Chemistry KS4 science GCSE/IGCSE Revision Notes
Chemical Equilibrium, Ammonia Synthesis by the Haber Process & Nitric acid, their Manufacture and Uses e.g. in
The first two sections introduce the concepts of reversible
reactions and a dynamic chemical equilibrium through detailed explanation and
example descriptions. The rules on the
effect of temperature, pressure and concentration on the position of an
equilibrium are also covered. Section 3. describes the manufacture (synthesis) of ammonia by
the Haber Process from nitrogen and hydrogen and how ammonia is converted to
nitric acid. Sections 4. & 5. describe some important uses of ammonia and nitric
acid are given including the use of ammonium and nitrate salts as artificial
fertilisers and problems with their use. Section 6. outlines the 'Nitrogen
1.Reversible Reactions * 2.Reversible reactions and Equilibrium
Haber Synthesis of ammonia * 4.The Uses of ammonia-nitric acid-fertilisers * 5.Fertilisers-environmental problems
* 6.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 Level Notes - Equilibrium
Advanced Level Chemistry Notes p-block
nitrogen & ammonia
- Introduction examples
- A reversible reaction is a chemical
change in which the products can be converted back to the original reactants under suitable conditions.
- In other words, you can change the
position of the chemical
- A reversible reaction is shown by the sign ,
- a half-arrow to the right (direction
of forward reaction),
- and a half-arrow to the left (direction
of backward reaction).
- Most reactions are not reversible (irreversible)
and have the usual complete arrow
only pointing to the right.
Five examples 1a to 1e of reversible reactions are
1a The thermal decomposition of ammonium
- On heating strongly above 340oC, the white solid ammonium chloride,
thermally decomposes into a mixture of two colourless gases
ammonia and hydrogen chloride.
- On cooling the reaction is reversed and solid ammonium chloride reforms.
- Reversing the reaction conditions reverses the direction of chemical change, typical of a reversible reaction.
- Thermal decomposition means using 'heat' to 'break down' a molecule into smaller ones.
- The decomposition is endothermic (heat absorbed or heat taken in) and the
formation of ammonium chloride is exothermic (heat released or heat given out).
- This means if the direction of chemical change is reversed, the energy change is also reversed.
- Ammonium fluoride (>?oC),
ammonium bromide (>450oC) and
ammonium iodide (>550oC), with a similar formula, all
sublime in a similar physical-chemical way
when heated, so the equations will be similar i.e. just swap F, Br or
I for the Cl.
- Similarly, ammonium sulphate
also sublimes when heated above 235oC and thermally
decomposes into ammonia gas and sulphuric acid vapour.
- Ammonium nitrate does not
undergo a reversible sublimation reaction, it melts and then
decomposes into nitrogen(I) oxide gas (dinitrogen oxide) and water
- This is very different reaction,
in fact its an irreversible redox reaction. The nitrate ion, NO3-,
or any nitric acid formed, HNO3, act as an oxidising agent and
oxidise the ammonium ion. If the products are cooled, ammonium nitrate is
- For more on
sublimation, see the States of Matter
1b The thermal decomposition of hydrated copper(II)
- The 5H2O in the formula of hydrated copper(II) sulphate is called the
water of crystallisation and forms part of the crystal structure when copper(II) sulphate solution is
evaporated and crystals form.
- This crystal structure is broken down on
heating and the water is given off.
- The thermal decomposition is endothermic as
heat is absorbed to drive off the water.
- The reverse reaction is exothermic i.e. on adding water to white anhydrous copper(II) sulphate the mixture heats
up as the blue crystals reform.
- The reverse reaction is used as a simple chemical test for
water i.e. white anhydrous copper(II) sulphate turns blue.
The reaction of bismuth chloride with
chloride (BiCl3) is added to water, it dissolves and then
reacts with water to form a white precipitate of bismuth oxychloride
(BiOCl) and a colourless solution of hydrochloric acid (HCl).
acid is added to the mixture, the bismuth oxychloride dissolves to
reform the bismuth chloride solution.
reaction is reversible and what is formed depends on the relative
amounts of hydrochloric acid and water, and so the reaction equation
should be written as:
The formation and decomposition of
The synthesis of
ammonia from nitrogen and hydrogen is a reversible reaction and by
changing the temperature and pressure of the reacting gases you can
make the reaction go one way more than another,
e.g. high pressure
makes more ammonia (forward reaction) and higher temperature causes ammonia to decompose
into hydrogen and nitrogen (backward reaction).
The formation and hydrolysis of
- Ethyl ethanoate, an ester, is formed by the reaction
of ethanoic acid with ethanol e.g.
- ethanoic acid + ethanol
ethyl ethanoate + water
- Its an equilibrium, and
starting with the pure acid plus pure alcohol you get about 2/3rds
conversion* to the ester, and the reaction is catalysed by a
few drops of concentrated sulphuric acid.
- *This means a yield of about 67% in the
- If the ester is warmed with water or
any dilute acid (faster), it changes back into the original acid and
alcohol. This reverse reaction is called hydrolysis (backward
- ethyl ethanoate + water ==> ethanoic acid + ethanol
- whereas esterification (forward
- ethanoic acid + ethanol ==>
ethyl ethanoate + water
- Esters are usually sweet smelling and widely used as fragrances
(e.g. perfumes) and
2. Reversible reactions and
- When a reversible reaction occurs in a closed system an equilibrium is formed, in which the original reactants and products formed coexist.
- In an equilibrium there is a state of balance between the concentrations of the reactants and products
and once a state of chemical equilibrium is reached there is no further
change in concentrations BUT the reactions don't stop!
- At equilibrium the rate at which the
reactants change into products is exactly equal to the rate at which the products change back to the original reactants.
- However the final relative equilibrium amounts of the reactants and products depend on the reaction conditions
e.g. the temperature and pressure.
For industrial processes, it is important to maximise the concentration of the desired products and minimise the 'leftover' reactants. A set of rules can be used to predict the best reaction conditions to give the highest possible yield of product.
The three rules outlined below
are known as Le Chatelier's Principle. This essentially states
that if a change is imposed on a system, the system will change to
minimise the enforced change to re-establish equilibrium.
Rule 1a: If the forward reaction forming the product is endothermic, raising the temperature favours its formation increasing the yield of product (lowering the temperature decreases the yield).
Rule 1b: If the forward reaction forming the product is exothermic, decreasing the temperature favours its formation (increasing temperature decreases the yield).
Rule 2a: Increasing the pressure favours the side of the equilibrium with the least number of gaseous molecules as shown by the balanced symbol equation.
Rule 2b: Decreasing the pressure favours the side of the equilibrium with the most number of gaseous molecules as shown by the balanced symbol equation.
Rules 1 above, and rule
3, below, apply to any reaction, BUT rule 2 above, ONLY applies to a reaction with one or gaseous reactants or products.
Rule 3a: If the concentration of a
reactant (on the left) is increased, then some of it must change to the products
(on the right) to maintain a balanced equilibrium position.
Rule 3b: If the concentration of a
reactant (on the left) is decreased, then some of the products (on the
right) must change back to reactants to maintain a balanced equilibrium
- e.g. nitrogen + hydrogen
- or N2(g) + 3H2(g)
- If the nitrogen or hydrogen concentration
was increased, some of this extra gas would change to
- If the nitrogen or hydrogen concentration
was decreased, some of ammonia would change back to nitrogen and
- At AS-A2 advanced level things can get more
complicated e.g. can you figure out why in terms of concentration to
maintain the equilibrium balance? (and if a gcse
don't worry if you can't) ...
- So in terms of enforced
change ==> system response:
- Increasing nitrogen ==>
decreases hydrogen and increases ammonia.
- Increasing hydrogen ==>
decreases nitrogen and increases ammonia.
- Increasing ammonia ==>
increases nitrogen and hydrogen.
- Decreasing ammonia ==>
decreases nitrogen and hydrogen.
- Decreasing nitrogen ==>
increases hydrogen and decreases ammonia.
- Decreasing hydrogen ==>
increases nitrogen and decreases ammonia.
Rule 4: A catalyst does NOT affect the
position of an equilibrium, you just get there faster! A catalyst usually
speeds up both the forward and reverse reaction but there is no way it can
influence the final 'balanced' concentrations. However, the importance of a
catalyst lies with economics e.g. (i) bringing about reactions with high
activation energies at lower temperatures and so saving energy or (ii) saving
time is saving money.
Notes on Chemical Equilibrium
Applying the rules 1 to 3
(a) The formation of calcium oxide (lime) and carbon dioxide from calcium carbonate (limestone)
CaO(s) + CO2(g)
The forward reaction is endothermic, 178kJ of heat energy is absorbed (taken in) for every mole of calcium oxide formed.
One mole of gas is formed in the process, so there is a net increase in the moles of gas in lime formation, since there are no gaseous reactants.
From rule 1: increasing the temperature will increase the yield of
calcium oxide or lime, CaO which is endothermically formed.
From rule 2: decreasing the pressure will favour the formation of more
gas molecules if possible, so more carbon dioxide formed, and hence more lime.
Lime is made commercially by heating limestone to a high temperature
(e.g. 1000oC) in a limekiln that is well ventilated (this reduces the carbon dioxide
pressure and so reduces the un-desired backward reaction).
(b) The formation of hydrogen
chloride from hydrogen and chlorine
H2(g) + Cl2(g)
The forward reaction is exothermic, 184kJ of heat energy is given out in forming hydrogen
bromide according to the above equation (184/2 = 92kJ per mole of HCl formed).
There is no net change in the moles of gas (2 moles reactants
2 moles of product)
From rule 1: decreasing the temperature favours the
exothermic formation of hydrogen chloride, so the equilibrium moves
proportionately to the right-hand side (more HCl, less H2 or Cl2).
From rule 2: since there is no net change in the
number of moles of gas on reaction, pressure has no effect on the yield of hydrogen
chloride and the proportions of HCl, H2 or Cl2 stay the
(c) The formation of ammonia (see section 3. below)
Synthesis of ammonia - The Haber Process
- Ammonia gas is synthesised in the chemical industry by
reacting nitrogen gas with hydrogen gas in what is known as the
Haber-Bosch Process, named after two highly inventive and subsequently
- The nitrogen is obtained from
liquified air (80% N2).
Air is cooled and compressed under high pressure to form liquid air (liquefaction).
The liquid air is fractionally distilled at low temperature to separate
oxygen (used in welding, hospitals etc.), nitrogen (for making ammonia),
Noble Gases e.g. argon for light bulbs, helium for balloons).
- The hydrogen is made by reacting methane (natural
gas) and water or from cracking hydrocarbons (both reactions are done at
high temperature with a catalyst).
- CH4 + H2O ==> 3H2
- e.g. C8H18 ==> C8H16
- The synthesis equation for the Haber
Synthesis reversible reaction is ...
N2(g) + 3H2(g)
- .. which means an equilibrium
will form, so there is no chance of 100% yield even if you use, as you
actually do, the theoretical reactant ratio of nitrogen : hydrogen of 1 : 3 !
- In forming ammonia 92kJ of heat energy is given out
exothermic, 46kJ of heat released per mole of ammonia formed).
- Four moles of 'reactant' gas form two moles of 'product' gas, so there is net decrease in gas molecules on forming ammonia.
- So applying the equilibrium rules from section 2
above, the formation of ammonia is favoured by ...
- (a) Using high pressure because you are
going from 4 to 2 gas molecules (the high pressure also speeds up the
reaction because it effectively increases the concentration of the gas
molecules), but higher pressure means more dangerous and more costly
- (b) Carrying out the reaction at a
low temperature, because it is an
exothermic reaction favoured by low temperature, but this may produce too
slow a rate of reaction,
- So, the idea is to use a set of optimum conditions to get the
most efficient yield of
ammonia and this involves getting a low % yield (e.g. 8% conversion)
- Described below are the conditions to give the most economic
production of ammonia.
- these arguments make the point that
yield* of an equilibrium reaction depends on
the conditions used.
The word 'yield' means how much
product you get compared to the theoretical maximum possible if the
reaction goes 100%.
- For more on chemical economics see
Industrial Chemistry page.
- In industry pressures of 200 - 300 times normal atmospheric pressure are used
in line with the theory.
- Theoretically a low temperature would give a high yield of ammonia BUT ...
is very stable molecule and not very
reactive i.e. chemically inert, so the rate of reaction is too slow
at low temperatures.
- To speed up the reaction
an iron catalyst is used as well as a
higher temperature (e.g.
- The higher temperature is an economic compromise,
i.e. it is more economic to get a low yield fast, than a high yield slowly!
- Note: a catalyst does NOT affect the yield of a
reaction, i.e. the equilibrium position BUT you do get there faster!
reference to the HABER SYNTHESIS chemical plant DIAGRAM
- Hydrogen and nitrogen gases are mixed
in the ratio 3:1 (to fit in with equation) and the gaseous mixture fed
into the top of the reaction chamber.
- The gases are pumped down through the
reaction chamber filled with lots of beds ('shelves') coated in the iron
- The hydrogen and nitrogen gases react on the
surface of the iron catalyst to form ammonia.
N2(g) + 3H2(g)
- At the end of the process, when the
gases emerge from the bottom of the iron catalyst reaction chamber, the gas mixture is cooled under high pressure, when
only the ammonia liquefies and is so can be removed, tapped off from
the cooled compression chamber and stored in cylinders for use e.g. making
- Any unreacted nitrogen and hydrogen (NOT liquified),
is recycled back through the reactor chamber, nothing is wasted!
- Nitrogen (-196oC) and hydrogen (-252oC) have much lower boiling
points than ammonia (-33oC) and stay as gases.
- Boiling points increase with pressure,
but these normal atmospheric pressure values offer a fair comparison and the
higher the boiling point of the liquid, the higher condensation point of
- The temperature in the lower chamber is
never low enough to condense out the unreacted hydrogen or nitrogen so
only the desired product, ammonia gas condenses out, then the liquid ammonia is drained off at the
bottom of the.
- Since the hydrogen and nitrogen are
still gases above the liquid ammonia, they are easily pumped around and
mixed with new hydrogen and nitrogen and hence recycled through the
- This means non of the original
hydrogen and nitrogen reactants is wasted, despite the reaction
being an equilibrium.
- In fact the yield of ammonia can be
as little as 6% conversion, but FAST, and the other 94% of reactant
gases is recycled FAST.
- To sum up: A low % yield of ammonia is produced quickly at
moderately high temperatures and pressure in the presence of an iron catalyst, and is more economic than getting a higher %
equilibrium yield of ammonia at a more costly high pressure and a slower lower
- Detailed notes on "Rates
of Reaction" for further reading.
- AND there are some more general notes on
Economics on the Industrial Chemistry page.
4. The Uses of Ammonia
4(a) Ammonia is used to manufacture nitric acid
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
- 4NO2(g)+ O2(g)
- 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).
4(b) Ammonia is used to manufacture 'artificial' nitrogenous 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 salts. (more on theory on the Acids,
Bases Salts page or on the Extra
Theoretical Acid-Base Chemistry page).
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.
The fertiliser salts are made by neutralising ammonia solution with the appropriate
acid (more method details on Acids, Bases and
Salts, but the equations are given below).
The resulting solution is heated,
evaporating the water to crystallise the salt e.g.
ammonia + sulphuric acid
2NH3(aq) + H2SO4(aq)
ammonia + nitric acid
NH3(aq) + HNO3(aq)
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 Extra
Aqueous Chemistry). Please remember these
are not strictly the correct equations!
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).
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
'problems' in section 5.).
contain compounds of three essential elements for healthy and
productive plant growth to increase crop yield. They replace
nutrient minerals used by a previous crop or enriches poor soil and more nitrogen
gets converted into plant protein.
ammonium or nitrate salts like ammonium sulphate, ammonium sulphate
or ammonium phosphate or urea (e.g. look for the N in the formula of
e.g. from potassium phosphate or ammonium phosphate
e.g. from potassium phosphate, potassium sulphate.
The fertiliser is
marked with an
'NPK' value, i.e. the
nitrogen : phosphorus : potassium
Fertilisers must be
soluble in water to be taken in by plant roots.
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.
- This increases levels of nitrate or
phosphate in rivers and lakes.
- This causes 'algal bloom' i.e. too much rapid growth of water plants on the surface
where the sunlight is the strongest.
- This prevents light from reaching plants lower in the water.
- These lower plants decay and the active
aerobic bacteria use up any dissolved oxygen.
- This means any microorganisms or higher
life forms relying on oxygen cannot respire.
- All the eco-cycles are affected and fish and other respiring aquatic animals die.
- The river or stream becomes 'dead' below the surface as all the food webs are disrupted.
- 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 acid rain on Oil
The Nitrogen Cycle
the gaseous element N2(g)
- Nitrogen is an extremely
important element for all plant or animal life! It is found in
important molecules such as amino acids, which are combined to
form proteins. Protein is used everywhere in living organisms from muscle
structure in animals to enzymes in plants/animals.
- Nitrogen from the
- Nitrifying bacteria,
e.g. in the root nodules of certain plants like peas/beans (the
legumes), can directly convert atmospheric nitrogen into nitrogen
compounds in plants e.g. nitrogen => ammonia => nitrates
which plants can absorb.
- However, most plants can't do
this conversion from nitrogen => ammonia, though they can all
absorb nitrates, so the 'conversion' or 'fixing' ability might be introduced into other
plant species by genetic
- The nitrogen from air is
converted into ammonia in the chemical industry, and from this
artificial fertilisers are manufactured to add to nutrient
deficient soils. However, some of the fertiliser is washed out of
the soil and can cause pollution (see above).
- The energy of lightning
causes nitrogen and oxygen to combine and form nitrogen oxides which
dissolve in rain that falls on the soil adding to its nitrogen
- N2(g) + O2(g)
==> 2NO(g), then
- then 2NO(g)
+ O2(g) ==> 2NO2(g)
- NO2(g) +
water ==> nitrates(aq) in rain/soil
- Incidentally, reactions 1.
and 2. can also happen in a car engine, and NO2 is
acidic and adds to the polluting acidity of rain as well as
providing nutrients for plants!
- Nitrogen recycling
apart from the atmosphere:
- Nitrogen compounds, e.g.
protein formed in plants or animals, are consumed by animals higher
up the food chain and then bacterial and fungal decomposers
break down animal waste and dead plants/animals to release nitrogen
nutrient compounds into the soil (e.g. in manure/compost)
which can then be re-taken up by plants.
- Nitrogen returned to the
- However, denitrifying
bacteria will break down proteins completely and release
nitrogen gas into the atmosphere.
Keywords & formulae: 1.Reversible Reactions * 2.Reversible
reactions and Equilibrium * 3.The Haber Synthesis of ammonia *
4.The Uses of ammonia-nitric acid-fertilisers * 5.Fertilisers-
environmental problems * 6.The nitrogen cycle NH4Cl(s) <=>
NH3(g) + HCl(g) * NH4NO3(s) <=> N2O(g) + 2H2O(g) *
CuSO4.5H2O(s) <=> CuSO4(s) + 5H2O(g) * BiCl3(aq) + H2O
(l) ==> BiOCl(s) + 2HCl(aq) * N2(g) + 3H2(g) <=> 2NH3(g) *
4NH3(g) + 5O2(g) ==> 4NO(g) + 6H2O(g) * 4NO2(g)+ O2(g)
+ 2H2O(l) ==> 4HNO3(aq) * 2NH3(aq) + H2SO4(aq) ==>
(NH4)2SO4(aq) * NH3(aq) + HNO3(aq) ==> NH4NO3(aq) *
NH3(aq) + HCl(aq) ==> NH4Cl(aq) * NH4Cl + NaOH ==>
NaCl + H2O + NH3 * 2NH4OH(aq) + H2SO4(aq) ==> (NH4)2SO4(aq) + 2H2O(l) *
NH4OH(aq) + HNO3(aq) ==> NH4NO3(aq) + H2O(l) * NH4OH(aq) + HCl(aq) ==>
NH4Cl(aq) + H2O(l) * NH4Cl <=>
NH3 + HCl * NH4NO3 <=> N2O + 2H2O *
CuSO4.5H2O <=> CuSO4 + 5H2O * BiCl3 + H2O
==> BiOCl + 2HCl * N2 + 3H2 <=> 2NH3 *
4NH3 + 5O2 ==> 4NO + 6H2O * 4NO2+ O2
+ 2H2O ==> 4HNO3 * 2NH3 + H2SO4 ==>
(NH4)2SO4 * NH3 + HNO3 ==> NH4NO3 *
NH3 + HCl ==> NH4Cl * NH4Cl + NaOH ==>
NaCl + H2O + NH3 * 2NH4OH + H2SO4 ==> (NH4)2SO4 + 2H2O * NH4OH + HNO3
==> NH4NO3 + H2O * NH4OH + HCl ==> NH4Cl + H2O *
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