Shielding Effect

[Qinyan Feng 1H]

Hello, I'm confused about the shielding effect. Can someone explain the concept of 'effective nuclear charge' and how to calculate it? Also how different orbitals have different shielding effect.
Thank you so much!

[Emily Ngo 1E]

Hi Qinyan,

Shielding is basically when electrons that occupy inner orbitals closer to the nucleus "shield" or block the electrons in outer orbitals. This will decrease the effective nuclear charge for electrons further away from the nucleus. For example, there is more shielding present in Strontium than Beryllium because there are more electrons between the valence electrons and the nucleus (bigger atom).

[Kathryn Heinemeier 3H]

so basically effective nuclear charge is Z-the shielding constant. when the energy level is closer to the nucleus it will have less shielding versus an energy level that is farther away and therefore has more distance and electrons shielding between.

[Jane Wang 1E]

Hi! Effective Nuclear Charge is like the charge from electrons that can really interact with the positive charge of the nucleus. Shielding effect is that inner electron will shield the interaction between outer electrons and the nucleus, and the shielding effect need to be deducted from the total nuclear charge to get the effective nuclear charge.

[Janice Hu 2L]

Effective nuclear charge (Z eff) = the positive charge of the nucleus that actually interacts with an electron

Z eff = Z (total nuclear charge, or number of protons in the nucleus) - S (shielding constant)

Shielding effect: electrons closer to the nucleus repel electrons further away from the nucleus, causing the nucleus to have less attraction to the outer electrons and hence lowering the effective nuclear charge on those electrons

You can think of the shielding effect as closer electrons "shielding" outer electrons from experiencing the full attractive positive charge of the nucleus :)

In general, orbitals with greater electron density closer to the nucleus have greater shielding power (meaning they are less shielded, have a greater effective nuclear charge, and can shield other electrons further away).

I think the trend for how much shielding electrons in an orbital experience follows the order of orbitals in electron configurations (or the diagonal rule):
1s < 2s < 2p < 3s < 3p < 4s < 3d < etc. where 1s experiences less shielding than 2s and so on

[Brooke Gushiken 1B]

Hi! From what I understand, shielding is when electrons that are in energy levels closer to the nucleus repel electrons in energy levels that are further away from the nucleus. This affects the effective nuclear charge, which is how much the nucleus actually interacts with any given electron. The further away from the nucleus the electron is, the more shielding it will experience, therefore lowering its effective nuclear charge.

[Triston Dinh 1D]

The shielding effect refers to the attraction between an electron and the nucleus in an atom. The effective nuclear charge is the net charge of a valence electron. As you go from left to right across the periodic table, the effective nuclear charge of an electron decreases. This is because the inner electrons "shield" the positive charge from the nucleus and thus the outer electrons will have a lower effective nuclear charge.

[Russell Chuang 1J]

The shielding effect is when electrons closer to the nucleus block further electrons and thereby decrease attraction. Because of this, valence shell electrons are further away from the nucleus and the larger the atomic radius is. Nuclear charge is the said attraction between the nucleus and the electrons.

[Emily Wan 1l]

Effective nuclear charge refers to the net amount of positive attraction that exists between the nucleus and valence electrons, after subtracting the shielding effect of inner electrons.