Nitrogen occurs in nature as nitrogen gas, N2. This gas makes up about 78% by volume of the atmosphere. There are two isotopes 14N and 15N (about 0.4%). Nitrogen is very unreactive at ambient condition. The only reaction of nitrogen gas at room temperature is with lithium to give Li3N. However nitrogen can be converted to important food compounds by bacteria.
Like carbon, nitrogen { [He] : 2s�, 2p� ] reacts by covalent bonding. It reacts according to the three hybrid atomic electron configurations.
| Hybrid Atomic Orbital | Example |
| sp� | Ammonia (NH3) |
| sp� | Nitrate (NO3‾) |
| sp | Nitrogen gas (N2) |
However with extremely positive electronegative elements, like lithium, it can give away three electrons to form a nitride ion, N‾�, as mentioned above. Also under special conditions it can receive electrons to give the amide ion, NH2‾, and the imide ion, NH‾�. These reactions are greatly aided by the proton ion. After reacting with the protons it actually needed only one or two electrons to complete an octet. When we progress to oxygen and chloride we will see that this is a favoured reaction.
AMMONIA, NH3
Ammonia is a pungent (smelling salt) gas with a boiling point of -33.3�C. When refrigerators were first manufactured ammonia was used as the coolant. Ammonia is hazardous so is soon replaced by other coolants.
In the laboratory be very careful when you open a bottle of ammonia solution. Very likely it is under slight pressure, as ammonia gas is released on standing at room temperature. The vapour can hit you right in your face if you do not take the necessary precaution. Always open a bottle of ammonia in the fumehood. Place a wet piece of cloth over the cap and unscrew slightly to release the pressure before opening the cap completely. If you come into contact with ammonia liquid wash the affected part with plenty of water.
Ammonia is a good molecule to explain the concept of "lone electron pairs", which we have come across in our earlier lessons. Consider a sp� electron configuration for nitrogen. There are five valence electrons and so each sp� would get at least one electron and one having two. Since an atomic orbital can only accommodate a maximum of two electrons (Pauli's Exclusion Principle), therefore only three of the sp� orbitals would be involved in covalent bonding.
The remaining filled sp� is known as a "lone electron pair". Being electron rich it readily comes to the aid of groups that need electrons but are unable to form covalent bonds. So you can imagine this pair as "The Lone Ranger and Tonto" to the rescue of the poor.
Water also has a lone electron pair. Whenever a salt dissolves in water, water molecules would surround the metallic cation, with the lone electron pair directed at the cation. This stabilising force is the most important driving force for the dissolution of compounds in water. The number of water molecules surrounding each cation is known as the coordination number. Six is the common coordination number.
When you are considering the stereochemistry (the structure of molecules) it is very important that you do not forget that a lone electron pair must be considered as a sort of a "bond". For example ammonia is tetrahedral with one lone electron pair plus three N-H sigma bonds. Experiment data showed that ammonia has a "three corner roof" structure. Needless to say the bond angle of "lone electron par"-N-H is much greater than H-N-H.
Tutorial 1
Work out the bondings in nitrogen gas. Why is nitrogen gas so unreactive? (Hint) Answer
HYDROGEN BONDING
      Ammonia |
Let us consider two molecules of ammonia in contact with each other in a solution. The nitrogen being electronegative would take more than its share of the electrons in the sigma-bond between nitrogen and hydrogen. The hydrogen is actually left with no electron part of the time. Since its neighbouring ammonia molecule has a pair of electrons (a rich source of electron) the "deprived" hydrogen is attracted to it. Hydrogen is the smallest atom and so it can approach another atom very closely. Experiments showed that the molecules are electrostatically attached to each other.
A hydrogen bond is not a strong bond (the bond energy is about 20 kJ/mol) and definitely not a permanent bond, like the covalent bond. Such bond exists when hydrogen is covalently bonded to an electronegative element that also has a lone electron pair. So hydrogen bond also exists between two molecules of water and is very important in biological activities (biochemistry).
AMMONIA in WATER
We have seen how hydrogen atoms are attracted to a lone electron pair, so it is not difficult to understand the reaction of a proton with ammonia.
So when ammonia dissolves in water the reaction would be
Tutorial 2
How would you define pKb? Answer
Ammonia is a strong base. It forms salts with mineral acids; NH4Cl, (NH4)2SO4, NH4NO3, and (NH4)2CO3. As a matter of fact the chemistry of the ammonium ion is very similar to that of the Group IA element.
The ammonium cation is a conjugate acid.
AMMONIA with OXYGEN
Ammonia can burn in a rich oxygen atmosphere to give nitrogen and water.
If a piece of hot platinum foil is inserted into a beaker of ammonia exposed to air the ammonia would react with the oxygen to give nitric oxide gas. NO.
The nitric oxide is readily oxidised by the oxygen in the air to brown nitrogen dioxide, NO2.
OXIDES of NITROGEN
NITROUS OXIDE, N2O
This is a linear molecule, NNO. It is relative unreactive. It is inert to halogens, alkali metals and oxygen at ambient condition. In the past it was used as an anesthesia commonly known as "laughing gas".
NITRIC OXIDE, NO
Nitric oxide is prepared by oxidising ammonia. At ambient condition it is easily oxidized to nitrogen dioxide. If left to stand by itself it decomposes to nitrous oxide and nitrogen dioxide.
NITROGEN DIOXIDE, NO2
This brown gas is actually a mixture of NO2 and its dimer (O2N-NO2). In solution it exists as NO+NO3‾. So when nitrogen dioxide dissolves in sulphuric acid the reaction is
The nitrosonium bisulphate, NO+HSO4‾, can actually be isolated. There are also nitrosonium salts known, like NO+ClO4‾.
Nitrogen dioxide dissolves in water to produce a mixture of nitric acid and nitrous acid.
HYDRIDES of NITROGEN
HYDRAZINE, H2N-NH2
The most known is ammonia. This would be followed by hydrazine (boiling point=114�C) . It is stable and when burnt in air it gives of a large amount of energy and very clean products, N2 and water.
This makes it an ideal source of energy, especially for rockets. However much work has to be done to produce hydrazine commercially for this purpose.
Hydrazine is a powerful reducing agent. For example;
HYDRAZOIC ACID, HN3
This is a colourless explosive liquid (b.pt. = 37�C). It is a weak acid and can give salts (nitrides) with electropositive elements like Li and sodium. These salts are reasonable stable and decompose on heating to the metal and nitrogen gas. However the nitrides of the heavier metals like barium, lead and mercury are not stable. It will explode on impact. So they are used as detonation caps in explosives.
OXO ACIDS of NITROGEN
NITRIC ACID, HNO3
Hydrogen nitrate is prepared by dissolving nitric oxide in water in the presence of oxygen. The concentrated acid (70% wt/vol; 0.51 molar) is colourless. However it decomposes easily to give nitrogen dioxide, resulting in the solution being slightly yellowish. The solution is actually a mixture;
So in non-aqueous reactions (like in Organic Chemistry) it is possible to add the nitronium ion, NO2+, to a molecule.
Hydrogen nitrate is a very powerful oxidizing agent. Only gold, platinum, Rhodium and Iridium can withstand hydrogen nitrate, although many metals would form a skin of oxide that can protect the metal from further reaction. If one part of hydrogen nitrate is mixed with three parts of hydrogen chloride (to form aqua regia) the solution can even attack gold and platinum.
However if hydrogen nitrate is diluted with water to below 2M it lost its oxidizing property. Below 0.1M (that is about 93% dissociated) it behaves as a strong acid.
All metals will form nitrates and all nitrates are soluble in water. Many of them will sublime when heated. When heated to high temperature (above 500�C) nitrates decompose to nitrites (NO2‾).
NITROUS ACID, HNO2
Nitrous acid is prepared by adding a solution of sulphuric acid to barium nitrite.
Nitrous acid is a very weak acid. It is unstable and decomposes rapidly to
Nitrites of alkali metals are best prepared by reduction of the nitrates.
Tutorial 3
Why is it not practical to use sodium nitrite in place of barium nitrite, in the preparation of nitrous acid? Answer
