$$ \begin{array}{|c|c|c|} \hline & \text{First Ionization} & \text{Atomic} \\ \text{Element} & \text{Energy } & \text{Radius} \\ & \text{(kJ/mol)} & \text{(pm)} \\ \hline \hline \ce{B} & 801 & 85 \\ \hline \ce{C} & 1086 & 77 \\ \hline \ce{N} & 1400 & 75 \\ \hline \ce{O} & 1314 & 73 \\ \hline \ce{F} & 1680 & 72 \\ \hline \ce{Ne} & 2080 & 70 \\ \hline \end{array} $$

The table above shows the first ionization energy and atomic radius of several elements. Which of the following best helps explain the deviation of the first ionization energy of oxygen from the overall trend?

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If electrons are closer to the core, it implies there is a greater force of attraction between the electron and the positive core. This is indeed why the radius decreases down a period. The positive charge of the core increases and each electron feels a greater force of attraction.

Oxygen breaks this trend, since the ionization energy is much smaller despite the molecule being smaller than elements nearby on the periodic table. This is because the fourth \(p\) orbital electron needs to be added an orbital that already contains an electron. Therefore the fourth electron experiences repulsive forces and is easier to remove from the atom.