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Neil Bartlett found that platinum hexa-fluoride reacts with oxygen to form a solid ionic compound of the formula [O₂]⁺ [Pt F₆]^-
O₂(g) + Pt F₆(g)→ O₂⁺ Pt F₆^–(s)
This reaction shows that Pt F₆ is a very strong oxidising agent which can attract electron even from molecular oxygen. Now since the ionization potentials of oxygen molecule (1182 kj mol^{– 1}) and Xenon atom (1170 kj mol^{– 1}) are comparable. Bartlett believed that if PtF₆ can oxidse can oxidise oxygen molecules, it should also be able to oxidise xenon atom.

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isostructural can be defined as having the same crystal structure but not necessarily a similar chemical composition.

(a) ICl₄^– = XeF₄


(b) IBr₂^– =XeF₂


(c) BrO₃^–=XeO₃ 

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Oxidation can be calculated as:
 H₃PO₃ 
oxidation state of hydrogen is +1
oxidation state of oxygen is -2

calculation for P oxidation state is :
1 x3+ P +(-2) x3 =0
3+P+(-6)=0
3+P=6
P=6-3 
=3
here oxidation state is +3

(ii) PCl₃
Oxidation state of chlorine is -3
thus
P+(-3) =0
P = 3
here oxidation state is +3

(iii) Ca₃P₂

oxdation state of calcium is +2
 2 x3 +2P =0
6+ 2P =
2P=-6
P= -6/2 =-3

oxidation state is -3

(iv)Na₃PO₄ 
here oxidation state of sodium is +1
oxidation state of oxygen is -2
1 x3 +P +(-2) x4=0
3 +P +(-8) =0
3+P =8
P=8-3 =5
oxidation state is +5

(v) POF₃
oxidation state fulorine is (-1)
oxidation state is (-2)
P+(-2) +(-1) x3 =0
P= 5
oxidation state is +5

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(i) Electronic configuration: The valence shell electronic configuration of these elements is ns² np³. The s orbital in these elements is completely filled and p orbitals are half filled, making their electronic configuration extra stable.

(ii) Atomic Size: Covalent and ionic (in a particular state) radii increase in size down the group. There is a considerable increase in covalent radius from N to P. However, from As to Bi only a small increase in covalent radius is observed. This is due to the presence of completely filled d and f or f orbitals in heavier members. 

(iii)Oxidation State: The common oxidation states of these elements are –3, + 3 and + 5. The tendency to exhibit –3 oxidation state decreases down the group due to increase in size and metallic group. In the last member of the group, bismuth hardly forms any compound in –3 oxidation state. The stability of + 5 oxidation state decreases down the group. The stability of + 5 oxidation state decreases and that of + 3 state increases (due to invert pair effect) down the group. Nitrogen exhibits + 1, + 2, + 4 oxidation states also when it reacts with oxygen. Phosphorus also shows +1 and + 4 oxidation states in some oxo acids.

(iv) Ionization enthalpy: Ionization enthalpy decreases down the group due to gradual increase in atomic size. Because of the extra stable half filled p orbitals electronic configuration and smaller size, the ionization enthalpy of the group 15 elements is much greater than that of group 14 elements in the corresponding periods. The order of successive ionization emthalpies are expected as ΔH₁, < ΔH₂ < ΔH₃.

(v) Electronegativity: The electronegativity value, in general, decreases down the group with increasing atomic size. However, amongst the heavier elements, the different is not that much pronounced.

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Stability of an ionic compound depend on its lattice energy. More the lattice energy of a compound more stable it will be.The electron gain enthalpy value for O → O^– is –141 kj mol^–1 and O → O^2– is 702 kj mol^–1. Combination of oxygen with other elements is often strongly exothermic which helps in sustaining the reaction. However, to initiate the reaction some external heating is required as bond dissociation enthalpy of oxygen-oxgen double bond is high (493.4 kj mol^–1). Hence, we find the formation of a large number of oxides having O^2– species and not O^–