The formal charges of N_{(1)},N_{(2)} and O a | vedclass

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(B) The formula for formal charge is:$	ext{Formal charge} = [	ext{Total number of valence electrons in the free atom}] - [	ext{Total number of non-

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$-1, +1, 0$

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(B) The formula for formal charge is:$	ext{Formal charge} = [	ext{Total number of valence electrons in the free atom}] - [	ext{Total number of non-bonding (lone pair) electrons}] - \frac{1}{2} [	ext{Total number of bonding (shared) electrons}]$For the end $N$ atom marked $(1)$:$	ext{Formal charge} = 5 - 4 - \frac{1}{2}(4) = 5 - 4 - 2 = -1$For the central $N$ atom marked $(2)$:$	ext{Formal charge} = 5 - 0 - \frac{1}{2}(8) = 5 - 4 = +1$For the end $O$ atom:$	ext{Formal charge} = 6 - 4 - \frac{1}{2}(4) = 6 - 4 - 2 = 0$Thus,the formal charges are $-1, +1, 0$ for $N_{(1)}, N_{(2)}$ and $O$ respectively.

Column-$I$	Column-$II$
$(P)$ $\underline PCl_5 \to PCl_3 + Cl_2$	$(i)$ Change in hybridisation
$(Q)$ $\underline SO_3 \to \underline SO_2 + \frac{1}{2}O_2$	$(ii)$ Change in bond angle
$(R)$ $\underline NH_3 + H^{+} \to \underline NH_4^+$	$(iii)$ Change in shape
$(S)$ $H_2\underline O + H^{+} \to H_3\underline O^{+}$	$(iv)$ Change in oxidation state

The formal charges of $N_{(1)}$,$N_{(2)}$ and $O$ atoms in the structure shown below are respectively:
$:N_{(1)}=N_{(2)}=\ddot{O}:$

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The formal charges of $N_{(1)}$,$N_{(2)}$ and $O$ atoms in the structure shown below are respectively:$:N_{(1)}=N_{(2)}=\ddot{O}:$