CHEMISTRY COMMUNITY

Hi!

I'm a bit confused about effective nuclear charge and the trends in the periodic table regarding the effective nuclear charge, and shielding. How does the effective nuclear charge relate to (if it relates to) shielding?

Effective nuclear charge is the net nuclear charge after taking into account the shielding caused by other electrons in the atom. Nuclear charge and shielding are related by the equation Z(eff)=Z-S, where Z(eff) is the effective nuclear charge, Z is the atomic number, and S is the number of shielding electrons. Effective nuclear charge increases across a period and decreases down a group.

Effective nuclear charge can tell us a lot about periodic trends. For example, as you move across a period, the effective nuclear charge increases as the number of protons increases while the relative number of shielding electrons stays the same. This signifies the notion that the proton's pull is becoming proportionally greater to the repulsion of the electrons. As a result, due to the increasing pull of the positive nucleus on the valence electrons of the atoms, atoms tend to decrease in size as you move from left to right of a period. In addition, the shielding effect is related to the idea that "core electrons", electrons that are not valence electrons, counteract the nucleus's attraction to the valence electrons because of their electron to electron repulsion.

Can you help explain why it would decrease down a group?

Effective nuclear charge decreases as you go down a group because you are adding orbitals as you go down, in other words the electron is being added into a new orbital that is farther away from the nucleus. This increases the shielding because there are more electrons in between the nucleus and the electron in question.