Effect of pressure
High pressure would increase the yield of ammonia since the reaction proceeds with a decrease in volume. There fore a pressure of 200-500 atm is used. At a high pressure, the reacting molecules collide more frequently thus increasing the rate of the reaction.
Effect of temperature
The formation of ammonia from its elements is an exothermic reaction and there fore low temperature will cause a better yield of ammonia. But at low temperature, the reaction will be slow thus a moderate temperature of 400-500˚C is used together with a catalyst.
Effect of a catalyst
To make the reaction proceed faster, a catalyst is used. The catalyst used is finely divided iron impregnated with alumina. The catalyst should be finely divided to increase on the surface area since the reaction occurs at the surface.
Effect of concentration
Since the reaction is reversible, we use in excess any of the reactants inorder to get a better yield of ammonia. However, in practice, nitrogen and oxygen are used in the ratio of 1:3 respectively.
Uses of ammonia
- It is used in the manufacture of fertilizers like ammonium phosphate and urea.
- Ammonia solution is used to soften hard water.
- Ammonia is used in the manufacture nitric acid.
- Liquid ammonia is used in large scale refrigerating plants such as in ships and ware houses.
- It is used in the manufacture of sodium carbonate in the Solvay process.
a) All common ammonium salts are white crystalline substances; soluble in water and are ionic compounds.
- Assemble the glass apparatus as shown in the diagram above.
- Put some potassium nitrate crystals or sodium nitrate crystals in the bulb of the retort with concentrated sulphuric acid.
- Heat the mixture gently and then collect the nitric acid in a water cooled receiver.
- The potassium nitrate crystal gradually dissolves and effervescence occurs.
The nitric acid distills and collects in the cooled receiver as a yellow liquid (the yellow colour is due to the presence of dissolved nitrogen dioxide gas), while drops of the acid can be seen running down the bulb and neck of the retort. The brown fumes are nitrogen dioxide formed by slight decomposition of the nitric acid by heat.
NB. The apparatus used must be glass because nitric acid quickly attacks other materials like cork and rubber tubing.
Any metallic nitrate when heated with concentrated sulphuric acid produces nitric acid.
Industrial Preparation of nitric acid (The Ostwald’s process)
Nitric acid is manufactured by the catalytic oxidation of ammonia and then dissolving the products in water.
- Ammonia from the Haber process is mixed with excess air and passed over a platinum (90%)/rhodium (10%) gauze catalyst. The catalyst is heated to red hot to start the reaction and since the reaction is exothermic, no heating is required once the reaction starts. Here ammonia is oxidized to colorless dinitrogen gas.
Properties of nitric acid
i) It is a colourless fuming liquid (if pure)
ii) It is corrosive just like other acids
iii) It turns blue litmus red and has no effect on red litmus paper
a) Thermal decomposition
Upon heating, concentrated nitric acid decomposes to give off brown fumes of nitrogen dioxide, oxygen and water vapour.
Other metals do not react with dilute nitric acid to produce hydrogen, they are simply oxidized to their corresponding nitrates and the nitric acid is reduced to nitrogen dioxide.
c) Oxidation reactions of nitric acid
Concentrated nitric acid is a powerful oxidizing agent i.e. it readily gives up its oxygen. It converts most metals to their corresponding nitrates and non metals such as carbon are oxidized to their oxides. In all cases, nitric acid is itself reduced to nitrogen dioxide. For example
i) Concentrated nitric acid oxidizes brown copper metal to copper (II) nitrate and the nitric acid is itself reduced to nitrogen dioxide.
Concentrated nitric acid renders iron and aluminium passive due to formation of an oxide layer which serves as a protective coating and therefore prevents the metals from reacting any further.
Uses of nitric acid
- It is used in the manufacture of dyes.
- It is used in the manufacture of fertilizers e.g. ammonium nitrate.
- Used in the manufacture of explosives.
- Since it is an oxidizing agent, it is used in the manufacture of nylon.
- It is used in the refining of precious metals.
These are salts of nitric acid. All nitrates are soluble in water.
Action of heat on nitrates
All the nitrates decompose on heating. The thermal decomposition of metal nitrates depends upon the position of the metal in the reactivity series.
When potassium and sodium nitrates are heated, they melt into colourless liquid decompose to give pale yellow nitrites and oxygen gas. E.g.
Test for nitrates
All nitrates irrespective of their position in the reactivity series under go the same reaction with iron(II) sulphate and concentrated sulphuric acid and the reaction is used as a test for all soluble nitrates.
To the nitrate solution in a test tube, add an equal volume of freshly prepared iron(II) sulphate solution. The test tube is tilted and concentrated sulphuric acid is carefully poured down the side of the test tube.
A brown ring is formed at a junction between sulphuric acid and iron(II) sulphate. The brown ring is of a compound with the formula FeSO4.NO.
The nitrogen cycle
This is a balance that exists between reactions that take nitrogen out of the air and out of the soil and reactions that put nitrogen into the air and into the soil. Nitrogen constitutes 78% by volume of air and it is an essential element in all living things.
Plants obtain nitrogen mainly in form of dissolved nitrates by absorption from the soil. When the nitrogen in the soil is not replenished, it leads to infertility and poor yield in crops. Animals obtain nitrogen by feeding on plants. Nitrogen is supplied to the soil through: death and decay of plants and animals; excretion by animals; nitrogen fixation by colonies of bacteria in root nodules and bacteria that live freely in the soil; lightning -causes some slight combination of nitrogen and oxygen which leads to passage of nitrogen into the soil as nitrates; industrial fixation through Haber process forming nitrogenous fertilizers such as ammonium sulphate. Denitrifying bacteria such as Pseudomonas denitrificans convert ammonium salts in the soil to atmospheric nitrogen
Sample question on Nitrogen and its compounds
- Outline briefly how a sample of nitrogen can be obtained from (a) air, (b) an ammonium salt.
Describe an outline of the industrial process showing how nitrogen is converted to ammonia. The reaction involved in this process is reversible. Indicate three ways in which the yield of ammonia can be made maximum. State two industrial uses of ammonia.
- Make a fully labeled drawing and give equation to show how you would prepare dry sample of ammonia in the laboratory starting from a named ammonium salt and a named alkali.
Giving equations and reaction conditions, outline how nitric acid is manufactured from ammonia.
- Describe the action of heat on ammonium chloride. By what reaction would you obtain (i) ammonia from ammonium chloride (ii) nitrogen from ammonia? How can you account for the fact that a solution of ammnia in water will turn litmus blue and give brown precipitate when mixed with a olution of iron(II) chloride?
- Describe how ammonia is manufactured from its elements. State the source of each element. Outline three differences between nitrogen/hydrogen mixture and ammonia. How would you show that ammonia is a very soluble gas?
- Explain how ammonia is converted into nitric acid on a large scale. Describe two reactions in which nitric acid (i) acts as an acid (ii) acts as an oxidizing agent. Outline some uses of nitric acid. Calculate the percentage of nitrogen in pure ammonium nitrate.
Describe with equations what happens when ammonia is passed (i) into dilute sulphuric acid, (ii) over heated copper(II) oxide.
- Nitric acid can be prepared in the laboratory by heating solid sodium nitrate with concentrated sulphuric acid. Make a labeled drawing of the apparatus you would use for the preparation and write equation for the reaction. Explain why sulphuric acid is used in this preparation rather than hydrochloric acid.
- Describe an experiment to show how nitrogen dioxide is prepared in the laboratory. Describe the oxidation reaction of nitrogen dioxide, use equations to illustrate. Mention two uses of nitrogen dioxide.
- With the aid of equations, describe the reactions of ammonia with: oxygen, hydrogen chloride, chlorine, water, metal oxides and metallic ions.