushes the reaction in the backward direction, i.e. to the reactants side. Decreasing the temperature drives the equilibrium forward.

Example 2: PCl₅ (endothermic forward)

Increasing the temperature favours the forward reaction: more PCl₅ decomposes.

Example 3: hydrated salts

Heating blue CuSO₄·5H₂O drives off water (endothermic) to give white CuSO₄; on cooling with moisture, the exothermic reverse change gives the blue salt back.

Speed versus yield

A higher temperature always makes both reactions faster, so equilibrium comes sooner, even when it lowers the yield. That is why ammonia is made at a compromise temperature of about 400 °C with a catalyst.

Observation

When blue copper sulphate pentahydrate is heated strongly, its colour changes to white. When white anhydrous copper sulphate absorbs moisture from the atmosphere, it turns blue again.

Why it is coloured

In the hydrated salt, water molecules are bonded around each Cu²⁺ ion (water of crystallisation). These water molecules change the energies of the copper ion's electrons so that it absorbs red-orange light; the light that is left looks blue.

Why it loses colour

Heating drives off the water molecules. Without them, the Cu²⁺ ion no longer absorbs visible light in the same way, so anhydrous copper sulphate looks white. The crystals also crumble into a powder, because the water held the crystal shape.

Reversibility

Adding water puts the water molecules back around Cu²⁺: the blue colour returns and heat is given out. The same idea explains pink CoCl₂·6H₂O, purple CoCl₂·2H₂O and blue CoCl₂.

The reaction

1. High pressure

4 moles of reacting gases produce 2 moles of product gas. When the pressure is increased, the reaction moves to the side with fewer moles, i.e. forward. The formation of ammonia is thus favoured at high pressure. About 200 atm is used; higher pressure gives more NH₃ but the plant becomes very costly and dangerous.

2. Temperature

The forward reaction is exothermic, so a lower temperature gives a higher yield. But at low temperature the reaction is too slow. About 400 °C is chosen as a compromise between yield and rate.

3. Catalyst

An iron catalyst increases both rates, so equilibrium is reached quickly at this temperature; it does not change the yield.

4. Remove ammonia

The ability of ammonia gas to be converted into its liquid form easily is used to drive the reaction to completion. The mixture is cooled, NH₃ liquefies and is removed, and the unreacted N₂ and H₂ are recycled. In this way practically 100 % conversion of N₂ and H₂ to NH₃ is achieved.

5. Concentration

Keep feeding fresh N₂ and H₂ in the ratio 1 : 3; adding reactant drives the reaction forward.

Irreversible reactions

In a chemical reaction the reactants react to give the products. The reaction continues until all the reactants or one of the reactants is converted into product. NaCl + AgNO₃ → AgCl + NaNO₃ takes place immediately in aqueous solution. The reaction goes to completion, and if stoichiometric amounts are used then no reactant species are present at the end. Such a reaction is called an irreversible reaction. It moves in the forward direction only.

Reversible reactions

In the majority of chemical reactions, however, the reaction does not go to completion. The products react among themselves to give back the reactants under the same conditions. Such a reaction moves in both the forward and backward directions. It is called a reversible reaction and is denoted by ⇌.

One mole of nitrogen reacts with three moles of hydrogen in a closed container to give two moles of ammonia. After its formation, the ammonia decomposes to give the reactants back. The reaction never goes to completion. At any time all three species are present.

Forcing completion

A reversible reaction goes to completion if one of the products is withdrawn or, being a gas, escapes into the atmosphere (CaCO₃ heated in an open vessel).

Physical changes

Like chemical changes, physical changes may also be reversible: copper sulphate pentahydrate turns white on heating and blue again with moisture.

Copper sulphate

Copper sulphate pentahydrate (CuSO₄·5H₂O) is an important salt which is blue in colour. When it is heated strongly, its colour changes to white.

When white anhydrous copper sulphate absorbs moisture from the atmosphere, it turns blue again.

Cobalt chloride

When cobalt chloride hexahydrate (CoCl₂·6H₂O), which is pink, is heated, it is converted into anhydrous CoCl₂, which is blue.

Stepwise rehydration

In the reverse reaction, with limited moisture, anhydrous cobalt chloride is first converted into a dihydrate which is purple. This dihydrate then absorbs four more water molecules to become a hexahydrate, which is pink.

Uses

Both are tests for water: anhydrous CuSO₄ turns blue, cobalt chloride paper turns from blue to pink.

How it is reached

If a reversible reaction is started by mixing the reactants, the reaction moves forward only. After some time, when enough products are built up, they react to give back the reactants. The reaction keeps going in both directions until the rate of forward reaction equals the rate of backward reaction.

Example

N₂ + 3H₂ ⇌ 2NH₃ in a closed vessel at 400 °C with a catalyst: within minutes N₂, H₂ and NH₃ are all present in fixed amounts, while molecules of NH₃ keep forming and decomposing.

Characteristics

Time taken

It depends on the nature of the reaction and the conditions; for very slow reactions it may take years.

At equilibrium

The concentrations of all the participating chemicals are constant at the state of equilibrium.

Adding nitrogen

If we add more N₂ gas, its concentration will increase and the reaction will no longer maintain its state of equilibrium. To restore equilibrium, nitrogen will react with hydrogen to produce more ammonia. This change goes on until a new equilibrium is reached, at which the concentrations will again become constant but different from those of the earlier equilibrium.

Withdrawing ammonia

If some ammonia is withdrawn, its concentration decreases. To restore equilibrium, more nitrogen and hydrogen react to produce ammonia.

Pressure

Change of pressure disturbs the equilibrium of only those gaseous reactions in which the number of moles of reacting gases differs from that of the gases produced. In this reaction 4 moles of reacting gases produce 2 moles of product gas. If the pressure is increased, the reaction moves to the side with fewer moles, i.e. forward. The formation of ammonia is favoured at high pressure.

Why equilibrium matters

Most industrial reactions are reversible. If left alone they give only an equilibrium mixture; by changing the conditions on purpose, chemists push the reaction towards the product and get a high yield quickly.

Haber process (ammonia)

Industrial production of ammonia in the Haber process is a very useful application of chemical equilibrium. Ammonia leads to the important fertilizer urea. High pressure (fewer gas moles), a moderate temperature of about 400 °C, an iron catalyst and liquefying and removing the ammonia give practically 100 % conversion.

Ethyl acetate

Ethyl acetate is commercially very important because it is used as a thinner in the paint industry. One way to get a better yield is to remove water from the reaction mixture as soon as it is formed (concentrated H₂SO₄ or a drying agent).

Lime

CaCO₃ ⇌ CaO + CO₂: heating in an open kiln lets CO₂ escape, giving lime for the steel, glass and paper industries.

Methanation

Coal + steam → CO + H₂ (water gas); CO + 3H₂ ⇌ CH₄ + H₂O over a catalyst: a possible use for Thar coal.