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Fertilisers

Fertilisers. Population and Food Needs. The ever increasing world population means more food is needed and fertilisers are used to grow plants efficiently. . Organic fertilisers such as animal manure can be used but additional chemical fertilisers are also needed.

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Fertilisers

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  1. Fertilisers

  2. Population and Food Needs • The ever increasing world population means more food is needed and fertilisers are used to grow plants efficiently.

  3. Organic fertilisers such as animal manure can be used but additional chemical fertilisers are also needed. • Growing plants need nutrients - soluble compounds containing nitrogen, phosphorus and potassium which are absorbed by roots. • This is why ammonium salts (such as NH4Cl), potassium salts and nitrates (such as KNO3) and phosphates (such as Na3PO4) are important fertilisers as they contain the essential elements that plants need and they are soluble in water.

  4. Fertilisers containing Nitrogen (N), Phosphorus (P) and Potassium (K) are known as NPK fertilisers. • When crops are harvested, these nutrients are removed from the soil in the plant and need to be replaced.

  5. Effects of Shortages of Nutrients on Plants

  6. Percentage Composition • You have learned to calculate the % of elements in a compound. Ammonium nitrate, NH4NO3 has a formula mass of 80. Mass of Nitrogen present in the formula mass = 14 x 2 = 28 Percentage of Nitrogen = (28/80) x 100 = 35%

  7. The Nitrogen Cycle

  8. The Nitrogen Cycle

  9. Clover, pea and bean plants change atmospheric nitrogen into soluble nitrogen compounds by 'fixing' nitrogen. They are called leguminous plants. • This happens because swellings (or nodules) on the roots contain nitrifying bacteria which add nitrogen to the soil, increasing fertility. • It is cheaper than chemical fertilisers and without pollution problems. • Some nitrifying bacteria are free-living.

  10. Problems Caused by the Use of Fertilisers • Fertilisers need to be soluble to get into plants. • Rain can wash fertilisers into rivers and lochs and cause harmful algae and bacteria to grow rapidly. • These remove oxygen from the water and kill fish and plants.

  11. Converting Unreactive Nitrogen into Ammonia - the Haber Process • Nitrogen (from the fractional distillation of liquid air) and Hydrogen (from natural gas) are mixed together at a moderately high temperature (400oC) and high pressure with an iron catalyst and nitrogen hydride (ammonia, NH3) is made. N2(g) + H2(g)  NH3(g) (Unbalanced)

  12. The previous reaction is reversible i.e. it can go in both directions. • If too high a temperature is used, the ammonia changes back to nitrogen. • This is why a moderately high, but not too high a temperature is used. • It also explains why not all of the nitrogen and hydrogen are changed into ammonia.

  13. Activity

  14. Making Ammonia in the Laboratory • Heating any ammonium compound with alkali produces ammonia NH4Cl + NaOH  NH3 + NaCl + H2O • This reaction can be used to prove that a compound is an ammonium compound

  15. Properties of Ammonia • Ammonia is a colourless gas which turns moist pH paper blue/purple i.e. it is an alkali. NH3(g) + H2O(l)  NH4OH(aq) • Ammonia has a characteristic smell. • Ammonia is very soluble in water as shown by the fountain experiment. • Ammonia reacts with acids to make salts. These salts can be used as fertilisers as they contain nitrogen. NH3 + HNO3 NH4NO3

  16. Activity

  17. Making Nitric Acid • Nitric acid is needed to make some fertilisers. • It can be made as follows: • Nitrogen dioxide, a brown gas, can be made when air (21% oxygen and 78% nitrogen) is sparked for about 20 minutes N2 + O2 NO2 (Unbalanced) N2 + 2O2 2NO2 (Balanced)

  18. Nitric acid is made if the brown gas formed shaken with water and oxygen. 2H2O + 4NO2 + O2 4HNO3 • The spark plug causes the same reaction in a petrol engine, as does lightning. • Time and cost make this method unsuitable to make large quantities of nitric acid

  19. Nitrogen is Unreactive • Many elements burn in air to form element oxides, but nitrogen is very unreactive. • In Topic 4 you discovered that there is a triple covalent bond holding nitrogen atoms tightly in the diatomic nitrogen molecule. • It is hard to break this bond and nitrogen does not burn in oxygen. • The spark plug and lightning provide more than enough Activation energy to cause this reaction to happen.

  20. The Catalytic Oxidation of Ammonia (Ostwald Process)

  21. This is an economic route to making nitrogen dioxide and then nitric acid. It can be done in the laboratory. • The reaction is an example of oxidation as ammonia reacts with oxygen on the platinum catalyst.

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