Introduction to Halides group:
We are familiar that group 15 elements form compounds mostly in two oxidation states, +3 and +5. Both kinds of their halides, that is, trihalides and also pentahalides are identified.
All the elements of Group 15 form trihalides, MX3. Of all nitrogen halides, just NF3 is stable because of the strong N-F bonds. All the other trihalides of nitrogen, NCl3, NBr3 and NI3 are unstable. Of such, NCl3 is explosive and NBr3 and NI3 exist just as their ammoniates, that is, as ammonia adducts and detonate on eliminating ammonia.
Apart from NX3, all the trihalides hydrolyze quickly to either hydrated oxides, example -As2O3·nH2O or acids, H3PO3 or oxochloride, SbOCl and BiOCl. NCl3 on hydrolysis provides NH3 and HOCl rather than HNO2 (nitrous acid) and HCl:
NCl3 + 3H2O → NH3 + 3HOCl
PCl3 + 3H2O → H3PO3 + 3HCl
SbCl3 + H2O → SbOCl + 2HCl
BiCl3 + H2O → BiOCl + 12HCl
The method or procedure of the hydrolysis of phosphorus trichloride is identical to that of silicon tetrachloride. It comprises the formation of an intermediate four coordinate species H2O→PCl3. The formation of this kind of intermediate is not possible for nitrogen due to the non-availability of d-orbitals. NCl3, thus, hydrolyses via a different method, which comprises the attack of the Lewis base, water, on the partially positive chlorine to eliminate it in the form of a protonated hypochlorous acid molecule. Proton exchange by the Cl2N- anion then produces one molecule of HOCl and NHCl2 each.
Substitution of two more chlorine atoms by hydrogen produces the NH3 molecule:
Fig: Hydrolysis of Nitrogen trichloride
The trihalides, apart from NX3, act as Lewis acids, that is, electron pair acceptors, forming halo anions, SbCl4, BiCl4, and so on. The trihalides can be oxidized to pentahalides. The easiness of oxidation reduces from PX3 to BiX3. Though, NX3 is not oxidized at all. The trihalides function as halogenating agents. PCl3 reacts by NH3, halogens, HI, S, O2 and AgNCO to provide P(NH2)3, PX2Cl3, PI3, PSCl3, POCl3 and P(NCO)3, correspondingly. The given chemical equations symbolize these reactions:
PCl3 + 3NH3 → P(NH2)3 + 3HCl
PCl3 + Br2 → PBr2Cl3
PCl3 + HI → PI3 + 3HCl
PCl3 + S → PSCl3
2PCl3 + O2 → 2POCl3
PCl3 + 3AgNCO → P(NCO)3 + 3AgCl
Trihalides of the lighter elements of the group are mostly covalent in nature, thus, they can exist as discrete molecules that encompass a pyramidal structure, such as NH3. The trihalides of heavier elements, example - BiF3, are ionic in nature.
Nitrogen doesn't form any pentahalides. The stability of other pentahalides of the group reduces in the order P > Sb > As > Bi and in the order F > Cl > Br for halogens. PF5, PCl5, PBr5, AsF5, SbF5, SbCl5, and BiF5 are stable. Though, PBr5 and SbCl5 readily lose halogen forming trihalides. The pentahalides get more oxidising as we go down in the group, apart from arsenic pentahalides, AsX5, which are unpredictably more oxidizing than those of the element following it in the group. Each and every pentahalides encompass trigonal bipyramidal structure in the gas phase. Though, their structures vary in the solid state, example - PCl5 divides into [PCl4] [PCl6] ¯ which encompass tetrahedral and octahedral structures, correspondingly; PBr5 divides to give [PBr4]+Br-. Pentahalides are obtained via the action of surplus of halogen on trihalides:
PX3 + X2 ↔ PX5
Pentahalides are good halogenating agents, example - PCl5 is employed for conversion of alcohols to alkyl halides and acids to acylchlorides as illustrated below:
ROH + PCl5 → RCl + POCl5 + HCl
RCOOH + PCl5 → RCOCl + POCl3 + HCl
They hydrolyze to give oxoacids, H3PO4, H3AsO4, or hydrated oxides such as Sb2O, nH2O or oxohalides such as BiOCl; partial hydrolysis of PCl5 gives POCl3. Nearly all the pentahalides can accept the electron pair and act as Lewis acids, as in PF6, PF5 and so on.
Introduction of Oxides group:
As stated, nitrogen forms a number of oxides, N2O, NO, N2O3, NO2 or N2O4 and N2O5, and as well very not stable NO3 and N2O6. All such oxides of nitrogen show pπ-pπ multiple bonding between nitrogen and oxygen. This doesn't take place with the heavier elements in the group and as a result nitrogen forms a number of oxides which encompass no P, As, Sb or Bi analogues.
Oxides of Nitrogen:
Significant features of oxides of nitrogen are illustrated in the table shown below. Let us first illustrate the preparation, properties and structures of these oxides.
Table: Oxides of Nitrogen
Extensively dissociated as gas
Dinitrogen tetra oxide
Extensively dissociated to NO2 as gas
Unstable as gas ; ionic solid
Not well characterized and quite unstable
N2O is obtained usually by heating NH4NO3:
NH4NO3 → N2O + 2H2O
NO is prepared by the reduction of 8M HNO3 by reducing agents such as Cu or by reduction of nitrous acid or nitrites by Fe2+ or I- ions:
3Cu + 8HNO3 → 3Cu(NO3)2 + 2NO + 4H2O
2NaNO2 + 2FeSO4 + 3H2SO4 → 2NaHSO4 + Fe2(SO4)3 + 2NO + 2H2O
2NaNO2 + 2NaI + 4H2SO4 → 4NaHSO4 + 2NO + I2 + 2H2O
NO is formed as an intermediate in the preparation of nitric acid via oxidation of NH3. N2O3 is obtained as an intense blue liquid or a pale blue solid on cooling the equimolar mixture of NO and NO2:
NO + NO2 → N2O3
On warming, its colour fades because of its dissociation into these two oxides.
NO2 can be manufactured by reduction of cone. HNO3 by Cu or through heating heavy metal nitrates:
Cu + 4HNO3 → Cu(NO3)2 + 2NO2 + 2H2O
2Pb(NO3)2 → 2PbO + 4NO2 + O2
N2O5 is the anhydride of HNO3. This is best made by dehydrating HNO3 by P4O10 at low temperatures:
4HNO3 + P4O10 → at 250K → 2N2O5 + 4HPO3
=> Properties and Structure:
The oxides of nitrogen are all oxidising agents, N2O even supporting the combustion of S and P. NO that is thermally more stable, supports the combustion of Mg and P however not of S. Sulphur flame is not hot adequate to decompose it. N2O and NO are neutral, whereas the other oxides are acidic.
N2O is isoelectronic by CO2 and as well consists of a linear structure. Though, dissimilar to CO2, N2O consists of a small dipole moment. The resonance structures of the molecule are given in the figure shown below.
Fig: Resonance Structures of Nitrous Oxide
NO consists of a total of 15 electrons. This is impossible for all of them to be paired and therefore this is an odd electron molecule. In the gaseous state, it is paramagnetic. Though, the liquid and solid states are diamagnetic as loose dimers (figure shown below) are formed cancelling out the magnetic effect of the unpaired electrons.
Fig: Resonance Structures of Nitric Oxide
The bonding in NO is best illustrated by the molecular orbital theory. Molecular orbital electronic configuration of NO molecule can be symbolized as σ1s2, σ*1s2, σ2s2, σ*2s2, σ(2pz)2, π(2px)2, π(2py)2, π*(2px)1. This provides a bond order of 2.5. If the unpaired electron occupying the anti-bonding π*2py orbital is removed, nitrosonium ion, NO+, is formed and the bond order becomes 3. This is reflected in the shortening of the bond length from 115 pm in NO to 106 pm in NO+. Nitrosonium ion is stable and forms salts such as NO+Cl¯. This is isoelectronic with CO and forms complexes with the transition elements. The brown ring made up in the test for nitrates is due to the formation of a complex of iron, [Fe(H2O)5NO]2+.
NO2 having 23 electrons is again an odd electron molecule. In the gaseous state it is paramagnetic. On cooling, the gas condenses to a brown liquid and ultimately to a colourless solid both of which are diamagnetic because of dimerisation. NO2 molecule is angular by ONO angle of 134°. The O-N bond length is 120 pm, intermediate between the single and double bond. The odd electron is on nitrogen. The dimer has been illustrated to encompass a planar structure. The N-N bond length is extremely large, 175 pm, making this an extremely weak bond.
Fig: Resonance Structure of (a) NO2 and (b) N2O4
Oxides of Phosphorus, Arsenic, Antimony and Bismuth:
As P, As and Sb form oxides in both +3 and +5 oxidation states while only one oxide of bismuth, Bi2O3 is known. Keep in mind that the stability of higher oxidation states decreases on going down the group. Apart from Bi2O3 others exist in dimeric form. The fundamental character of oxides increases on descending the group. The oxides of P and As are acidic, those of Sb amphoteric and of Bi completely basic. As well the higher oxidation states are more acidic.
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