Uses of ammonia, nitric
acid, fertiliser salts
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
A Reversible Reactions * B Reversible reactions and Equilibrium
Haber Synthesis of ammonia
D(a) The Uses of ammonia-nitric acid-fertilisers
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
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
is oxidised with oxygen from air using a
catalyst to form nitrogen monoxide
- 4NH3(g) +
5O2(g) ====> 4NO(g) + 6H2O(g)
- The gas is cooled and reacted with more oxygen to form
- This is reacted with more oxygen and water to form
(b) Ammonia is used to manufacture 'artificial' nitrogenous fertilisers
Raw materials the composition
and use of fertilisers
The production and uses of NPK
fertilisers are an important sector of the chemical industry.
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
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.
acid can then be neutralised with ammonia to produce the salt ammonium
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).
To make ammonium sulfate from
ammonia and sulfuric acid, an important component of many NPK
fertilisers you need quite a few raw materials:
(i) The sulfuric acid is made
Contact Process, which uses sulfur, oxygen
(from air) and water.
(ii) The ammonia is made by
Haber process, which
uses nitrogen from air and hydrogen from hydrocarbons in oil.
It takes a lot of energy and raw materials
to manufacture nitrogen based inorganic fertilisers.
Ammonium sulfate and ammonium nitrate are
synthetic - artificial fertilisers.
Composted plant or animal waste is an
Ammonia gas and ammonium salts -
chemistry of making fertilisers
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
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
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 ...
ammonia + hydrochloric acid
==> ammonium chloride (not used in fertilisers)
NH3 + HCl ====> NH4Cl
NH3(aq) + HCl(aq) ====> NH4Cl(aq)
ammonia + sulphuric acid
ammonium sulfate (used in fertilisers)
2NH3 + H2SO4 ====>
2NH3(aq) + H2SO4(aq)
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.
ammonia + nitric acid
ammonium nitrate (used in fertilisers)
NH3 + HNO3 ====>
NH3(aq) + HNO3(aq)
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
ammonium hydroxide + nitric acid ==>
ammonium nitrate + water
ammonia + hydrochloric acid ==>
ammonium compounds (apart from fertilisers - in next section)
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).
More on NPK
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 - bigger crop yields, 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
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).
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.
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
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).
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
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 ...
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.
e.g. from potassium phosphate or ammonium phosphate.
e.g. from potassium phosphate or potassium sulphate.
The fertiliser is
marked with an
'NPK' value, i.e. the
nitrogen : phosphorus : potassium
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.
As you can see below, the most
modern fertilisers contain a lot more elements than the NPK
Modern fertilisers supply a whole
range of elements that plants need to grow healthily and give high
The fertiliser bag information
indicates the following elements are present in its composition:
nitrogen N, phosphorus P,
potassium K, magnesium Mg, sodium Na, sulfur S, calcium Ca,
boron B, copper Cu, zinc Zn, manganese Mn, iron Fe, cobalt Co
and selenium Se
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
- Excess fertilisers in streams and rivers cause
- 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
- 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
Aqueous Chemistry page and acid rain on
- There are issues and controversies over
the use of artificial fertilisers and it isn't just about eutrophication
- 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.
preparation of an ammonium salt
of artificial fertilisers by neutralisation
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).
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!
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.
end-point is an orange colour, that is when all the ammonia
is neutralised by which ever acid you are using.
you 'overshoot' the titration with excess acid, the methyl
orange indicator turns red.
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.
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)
is 'briefly' illustrated in the right-hand diagram below.
(i) You can make potassium
nitrate this way, but you would use phenolphthalein
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)
equations have already
been given in section D(b)
(iii) This is a small scale school
In the chemical industry the salt
will be done in several ways e.g.
(i) automatically by mixing the
acid and alkali (ammonia) solutions in large quantities
in big reaction vessel.
The neutralised salt solution
will be continuously fed into an evaporation and
(ii) in making the fertiliser
ingredient ammonium sulfate, ammonia gas straight
from the Haber process is mixed with a spray of sulfuric
acid and the salt is precipitated as a white powder.
The sulfuric acid is made by the
A variety of
apparatus you might come across
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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.
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 :
More on Acid–Base Theory and Weak and Strong Acids
How to do acid-alkali titration calculations, diagrams of apparatus,
details of procedures
A Reversible Reactions * B
Reversible reactions and Equilibrium
Haber Synthesis of ammonia * D(a) The Uses of ammonia-nitric acid-fertilisers
Fertilisers-environmental problems (this page)
* 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
Advanced A Level Chemistry Notes p-block nitrogen & ammonia
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chemistry Keywords & formulae: NH3 O2 NO HNO3 NO2 4NH3 + 5O2
==> 4NO + 6H2O * 4NO2+ O2 + 2H2O ==> 4HNO3 * 2NH3 + H2SO4
==> (NH4)2SO4 * NH3 + HNO3 ==> NH4NO3 * NH3 + HCl ==> NH4Cl
+ NaOH ==> NaCl + H2O + NH3 * 2NH4OH + H2SO4 ==> (NH4)2SO4 +
2H2O * NH4OH + HNO3 ==> NH4NO3 + H2O * NH4OH + HCl ==> NH4Cl
+ H2O * 4NH3 + 5O2 ==> 4NO + 6H2O * 4NO2+ O2 + 2H2O ==>
4HNO3 * 2NH3 + H2SO4 ==> (NH4)2SO4 * NH3 + HNO3 ==> NH4NO3 *
NH3 + HCl ==> NH4Cl + NaOH
==> NaCl + H2O + NH3 * 2NH4OH + H2SO4 ==> (NH4)2SO4 + 2H2O * NH4OH +
HNO3 ==> NH4NO3 + H2O * NH4OH + HCl ==> NH4Cl + H2O GCSE chemistry IGCSE chemistry revision
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fertiliser? why are ammonium salts good artificial
fertilisers? what is an NPK fertiliser?
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