Electronic configurations of four elements $A$,$B$,$C$,$D$ are given below:
$A$) $1s^2 2s^2 2p^6 3s^1$
$B$) $1s^2 2s^2 2p^6 3s^2 3p^1$
$C$) $1s^2 2s^2 2p^6 3s^2$
$D$) $1s^2 2s^2 2p^6 3s^2 3p^2$
The correct order of first ionization enthalpy of these elements is:

Given below are two statements: one is labelled as Assertion $A$ and the other is labelled as Reason $R$:
Assertion $A:$ The first ionisation enthalpy decreases across a period.
Reason $R:$ The increasing nuclear charge outweighs the shielding across the period.
In the light of the above statements,choose the most appropriate from the options given below:

What is the decreasing order of the ionization energy for the following elements?

Which of the following is the correct order of the first ionization energy for $B$,$C$,and $N$?

The first ionization enthalpies of $C$,$N$,and $F$ are $1086$,$1402$,and $1681 \, kJ \, mol^{-1}$ respectively. Which of the following values is the most likely first ionization enthalpy of $O$ ($1310$ or $1510 \, kJ \, mol^{-1}$)?

Which of the following configurations represents atoms of the elements having the highest second ionization energy?