Lack of Reactivity of Nitrogen

  • Nitrogen molecule has three strong covalent bonds
  • Bond is very strong and requires high energy for splitting the two nitrogen atoms of a molecule.
  • It reacts only under extreme temperature or pressure or in presence of catalyst.
  • 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 mol¹) and reactions involving nitrogen tend to break the entire bond.
  • However, nitrogen still undergoes the following reactions:
    • When nitrogen and oxygen are struck by lightning in the atmosphere, nitrogen monoxide, NO is produced. In this case, the lightning provides the activation energy required to start the reaction.

 

N2(g) + O2(g) 2NO(g) ; ΔH = +181 kJ mol¹ ii. 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 mol¹

  • Carbon monoxide, CO with a triple bond and similarly high bond energy is more reactive because:
    • 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.
    • 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.

 

Ammonium

  • Lone pair of es of nitrogen forms a coordinate bond with the H+ ion
  • Formation: NH3(g) + H+ NH4+
  • Shape: tetrahedral
  • Bond angle: 109.5o
  • Bond length: equal lengths

Displacement of ammonia from its salts:

Any Ammonium Salt + Any Base >​ Ammonia Gas + salt + water (conditions :warm)

Uses of Ammonia & its Compounds

  • Used in the production of nitric acid
  • Used in the production of inorganic fertilizers
  • Used in the production of nylon
  • Used in the production of explosives

 

Manufacture of ammonia – the Haber process

The Haber process is used to manufacture ammonia

3H2(g) + N2(g) 2NH3(g) ; ΔH = -92 kJ mol¹


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:

  • (400 – 450) ºC.
  • 200 atm (equivalent to 20000 kPa).
  • Presence of fine 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.

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.

  • 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.
  • Hence, (400 – 450) ºC is the compromise temperature that produces a good enough yield in a short time.

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. 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. iii. 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 its derivatives

Ammonia can be used to make fertilisers.

  • Common fertilisers include ammonium sulfate, ammonium nitrate, ammonium phosphate and urea, CO(NH2)2.
  • 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. Example is the manufacture of nitric acid, HNO3 by the oxidation of ammonia in the Ostwald process.

  • Nitric acid has several uses:
    • To make fertilisers such as ammonium nitrate (the main use).
    • To make explosives such as TNT.
    • To be used in the manufacture of dyes, polymers and drugs.

 

Oxides of Nitrogen

N2(g) + O2(g) → 2NO(g) or N2(g) + ½O2(g)→ NO2(g)

  • Naturally: during lightning, EA provided for N2 to react
  • Man-made: in car engine, high temp. and pressure
  • Catalytic convertors: exhaust gases passed through catalytic convertors containing a catalyst (platinum/ palladium/nickel) helping to reduce oxides to nitrogen.
  • Catalytic role in oxidation of sulphur dioxide:

Pollution

Acid Rain: SO3 + H2O→ H2SO4

2NO2 + H2O → HNO3 + HNO2 or NO2 + H2O + ½O2 → HNO3

  • Damages trees & plants, kills fish and other river life, buildings, statues and metal structures

 

Combustion Pollutants:

  • Nitrogen oxide (NO): formed by reaction of N2 and O2 in the engine, forms acid rain and respiratory problems
  • Atmospheric oxides of nitrogen (NO & NO2) can react with unburned hydrocarbons to form peroxyacetyl nitrate (PAN) which is a component of photochemical smog
  • Carbon monoxide (CO): source: incomplete combustion of hydrocarbon fuel, toxic effect on hemoglobin

 

Pollution formation equations

Nitrogen combines with oxygen – N2 (g) + O2 (g) ——> 2NO(g)

Nitrogen monoxide is oxidised – 2NO(g) + O2 (g) ——> 2NO2 (g)

Incomplete hydrocarbon combustion – C8H18 (g) + 8½O2(g) ——> 8CO(g) + 9H2O(l)

 

Environmental Consequence of Using Nitrogen Compounds

  • When excessive nitrate or ammonium fertilisers are used, the unabsorbed ones will dissolve in rain water and it leaches into lakes and rivers.
  • An excess of these chemicals in the waters can promote the growth of algae, eventually causing an algae bloom. The algae grow exponentially across the surface of water, blocking sunlight from the reach of aquatic plants and causes the plants to die.
  • The algae grow faster than being consumed, eventually a large number of algae die without being consumed. When their remains decompose, the process takes up a lot of oxygen from the water. The oxygen level in the water will eventually reach a level where no life can sustain.
  • This process of excess growth leading to the destruction of life in the water – eutrophication.
  • Since nitrates are soluble in water, removing them from drinking water is very expensive. High levels of nitrates in drinking water can cause a disease in young babies called ‘blue baby syndrome’.
  • Nitrates in water can also potentially cause stomach cancer.

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