GCSE Chemistry Notes: Uses of ammonia and nitric acid - fertilisers
Uses of ammonia, nitric acid, fertiliser salts
Doc Brown's Chemistry KS4 science GCSE/IGCSE/O level Revision Notes
PART D Uses of Ammonia, fertilisers and pollution
The manufacture of nitric acid is described and the preparation of ammonium salts from ammonia and an acid are also explained. The uses of nitric acid, ammonia and ammonium salts are described and some of the pollution problems from overuse of artificial fertilisers. These revision notes on the uses of ammonia, the value of fertilisers, pollution from fertilisers, should prove useful for the new AQA GCSE chemistry, Edexcel GCSE chemistry & OCR GCSE chemistry (Gateway & 21st Century) GCSE (9–1), (9-5) & (5-1) science courses.
Index: A Reversible Reactions * B Reversible reactions and Equilibrium * C The Haber Synthesis of ammonia
D(a) The Uses of ammonia-nitric acid-fertilisers (this page) * D(b) Fertilisers-environmental problems (this page)
E The nitrogen cycle
Advanced A Level Notes - Chemical Equilibrium * Advanced A Level Notes - Nitrogen & Ammonia
PART D 4. The Uses of Ammonia
Ammonia is used to make ammonium salts (mostly artificial fertilisers) and is used in the chemical industry in the manufacture of explosives, nitric acid and nitrates, pharmaceutical products and plastics.
(a) Ammonia is used to manufacture nitric acid
Ammonia is oxidised with oxygen from air using a hot platinum catalyst to form nitrogen monoxide and water.
4NH3(g) + 5O2(g) ====> 4NO(g) + 6H2O(g)
The gas is cooled and reacted with more oxygen to form nitrogen dioxide.
2NO(g) + O2(g) ====> 2NO2(g)
This is reacted with more oxygen and water to form nitric acid.
4NO2(g)+ O2(g) + 2H2O(l) ====> 4HNO3(aq)
Nitric acid is used to make nitro-aromatic compounds from which dyes are made.
It is also used in the manufacture of artificial nitrogenous fertilisers (like ammonium nitrate, see below).
See also Haber synthesis of ammonia
(b) Ammonia is used to manufacture 'artificial' nitrogenous fertilisers
The production and uses of NPK fertilisers are an important sector of the chemical industry. Compounds of the three elements nitrogen, phosphorus and potassium, essential for plant growth, are mixed together and used as 'artificial' fertilisers to improve agricultural productivity, though not without some environmental concerns. The industrial production of NPK fertilisers can be achieved using a variety of raw materials to produce formulations of various salts containing appropriate percentages of the elements.
Ammonia can be used to manufacture ammonium salts and nitric acid. In industry this produced in batches using large vats in which the alkaline ammonia is neutralised with the appropriate acid. The reaction is quite exothermic and the heat released is used to help evaporate and concentrate the solution. The solution is further heated to evaporate water and allow the fertiliser salts to crystallise out of solution. I have described simple methods of preparing ammonium salts in the school/college laboratory.
Soluble potassium chloride (for potassium), potassium sulfate (source of potassium and sulfur) and phosphate rock (source of phosphorus) are obtained by mining. However, phosphate rock cannot be used directly as a fertiliser because it is insoluble and can't be taken up by plants as nutrient.
When phosphate rock is treated with nitric acid you can produce phosphoric acid and the salt calcium nitrate. Phosphoric acid can then be neutralised with ammonia to produce the salt ammonium phosphate.
If phosphate rock is treated with sulfuric acid you make calcium sulfate and calcium phosphate and the mixture is known as 'single super phosphate'.
If phosphate rock is treated with phosphoric acid you produce triple superphosphate (calcium phosphate).
Ammonia is a pungent smelling alkaline gas that is very soluble in water.
The gas or solution turns litmus or universal indicator blue because it is a soluble weak base or weak alkali and is neutralised by acids to form ammonium salts.
More on theory on the Introduction to ionic acid bases theory page or on the Extra Theoretical Acid-Base Chemistry page.
Ammonia is a synthetic rich source of artificial nitrogenous fertilisers essential for increased growth of plants e.g. cereal crops.
Ammonium salts are used as 'artificial' or 'synthesised' fertilisers i.e. nitrogenous fertilisers 'man-made' in a chemical works, and used as an alternative to natural manure or compost etc.
Ammonia is a base, and fertiliser salts are made by neutralising ammonia solution with the appropriate acid.
The method for preparing fertiliser salts in the school/college laboratory are given in the APPENDIX, but the equations are quoted below.
The resulting solution is heated, evaporating the water to crystallise the salt e.g. with correct equations, with and without state symbols ...
(i) ammonia + hydrochloric acid ==> ammonium chloride
NH3 + HCl ====> NH4Cl
NH3(aq) + HCl(aq) ====> NH4Cl(aq) correct equations
(ii) ammonia + sulphuric acid ==> ammonium sulfate
2NH3 + H2SO4 ====> (NH4)2SO4
2NH3(aq) + H2SO4(aq) ====> (NH4)2SO4(aq) correct equations
Before you can make this salt for use in artificial fertilisers the chemical industry must produce, from the appropriate raw materials, large quantities of ammonia by the Haber process and sulfuric acid from the Contact process. So there are several stages in the manufacture of ammonium sulfate.
(iii) ammonia + nitric acid ====> ammonium nitrate
NH3 + HNO3 ====> NH4NO3
NH3(aq) + HNO3(aq) ====> NH4NO3(aq) correct equations
Before you can make this salt for use in artificial fertilisers the chemical industry must produce, from the appropriate raw materials, large quantities of ammonia by the Haber process and nitric acid from the oxidation of ammonia. So there are several stages in the manufacture of ammonium nitrate.
Reactions (ii) and (iii) are used in fertiliser production, as is reaction (iv)
ammonia + phosphoric acid ====> ammonium phosphate
Some of these equations are sometimes written in terms of the fictitious 'ammonium hydroxide' (shown below). The above equations are however, more correct! Quite simply, we are dealing with an aqueous solution of ammonia NH3(aq), but NH4OH is used in some textbooks! Only about 2% of the dissolved ammonia forms ammonium and hydroxide ions (more on this on Theory and Weak and Strong Acids). Please remember these are not strictly the correct equations!
ammonium hydroxide + sulphuric acid ==> ammonium sulphate + water
2NH4OH(aq) + H2SO4(aq) ==> (NH4)2SO4(aq) + 2H2O(l) incorrect equation
ammonium hydroxide + nitric acid ==> ammonium nitrate + water
NH4OH(aq) + HNO3(aq) ==> NH4NO3(aq) + H2O(l) incorrect equation
ammonia + hydrochloric acid ==> ammonium chloride
NH4OH(aq) + HCl(aq) ==> NH4Cl(aq) + H2O(l) incorrect equation
The salt Ammonium chloride is used in zinc-carbon dry cell batteries. The slightly acid paste, made from the salt, slowly reacts with the zinc to provide the electrical energy from the chemical reaction.
If ammonium salts are mixed with sodium hydroxide solution, free ammonia is formed (detected by smell and damp red litmus turning blue).
e.g. ammonium chloride + sodium hydroxide ==> sodium chloride + water + ammonia
NH4Cl + NaOH ==> NaCl + H2O + NH3
NPK Fertilisers
NPK fertilisers contain compounds of three elements essential for the healthy growth of plants, namely nitrogen, phosphorus and potassium. The compounds are mixed together in appropriate proportions and used as 'artificial' fertilisers to improve agricultural productivity, though not without some environmental concerns (see last section).
As already mentioned, ammonia is used to manufacture ammonium salts and nitric acid, and both are involved in producing NPK fertilisers. Salts such as potassium chloride, potassium sulfate and phosphate rock are obtained by conventional mining, however phosphate rock cannot be used directly as a fertiliser because it is insoluble and difficult for plant roots to absorb.
Therefore phosphate rock is reacted with nitric acid to can produce phosphoric acid and the salt calcium nitrate. The phosphoric acid can then be neutralised with ammonia to produce the salt ammonium phosphate.
Phosphate rock can also be reacted with sulfuric acid to produce single superphosphate (a mix of calcium phosphate and calcium sulfate) or reacted with phosphoric acid to produce triple superphosphate (calcium phosphate).
See Contact Process for the manufacture of sulfuric acid
From the above descriptions you can see that a variety of NPK fertilisers can be produced for the agricultural industry and the different formulations can be matched to the particular needs of a farmer's particular crop.
Ammonium sulphate or nitrate salts are widely used as 'artificial or synthetic fertilisers (preparation reactions above). There are several advantages to using artificial fertilisers in the absence of sufficient manure-silage etc. e.g. relatively cheap mass production, easily used to make poor soils fertile or quickly enrich multi-cropped fields.
Artificial fertilisers are important to agriculture and used on fields to increase crop yields but they should be applied in a balanced manner (see next section).
Fertilisers usually contain compounds of three essential elements required for healthy and productive plant growth to increase crop yield.
They replace nutrient minerals used by a previous crop or enrich poor soils so more nitrogen gets converted into plant protein.
Without sufficient of all three of these vital elements plant growth is restricted leading to weak plants and low crop yields.
The fertiliser must be soluble in water so that plants can take up the nutrients through their roots.
However, you don't want it to dissolve too fast or pollution problems will quickly arise and much of the fertiliser will be wasted as well as being an environmental nuisance (see next section).
The fertiliser can be applied in pellet form that slowly breaks down allowing the fertiliser to be absorbed into the soil.
The three important elements in most fertilisers are ...
Nitrogen (N) e.g. from ammonium or nitrate salts like ammonium sulphate, ammonium sulphate or ammonium phosphate (e.g. look for the N in the formula of ammonium salts) or urea, formula CO(NH2)2.
Phosphorus (P) e.g. from potassium phosphate or ammonium phosphate.
Potassium (K) e.g. from potassium phosphate or potassium sulphate.
The fertiliser is marked with an 'NPK' value, i.e. the nitrogen : phosphorus : potassium ratio
Different NPK formulations are made up to suite a particular soil or crop and often applied in slowly dissolving pellet form.
Whatever the particular 'NPK' compounds used, the fertiliser components must be soluble in water to be taken in by plant roots.
PART D contd. 5. Problems with using 'artificial' fertilisers
Overuse of ammonia fertilisers on fields can cause major environmental problems as well as being uneconomic.
Ammonium salts are water soluble and get washed into the groundwater, rivers and streams by rain contaminating them with ammonium ions and nitrate ions.
This contamination causes several problems.
Excess fertilisers in streams and rivers cause eutrophication.
Overuse of fertilisers results in appreciable amounts of them dissolving in rain water and running off into streams, lakes and rivers.
This increases levels of nitrate or phosphate in rivers and lakes, i.e. it significantly increases the levels of plant nutrients in the water.
Naturally occurring green algae living in the water can then feed on these excessive nutrients.
This causes an 'algal bloom' i.e. too much rapid growth of water plants like algae on the surface where the sunlight is the strongest.
This prevents light from reaching plants lower in the water and reduces and eventually stops photosynthesis by the shaded plants on the bed of the lake or river.
Decomposer bacteria can then feed on the dead plants low in the water.
But these lower plants decay via aerobic bacteria which use up any dissolved oxygen.
This means any microorganisms or higher life forms relying on oxygen cannot respire and die.
All the eco-cycles are affected so not only fish, but other respiring aquatic animals like insects die too.
The river or stream becomes 'dead' below the surface as all the food webs are disrupted i.e. lack of plant food for smaller animals, lack of insects for fish etc.
That sums up eutrophication, i.e. what happens in water systems that cannot sustain photosynthesising plants!
Nitrates are potentially carcinogenic (cancer or tumour forming).
The presence in drinking water is a health hazard.
Rivers and lakes can be used as initial sources for domestic water supply.
You cannot easily remove the nitrate from the water, it costs too much!
So levels of nitrate are carefully monitored in our water supply.
More on water pollution on the Extra Aqueous Chemistry page and acid rain on Oil Products page.
There are issues and controversies over the use of artificial fertilisers and it isn't just about eutrophication e.g.
There are great pressures to use mass manufactured 'agro-chemicals' to increase food production to feed the World's growing population.
There are concerns about the quality of soils repeatedly fed artificial fertilisers rather than organic fertilisers like good old fashioned 'muck' from animal waste.
Small scale preparation of an ammonium salt
Preparation of artificial fertilisers by neutralisation
The preparation procedure involves titrating ammonia (pipetted) with standardised hydrochloric, nitric or sulfuric acid (in burette) using methyl orange indicator
The apparatus, chemicals and indicator colours are illustrated in the diagram on the left and another further down, also on the left, for the laboratory preparation of an ammonium salt e.g. ammonium chloride (not used as a fertiliser), ammonium sulfate and ammonium nitrate (both used in synthetic fertilisers). On the bottom right is a diagram of all sorts of apparatus you might come across.
Initially the burette is clamped carefully in position and filled with acid solution - hydrochloric acid, nitric acid or sulfuric acid, depending on which ammonium salt you wish to make. The acid is run through until the reading below the meniscus is 0.00 cm3 (the reading in the diagram is 7.00 cm3, which could represent a titration value). The burette is usually calibrated to 50.00 cm3 (only 10.00 cm3 in diagram - couldn't fit rest of scale on!). The titration is between a weak base (ammonia) and a strong acid (sulfuric, hydrochloric or nitric acid).
The ammonia solution is accurately measured out into the conical flask with e.g. a 25 cm3 pipette and suction bulb (see diagram further down). Add a few drops of methyl orange indicator to the ammonia solution and it should turn yellow for an alkali. Carefully place the conical flask under the tip of the burette so drops don't go astray!
The titration: You carefully add small portions of the acid, swirling after each addition and checking the colour of the indicator (not shown in the diagram, but its good to stand the flask on white tile). At the start of the titration the methyl orange indicator is yellow. As you add the acid you get 'splurges' of reddish-orange colour until the mixture is swirled in the conical flask. Try to add dropwise when you seem to be near the orange colour at the endpoint.
The end-point is an orange colour, that is when all the ammonia is neutralised by which ever acid you are using.
If you 'overshoot' the titration with excess acid, the methyl orange indicator turns red.
The whole procedure of (1) to (3) is repeated without the methyl orange indicator, that is you measure out the same amount of ammonia solution (e.g. 25 cm3 from the pipette) and add to it the titration volume of acid from the end-point first time round (3a). The colourless neutralised solution is then transferred to an evaporating dish. The solution is gently heated to evaporate some of the water.
On leaving to cool the crystals of the ammonium salt should form - crystallisation. The residual liquid can be decanted away and the crystals can be carefully collected and dried by 'dabbing' with a filter paper OR the crystals can be collected by filtration left to dry.
The whole procedure (1) to (5) is 'briefly' illustrated in the left-hand diagram below.
(i) You can make potassium nitrate this way, but you would use phenolphthalein indicator.
Potassium nitrate is also used in fertilisers.
to make it in the school/college laboratory you would put potassium hydroxide in the conical flask and titrate it with nitric acid solution.
The end-point is when the pink colour of the phenolphthalein indicator is discharged and the solution is colourless.
potassium hydroxide + nitric acid ==> potassium nitrate + water
KOH + HNO3 ==> KNO3 + H2O
KOH(aq) + HNO3(aq) ==> KNO3(aq) + H2O(l)
(ii) ?
The equations have already been given in section D(b)
A variety of apparatus you might come across
See also GCSE/IGCSE Acids & Alkalis revision notes sub–index: Index of all pH, Acids, Alkalis, Salts Notes 1. Examples of everyday acids, alkalis, salts, pH of solution, hazard warning signs : 2. pH scale, indicators, ionic theory of acids–alkali neutralisation : 4. Reactions of acids with metals/oxides/hydroxides/carbonates, neutralisation reactions : 5. Reactions of bases–alkalis like ammonia & sodium hydroxide : 6. Four methods of making salts : 7. Changes in pH in a neutralisation, choice and use of indicators : 8. Important formulae of compounds, salt solubility and water of crystallisation : 10. More on Acid–Base Theory and Weak and Strong Acids AND How to do acid-alkali titration calculations, diagrams of apparatus, details of procedures
Index: A Reversible Reactions * B Reversible reactions and Equilibrium
C The Haber Synthesis of ammonia * D(a) The Uses of ammonia-nitric acid-fertilisers
D(b) Fertilisers-environmental problems * E The nitrogen cycle
Foundation tier (easier) multiple choice QUIZ on ammonia, nitric acid and fertilisers etc.
Higher tier (harder) multiple choice QUIZ on ammonia, nitric acid and fertilisers etc.
Advanced A Level Notes - Equilibrium (use indexes)
Advanced A Level Chemistry Notes p-block nitrogen & ammonia
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