In which of the following electronic configurations will there be a large difference between the second and third ionization energy?

Among the second period elements,the actual ionization enthalpies are in the order $Li < B < Be < C < O < N < F < Ne$. Explain why:
$(i)$ $Be$ has higher $\Delta_{i}H$ than $B$
$(ii)$ $O$ has lower $\Delta_{i}H$ than $N$ and $F$?

If the first ionisation enthalpy of $Li$,$Be$,and $C$ respectively are $520$,$899$,and $1086 \ kJ \ mol^{-1}$,the first ionisation enthalpy (in $kJ \ mol^{-1}$) of $B$ will be:

Given below are two statements. One is labelled as Assertion $A$ and the other is labelled as Reason $R$.
Assertion $A$: The first ionization enthalpy for oxygen is lower than that of nitrogen.
Reason $R$: The four electrons in $2p$ orbitals of oxygen experience more electron-electron repulsion.
In the light of the above statements,choose the correct answer from the options given below.

What is the increasing order of the first ionization enthalpy of $B, P, S$ and $F$ (lowest first)?

The correct order of $2^{nd}$ ionisation potential of carbon,nitrogen,oxygen and fluorine is