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Question: The formal charges of \({{{N}}_{\left( {{1}} \right)}}\), \({{{N}}_{\left( {{2}} \right)}}\) and \({...

The formal charges of N(1){{{N}}_{\left( {{1}} \right)}}, N(2){{{N}}_{\left( {{2}} \right)}} and O{{O}} atoms in :N¨(1)=N(2)=O¨:{{:}}{{{\ddot N}}_{\left( {{1}} \right)}} = {{{N}}_{\left( {{2}} \right)}} = {{\ddot O:}}: are respectively.
A) +1,−1,0 + 1, - 1,0
B) −1,+1,0 - 1, + 1,0
C) +1,+1,0 + 1, + 1,0
D) −1,−1,0 - 1, - 1,0

Explanation

Solution

To solve this we must know that the formal charge is the charge assigned to an atom in a molecule assuming that the bonded atoms equally share the electrons regardless of their electronegativities. Determine the formal charge using the formula to calculate the formal charge.

Formula Used: Formal charge=Valence electrons in free atom−Non - bonding electrons−12Bonding electrons{{\text{Formal charge}}} = {{\text{Valence electrons in free atom}}} - {{\text{Non - bonding electrons}}} - \dfrac{1}{2}{{\text{Bonding electrons}}}

Complete step-by-step answer:
We are given a structure as follows:
:N¨(1)=N(2)=O¨:{{:}}{{{\ddot N}}_{\left( {{1}} \right)}} = {{{N}}_{\left( {{2}} \right)}} = {{\ddot O:}}
We have to calculate the formal charges of N(1){{{N}}_{\left( {{1}} \right)}}, N(2){{{N}}_{\left( {{2}} \right)}} and O{{O}} atoms.
The formula to calculate the formal charge is as follows:
Formal charge=Valence electrons in free atom−Non - bonding electrons−12Bonding electrons{\text{Formal charge}} = {{\text {Valence electrons in free atom}}} - {{\text {Non - bonding electrons}}} - \dfrac{1}{2}{{\text {Bonding electrons}}}
The non-bonding electrons are the lone pairs of electrons that are shown on the atoms.
Now, calculate the formal charge of N(1){{{N}}_{\left( {{1}} \right)}} atom as follows:
We know that nitrogen has 5 valence electrons. From the given structure, non-bonding electrons of N(1){{{N}}_{\left( {{1}} \right)}} are 4 and bonding electrons are 4. Thus,
Formal charge ofN(1)=5−4−12×4{{\text{Formal charge of} }}{{{N}}_{\left( {{1}} \right)}} = {{5}} - {{4}} - \dfrac{1}{2}\times 4
Formal charge of N(1)=5−4−2{{\text{Formal charge of }}}{{{N}}_{\left( {{1}} \right)}} = {{5}} - {{4}} - 2
Formal charge ofN(1)=−1{{\text{Formal charge of} }}{{{N}}_{\left( {{1}} \right)}} = - 1
Thus, the formal charge of N(1){{{N}}_{\left( {{1}} \right)}} atom is −1 - 1.

Now, calculate the formal charge of N(2){{{N}}_{\left( {{2}} \right)}} atom as follows:
We know that nitrogen has 5 valence electrons. From the given structure, non-bonding electrons of N(2){{{N}}_{\left( {{2}} \right)}} are 0 and bonding electrons are 8. Thus,
Formal charge ofN(2)=5−0−12×8{{\text{Formal charge of} }}{{{N}}_{\left( {{2}} \right)}} = {{5}} - {{0}} - \dfrac{1}{2}\times 8
Formal charge ofN(2)=5−0−4{{\text{Formal charge of} }}{{{N}}_{\left( {{2}} \right)}} = {{5}} - {{0}} - 4
Formal charge ofN(2)=+1{{\text{Formal charge of} }}{{{N}}_{\left( {{2}} \right)}} = + 1

Thus, the formal charge of N(2){{{N}}_{\left( {{2}} \right)}} atom is +1 + 1.
Now, calculate the formal charge of O{{O}} atom as follows:
We know that oxygen has 6 valence electrons. From the given structure, non-bonding electrons of O{{O}} are 4 and bonding electrons are 4. Thus,
Formal charge of O=6−4−12×4{{\text{Formal charge of O}}} = 6 - 4 - \dfrac{1}{2}\times 4
Formal charge of O=6−4−2{{\text{Formal charge of O}}} = {{6}} - {{4}} - 2
Formal charge of O=0{{\text {Formal charge of O}}} = 0
Thus, the formal charge of O{{O}} atom is 0.
Thus, the formal charges of N(1){{{N}}_{\left( {{1}} \right)}}, N(2){{{N}}_{\left( {{2}} \right)}} and O{{O}} atoms in :N¨(1)=N(2)=O¨:{{:}}{{{\ddot N}}_{\left( {{1}} \right)}} = {{{N}}_{\left( {{2}} \right)}} = {{\ddot O:}}: are respectively −1,+1,0 - 1, + 1,0.

Thus, the correct option is (B) −1,+1,0 - 1, + 1,0.

Note: The formal charge keeps a track of all the electrons for every atom in a molecule. Thus, formal charge helps in predicting reactivity of a molecule. Formal charge helps in determining the lowest energy Lewis structures.