Download Unit 3c File - Uddingston Grammar School

National 4/5
Sub-topic 3C – Fertilisers
Summary
Fertilisers

Increasing world population means that we need more efficient means of food
production.

Growing plants take nutrients from the soil.

y
These nutrients include compounds of nitrogen,
phosphorous and potassium – these are
called essential elements.

Fertilisers are substances which are added to the soil to replace the essential
elements needed for plant growth.
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
Different plants require fertilisers containing different proportions of these
nutrient elements. (N, P and K).
•
Natural fertilisers are made from plant/animal waste i.e. compost and manure.
•
Synthetic fertilisers are made by chemists and need to be soluble because plants take
up nutrients through their roots.
•
Synthetic Fertilisers also need to contain at least one of the three essential
elements. (N, P or K)
•
Synthetic fertilisers are compounds which can contain:
•
Nitrate ion : NO3-
•
Ammonium ion : NH4+
•
Potassium ion : K+
•
Phosphate ion: PO43-
this contains nitrogen
this contains nitrogen
this contains potassium
contains phosphorus
1
Bacterial Methods of Increasing Nitrogen in Soil
•
Certain plants have nitrifying bacteria present in nodules in their roots.
•
These bacteria can convert atmospheric nitrogen (called free nitrogen)
into nitrogen compounds (fixed nitrogen).
•
These nitrogen compounds increase the fertility of the soil.
•
Examples of these plants are peas, beans and clover.
•
This bacterial method of increasing the nitrogen content of
soil is cheaper than chemical methods i.e. using fertilisers.
Ammonia

Formula is NH3

colourless gas, distinctive sharp smell, highly soluble in water producing an alkaline
solution (ammonium hydroxide), turns moist pH paper blue
Fountain Experiment

This experiment shows how soluble ammonia gas is in water and also that it is an alkali.

The water rises up the tube showing that
the ammonia dissolves in water.

The universal indicator changed to blue in
the upper flask showing that ammonia is an
alkaline gas.
Plastics can be made by either:1. Addition Polymerisation
2. Condensation Polymerisation
2
Ammonia

Ammonia can be prepared by the reaction of ammonium compound with alkali:
Ammonium
+
chloride
NH4Cl
+
Sodium
Sodium
hydroxide
chloride
NaOH
NaCl
+ Water + Ammonia
+ H2O

This is also used as a test for the ammonium ion NH4+.

Ammonia gas will be produced which will turn the pH paper blue.
3
+
NH3
Making Fertilisers

Ammonia can be converted into an ammonium compound by its reaction with acid:
2NH3
+
ammonia
H2SO4
(NH4) 2SO4

sulphuric acid
ammonium
sulphate

This new ammonium compound can be used as a fertiliser.

These ammonium compounds can be used as fertilisers as they contain nitrogen.

This is a neutralisation reaction since the ammonia is an alkali and it is reacting with
an acid.

ESTER LINK
More example include:
2NH3
+
ammonia
HNO3
NH4NO3

Nitric acid
ammonium
nitrate
2NH3
+
ammonia
HCl
These are
fertilisers
NH4Cl

Hydrochloric
ammonium
acid
chloride
Percentage Composition of Fertilisers
The percentage of each of the essential elements in a fertiliser can be varied to suit the
crop being grown. Percentage composition allows us to calculate the percentage of any
element present in a compound.
Example:
Calculate the percentage of nitrogen in ammonium sulphate, (NH4) 2SO4.
Step 1.
Write the formula for the compound.
Step 2.
Calculate the formula mass.(2 x 14) +(8 X 1)+(1 x 32)+(4 x16) = 132
Step 3.
Work out the mass of the required element (nitrogen)
Step 4.
Calculate the percentage of the element
(NH4) 2SO4
= 28
28
x 100
132
4
= 21.2%
Problems with Synthetic Fertilisers

They are washed into lakes and rivers by rainwater where they encourage bacteria and
algae to grow.

Bacteria and algae use up dissolved oxygen in the water.

This results in the death of fish as they cannot get enough oxygen to live.
5
Industrial Manufacture of Ammonia – The Haber Process

Ammonia and nitric acid are nitrogen compounds which are used to make fertilisers.

We will look at the industrial production of ammonia first.

Ammonia gas is produced industrially by the Haber Process.


Nitrogen gas comes from the air.
This combines with hydrogen ( which comes from chemicals in the petrochemical
industry).
•
nitrogen
+
N2
+
 ammonia
3H2
 2NH3
Nitrogen and hydrogen react over an iron catalyst to produce ammonia:
N2
•
hydrogen
+ 3H2
2NH3
The formation of ammonia is a reversible reaction and so not all of the nitrogen and
hydrogen are converted into ammonia.
•
At low temperatures a large amount of ammonia is produced slowly.
•
At high temperatures a smaller amount of ammonia is produced more quickly.
•
So the Haber Process is carried out at a moderately high temperature to produce
ammonia at the most economical rate.
•
Temperature too low, ammonia production is too slow.
•
Temperature too high, yield of ammonia is too low.
6
Summary Diagram of The Haber Process
air
natural gas
N2
CH4
+
O2
fractional distillation
of liquid air
hydrogen
nitrogen
H2
N2
* a mixture of nitrogen and hydrogen,
* high pressure,
Unreacted
nitrogen and
* moderately high temperature,
* iron catalyst
hydrogen are
recycled.
The ammonia formed is
separated from unreacted
nitrogen and hydrogen by
liquid ammonia
NH3
cooling.
ammonia gas
NH3 (g)
Production of Nitric Acid For Fertilisers

Nitric acid is also used in the industrial manufacture of fertilisers.

This acid is formed when nitrogen dioxide dissolves in water.

Nitrogen and oxygen can be obtained from the air.

However, nitrogen is not a very reactive gas due to the energy required to break the
triple bonds in the molecules and only combines with oxygen in the presence of a spark,
e.g. during lightning conditions or at the spark-plugs in car engines.

Due to the energy involved in the reaction, it does not provide an economical route
to nitric acid.

The industrial manufacture of nitric acid is by the catalytic oxidation of ammonia, a
route known as the Ostwald Process.
7
The Ostwald Process

The first step is the catalytic oxidation of ammonia : ( Ammonia reacting with
oxygen)
2NH3 (g) + 31/2O2(g)

2NO(g) + 3H20(g)
This requires a platinum catalyst, atmospheric pressure and 600-900oC
Catalytic Oxidation of Ammonia in the Lab

This process is carried out at a moderately high temperature
to allow it to proceed fairly quickly and produce a good yield
of nitrogen monoxide.

A platinum catalyst is used.

Since the reaction is exothermic it is not
necessary to continue heating it after the
reaction has started since it will supply
sufficient energy to continue at a reasonable
rate.

The second step is : nitrogen monoxide reacts
with oxygen to produce nitrogen dioxide.
2NO(g)
•
+
O2(g)

2NO2(g)
The third step involves the nitrogen dioxide reacting
with more oxygen and water to form nitric acid.
4NO2(g)
+
O2(g)
+
2H2O(l)
4HNO3(aq)

8

Since the reaction is exothermic (heat is given out) there is no need to keep
heating once the reaction has started.

As with the Haber Process, the higher the temperature the faster the rate of the
reaction but the lower the yield … so a moderately high temperature is used.

In the manufacture of nitrogen fertilisers, nitric acid is converted to nitrate
compounds (solids) by the reaction with alkalis.
Summary of The Ostwald Process
9
What Happens to the Nitric acid ?
•
The nitric acid produced in the Ostwald process is used to make fertilisers.
•
It can be reacted with ammonia to produce the fertiliser ammonium nitrate.
Ammonia
+
Nitric Acid
Ammonium nitrate
NH3
+
H+NO3-
NH4+NO3-
Fertiliser
•
This is a neutralisation reaction.
•
The nitric acid can react with other bases in neutralisation reactions to from
more fertilisers.
Nitric acid +
H+NO3-
+
Potassium
Potassium
hydroxide
nitrate
K+OH-
K+NO3-
Fertiliser
10
+
+
Water
H 2O