Equilibrium is reached when the forward and reverse reactions, of a reversible reaction, occur at the same rate (the reaction must be in a closed system).
When the forward and reverse rates balance, and are equal, it can appear as if the reaction isn't doing anything - and is finished. What actually is happening is that the reactants are being converted into products at the same rate the products are converted back into reactants.
The concentrations of the reactants vs products are not necessarily equal, and there is usually more of one than the other:
if the concentration of the reactants is higher than products, we say the equilibrium lies to the left
if the concentration of the products is higher than the reactants, we say the equilibrium lies to the right
Dynamic equilibrium is when the forward and reverse reactions are still occurring, however, they are occurring at a different (changed) rate.
If a system is at equilibrium and a change is made to any of the conditions (temperature, concentration, pressure), then that system will respond to counteract the change. This is called Le Chatelier's Principle.
Changing the concentration, temperature and pressure of a reaction system can make a big change to where the equilibrium lies, and industry uses this principle regularly to increase the amount of product they make (for the best profits!).
if you add more reactant, the equilibrium will shift to the right to reduce the concentration of reactant (and make more product)
if you remove some of the product, the equilibrium will shift to the right to increase the concentration of the product
if the temperature is increased then the equilibrium position will shift to reduce the temperature (so will favour the endothermic reaction)
if the temperature is decreased then the equilibrium position will shift to increase the temperature (so will favour the exothermic reaction)
if you increase the pressure then the equilibrium will shift to reduce it (by favouring which ever side of the reaction has the fewest molecules of gas)
The Haber Process
nitrogen + hydrogen ⇌ ammonia
N₂ + 3H₂ ⇌ 2NH₃
The Haber Process is used to produce ammonia, which can be used to manufacture nitrogen-based fertilisers. This process is a reversible reaction between nitrogen (extracted from the air) and hydrogen (obtained from natural gas). This reaction can reach a dynamic equilibrium.
The conditions for the Haber Process are:
temperature of 450 °C
pressure of 200 atmospheres (200 atm)
A catalyst will increase the rate of both reactions equally, and so does not change the yield of the reaction. The reasons why the other above conditions are used, is given below:
Temperature of 450 °C
forward reaction is exothermic, so a lower temperature favours the forward reaction to maximise yield
through heating particles will have more energy, so more frequent, successful collisions between particles
a compromise has to be made on the temperature
Pressure of 200 atm
the forward reaction turns 4 moles ('volumes') of gas into 2 moles of gas
increasing the pressure will favour the forward reaction to maximise the yield (as this forces the particles to take up less volume)
to maintain this high pressure, lots of energy is needed - so high costs
Fertilisers are made from compounds of nitrogen, phosphorus and potassium, and are used to improve agricultural productivity. NPK fertilisers contain compounds of all three elements. Industrial production of NPK fertilisers can be achieved using different raw materials in several different steps. NPK fertilisers are formulations of salts containing defined percentages of these elements.
Ammonia is made in the Haber process, and is used to manufacture ammonium salts as well as nitric acid.
Nitric acid is made from ammonia. Several stages are involved, but overall the process is:
ammonia + oxygen → nitric acid + water
NH₃(g) + 2O₂(g) → HNO₃(aq) + H₂O(l)
Ammonium nitrate is a salt used as a fertiliser. It is made by reacting ammonia with nitric acid:
ammonia + nitric acid → ammonium nitrate
NH₃(g) + HNO₃(aq) → NH₄NO₃(aq)
Industrial production of ammonium sulfate
The industrial production of ammonium sulfate happens on a much larger scale than its production in the lab. A fertiliser factory begins with the raw materials needed to make ammonia and sulfuric acid, rather than buying these two reactants from elsewhere.
Several stages are required to produce ammonia and sulfuric acid form their raw materials and the production is carried out on a large scale.
The industrial production of ammonium sulfate is a continuous process. The product is made quickly all the time, as long as raw materials are provided.
Making ammonium sulfate in the lab
Ammonium sulfate can be made in the lab using dilute ammonia solution and dilute
sulfuric acid. Both reactants are soluble, so a titration must be used to carry out this reaction:
ammonia + sulfuric acid → ammonium sulfate
2NH₃(aq) + H₂SO₄(aq) → (NH₄)₂SO₄(aq)
The lab preparation of ammonium sulfate is a batch process. This means that only a small amount of product at any one time, and the apparatus needs to be cleaned before each new batch.
pour dilute sulfuric acid into a beaker
add a few drops of methyl orange indicator
add dilute ammonia solution drop by drop, stirring in between each addition
continue step 3 until the colour permanently changes from red to orange
pour the reaction mixture into an evaporating basin, and heat carefully over a boiling water bath
stop heating before all the water has evaporated. Leave aside for crystals to form
pour away excess water and leave the crystals to dry in a warm oven (or pat dry with filter paper)