The trend of increasing ionization energy across the alkali metals (Li, Na, K) can be explained by the concept of electron shielding and effective nuclear charge.
Ionization energy is the energy required to remove an electron from an atom or ion in the gaseous state. In the case of alkali metals, such as lithium (Li), sodium (Na), and potassium (K), their atoms have a single valence electron in the outermost energy level.
The first ionization energy refers to the energy required to remove the first valence electron from an atom, while the second ionization energy is the energy needed to remove the second valence electron, and so on.
In the case of potassium, it has a higher first ionization energy compared to lithium and sodium because potassium has an additional energy level (shell) compared to them. The extra energy level results in increased electron shielding. Electron shielding refers to the repulsion between electrons in different energy levels, which reduces the effective nuclear charge experienced by the valence electron.
The effective nuclear charge is the net positive charge experienced by an electron and is determined by the number of protons in the nucleus and the amount of shielding from inner electrons. With increased shielding in potassium due to the extra energy level, the valence electron experiences less attraction to the nucleus, making it easier to remove compared to lithium and sodium. This results in a lower second and third ionization energy for potassium because the removal of subsequent electrons requires overcoming a reduced effective nuclear charge.
In summary, the high first ionization energy of potassium compared to other alkali metals is due to the increased electron shielding resulting from an extra energy level. The subsequent ionization energies are lower because the effective nuclear charge experienced by the valence electron is reduced, making it easier to remove additional electrons.