reversible reactions and dynamic equilibrium

are chemical reactions reversible

• yes - the sign is ⇌

• the direction of some reversible reactions can be altered by changing the reaction conditions

• forward reaction = left to right

• backward reaction = right to left

dynamic equilibrium

forward and backward reactions are occurring at the same rate, so the amount of reactant and product doesn’t change

position of equilibrium

‘lies to the left’ = more reactant than product at equilibrium

‘lies to the right’ = more product than reactant at equilibrium

how can equilibrium be affected

• if there is a change is the surroundings, equilibrium moves to counteract the change

• e.g. if the temperature increases, the opposing action will be to decrease temperature and shift towards the endothermic side of the reaction

• if pressure increases, the opposing action will be to decrease pressure and shift to the side with less gas molecules

• same works with increasing reactant/product concentration

how can rate of attainment of equilibrium be affected

catalysts increase speed at which equilibrium is reached (but not equilibrium position)

changing temperature/pressure/concentration can also shift the position of equilibrium and so change how long it takes to achieve it

haber process

reversible reaction (so it can obtain equilibrium) between nitrogen and hydrogen to form ammonia

3H2 + N2 —> 2NH3

nitrogen = extracted from air

hydrogen = obtained from natural gas

conditions for haber process

• iron catalyst

• temperature 450 degrees C

• pressure 200 atmospheres

explain temperature for haber process

ideal = lower temperature would favour forward reaction as it is exothermic so the equilibrium would shift to the right to increase temperature, increasing the yield of ammonia

however, this would mean that the rate of reaction is very slow at low temperatures

450 degrees C is a compromise; reduces yield but increases rate of reaction

explain pressure of haber process

ideal = high pressure would favour forward reaction as the equilibrium would shift to the right to decrease pressure as there are less gas molecules, increasing the yield of ammonia

however high pressures can be very dangerous and expensive to maintain, buy equipment for, and contain

so 200 atm is a compromise between a lower yield of products being made more safely and economically

fertilisers

may contain nitrogen (found in soluble ions such as ammonium and nitrate), phosphorus, and potassium compounds to promote plant growth

ammonia being used for fertiliser:

ammonia + nitric acid —> ammonium nitrate

making ammonium sulfate in the laboratory

preparing ammonium sulfate by titration:

• add exact volume of ammonia to conical flask placed on the white tile

• add a few drops of indicator and swirl - should turn yellow

• add sulfuric acid to flask solution drop by drop, swirling the falsk in between

• continue until the colour turns red sharply and record the titre

• repeat by adding exactly the same amount of acid but this time without the indicator (is an impurity(

• pour reaction mixture in an evaporating dish, gently heat in water bath to remove some water

• leave in dry place so remaining water evaporates, allowing crystallisation to occur

• after a few days ammonium sulfate crystals should appear

industrial preparation of ammonium sulfate

• large scale operation

• hugely expensive and complex

• ammonia prepared by haber process, sulfuric acid by contact process

comparing industrial and laboratory productions of ammonium sulfate

• in laboratory more simple equipment needed, industrial is hugely expensive and complex

• industrial has more stages

• in laboratory lower concentrations, industrial high concentrations

• laboratory less heat given off, industrial highly exothermic

• in laboratory product separated by crystallisation which is a very slow process, industrial heat produced is used to evaporate water from the reaction mixture to make very concentrated ammonium sulfate

• industrial faster