Question

I've read that the effective nuclear charge increases down the group.

Zeff = Z-S

Zeff effective nuclear charge



Z= number of protons in the nucleus (atomic number)

S= shielding constant
This seems incorrect. As we go down the group the number of protons increases and the shielding constant also increases. We can approximate the shielding constant with the number of electrons between the nucleus and the valence electrons. Hence it seems that the effective nuclear charge is the same within a given group.


The effective nuclear charge doesn't depend on the distance ( Shown through the formula above ). So the valence electrons are further away from the nucleus and experience the same effective nuclear charge down the group. Hence they have less attraction to the nucleus and the atomic radius increases down the group.

Is my intuition correct? Or does the effective nuclear charge increase down the group?

Answer 1

The effective nuclear charge does not increase down the group. The increase in atomic radius down the group is primarily due to the addition of new electron shells.

Step-by-step explanation:

The effective nuclear charge (Zeff) does not increase down the group in the periodic table. Zeff is determined by the number of protons in the nucleus (Z) and the shielding constant (S), which is associated with the number of electrons between the nucleus and the valence electrons. As you correctly mentioned, both Z and S increase down the group, resulting in the effective nuclear charge remaining relatively constant within a given group.

The increase in atomic radius down the group is primarily due to the addition of new electron shells as we move from one element to the next. This results in the valence electrons being further away from the nucleus and experiencing a weaker attraction to the positive charge. Hence, the atomic radius increases down the group.