13.1Nitrogen Compounds

13.1    Nitrogen Compounds

The lack of  reactivity of nitrogen

  • Nitrogen, N2 exists as a diatomic molecule, two nitrogen atoms are bonded by a triple bond
  • Nitrogen is very unreactive because the bond energy is very high (about +944 kJ mol⁻¹) and reactions involving nitrogen tend to break the entire bond.
  • However, nitrogen still undergoes the following reactions:
    1. When nitrogen and oxygen are struck by lightning in the atmosphere, nitrogen monoxide, NO is In this case, the lightning provides the activation energy required to start the reaction.
      N2(g) + O2(g) → 2NO(g)     ; ΔH = +181 kJ mol⁻¹
    2. Magnesium nitride, Mg3N2 is formed when magnesium is heated in nitrogen. The reaction is exothermic because the ionic bond formed is much stronger than the original bonds and a net energy is released
      3Mg(s) + N2(g) → Mg3N2(s)    ; ΔH = -461 kJ mol⁻¹
  • Carbon monoxide, CO with a triple bond and similarly high bond energy is more reactive because:
    1. it has a dipole moment hence the molecule is polar. They are more attractive to nucleophiles or electrophiles and this initiates a reaction to occur
    2. the reaction involving carbon monoxide will normally not break the entire triple bond. Instead, the bond is partially broken to produce a double-bonded carbon dioxide, CO2.


Ammonia, NH3 and its  reactions

  • Ammonia, NH3 is a trigonal pyramidal molecule with a net dipole moment, hence the molecule is polar


  • Ammonia is a weak base, it is also a Brønsted-Lowry base, hence it is capable of  accepting a hydrogen ion to form ammonium ion,  NH4⁺.
    1. Ammonia reacts with acids to form ammonium salts. For example, the reaction between ammonia and hydrogen chloride, HCl gas:NH3(g) + HCl(g) → NH4⁺Cl⁻
    2. Ammonium salts react with bases to liberate ammonia gas, salt and water is also formed. This is because ammonia is a weak base, the proton accepted is easily removed For example, the reaction between ammonium sulfate, (NH4)2SO4 and calcium oxide, CaO:
      2NH4Cl(s) + Ca(OH)2(s) → CaCl(s) + 2NH3(g) + 2H2O(l)
      The ionic equation is:
      NH4⁺(aq or s) + OH⁻(aq) → NH3(g) + H2O(l)
      This is also a common test for ammonium ions in a compound. When a suspected compound is warmed with sodium hydroxide, NaOH solution, ammonia gas will be released if  it contains ammonium ions. The ammonia gas can be confirmed by using a red litmus paper.
    3. This reaction can also be used to prepare ammonia in school laboratories, the setup is as follow:

      1. This reaction is also known as the displacement of ammonia
      2. Calcium oxide, CaO is used as a drying agent. Other drying agents like calcium chloride, CaCl2 and sulfuric acid, H2SO4 are not used because they react with ammonia


Manufacture of  ammonia – the Haber  process

  • The Haber process is used to manufacture ammonia on a large A brief summary of the Haber process:
    3H2(g) + N2(g) ⇌ 2NH3(g)      ; ΔH = -92 kJ mol⁻¹

    • Hydrogen gas is obtained by reacting methane, CH4(natural gas) with steam at around 700 ºC and the presence of nickel as catalyst.
      CH4(g) + H2O(g) → CO(g) + 3H2(g)
    • Nitrogen gas is obtained by the purification of air. Air which contains mostly a mixture of nitrogen and oxygen gas is reacted with hydrogen gas at high temperature. Oxygen from the air will react with hydrogen to form water.
      2H2(g) + O2(g) → 2H2O(g)
      Oxygen gas is removed, leaving only nitrogen gas behind.
  • The required conditions for optimum yield are:
    i. (400 – 450) ºC.
    ii.   200 atm (equivalent to 20000 kPa).
    iii. Presence of ftne iron as catalyst.
  • Nitrogen and oxygen gas are fed into the reactor in a ratio of 1:3, which is the one demanded by the equation. Excess of reactants are not used because it wastes the space in the reactor and decrease the efficiency of the catalyst, since the excess reactants will have nothing to react with
    1. The production of ammonia is an exothermic reaction in equilibrium. According to Le Chatelier’s principle, in order to shift the position of equilibrium to the right as much as possible(to increase the yield), a low temperature should be used. However, (400 – 450) ºC is not a low temperature.
    2. A low temperature will decrease the rate of reaction albeit having a high yield. The reaction will take a long time to complete and it is not economically plausible
    3. Hence, (400 – 450) ºC is the compromise temperature that produces a good enough yield in a short time
    1. According to Le Chatelier’s principle, the position of equilibrium will shift to the right if the pressure is increased because there are less molecules on the right of the equation. Besides, a high pressure can also increase the rate of  reaction. Hence, a high pressure, 200 atm is  used.
    2. Higher pressures are not used because:
      • it is expensive to build and maintain the pipes and generators to withstand the pressure, this increases the production cost
      • there is a risk of the pipes exploding
    • Hence, 200 atm is the compromise pressure chosen on economic grounds
  • A catalyst of fine iron is used to increase the rate of  reaction. Although it   has no effect on the position of equilibrium, it is essential because without it, the reaction will too long to complete
  • Under these conditions, about 15% of nitrogen and hydrogen converts to ammonia. Unreacted molecules are recycled again so that the overall percentage conversion is about 98%.


Industrial use of  ammonia and nitrogen compounds derived from  ammonia

  • i. Ammonia can be used to make fertilisers. Common fertilisers include ammonium sulfate, ammonium nitrate, ammonium phosphate and urea, CO(NH2)2.
    ii. This is because they contain the element nitrogen. Nitrogen is essential for plants to grow healthy.


  • Ammonia is also a precursor for most nitrogen-containing compounds. One famous example is the manufacture of nitric acid, HNO3 by the oxidation of ammonia in the Ostwald process
  • Nitric acid has several uses:
    1. To make fertilisers such as ammonium nitrate(the main use).
    2. To make explosives such as TNT
  • To be used in the manufacture of dyes, polymers and TNT