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Mahir Pepic 3F
Posts: 15
Joined: Fri Sep 25, 2015 3:00 am


Postby Mahir Pepic 3F » Sat Oct 10, 2015 4:14 pm

So basically:
The effective nuclear charge of an electron in a p orbital would be lower than that of one in an s orbital because the s orbital is shielding the p orbital in that situation? And it would be even lower in a d orbital because it has the s orbital AND the p orbital shielding it? Am I on the right track here?

Annie Qing 2F
Posts: 28
Joined: Fri Sep 25, 2015 3:00 am

Re: Shielding

Postby Annie Qing 2F » Sat Oct 10, 2015 4:35 pm

From what I understand, shielding is not a phenomenon that occurs as orbitals change but as the number of shells (quantum number n) changes. So, as you go down the periodic table, the effective charge on an electron decreases because the shells are increasingly farther away from the nucleus and closer shells "shield" outer electrons.

This can be demonstrated by Coulomb's Law Equation

where distance d is inversely proportional to the electrical force between two charged objects. When distance increases (or as you get farther from the nucleus and n increases), the electrical force between the protons in the nucleus and an electron decreases.

Daniel Fu 2K
Posts: 10
Joined: Fri Sep 25, 2015 3:00 am

Re: Shielding

Postby Daniel Fu 2K » Sat Oct 10, 2015 5:53 pm

Starting from the basics, effective nuclear charge [also written as Zeff (eff is subscript)]is the measurement of attractive forces between protons in the nucleus and the electrons in the valence shell. in order to calculate this value we have to use the following equation of Zeff=Z(#of protons)-[#of non-valence electrons]. Using this equation we can calculate the effective nuclear charge for any element. Below, I will give you examples of how to calculate the Zeff. I will show you the trends in the process as well.
Examples (going across the element period, aka moving from left to right on the periodic table):
Sodium: Zeff=11-10=+1
Magnesium: Zeff=12-10=+2
Aluminum: Zeff=13-10=+3

As you can see, when we go across the periodic table from left to right, the effective nuclear charge increases.

Examples (going down element group, aka moving from top to down):
Sodium: Zeff=11-10=+1
Potassium: Zeff=19-18=+1
Rubidium: Zeff=37-36=+1

So in effect, the Zeff for anyone group will be the same going down, however, since Rubidium is father away than Potassium which is farther away than Sodium from the nucleus, the shielding effect comes into play as Annie mentioned. Which means that although the calculated Zeff is the same, there will be less attraction between the valence electrons from the nucleus the further you go down any group of elements.

Hope this helps. I'm pretty rusty at the moment, so if I botched anything up or didn't explain something clear enough, please let me know :)

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