Electron Configuration Exceptions

[eurekameh]

This is the electron configuration for Ag (Silver) found on the wikipedia page:
1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 4d10 5s1

I used the Aufbau Principle

and got this instead:
1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 5s2 4d9

1. I can see that Ag is an exception to the rule, but are there any generalizations I can make so that I don't write down the wrong configurations just because of an exception?

2. Also, is it traditional to write these in order of increasing quantum number (just like it's written above), because the Aufbau Principle is meant for the orbitals to be filled in increasing energy?

For example, Br is written, on wikipedia, as 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p5, but written using Aufbau Principle, it's actually
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p5. Which one is more correct?

 

[AGNuke]

1.) In case of configuration like (n-1)d⁴ ns² OR (n-1)d⁹ ns², to attain a more stable configuration, the electron from s orbital jumps to the previous d orbital. So, the new configuration becomes (n-1)d⁵ ns¹ and (n-1)d¹⁰ ns¹ respectively.

This phenomenon is called pseudo-inert effect, and according to this, half filled or fully filled orbitals cluster (like 2p, 3d orbitals) are more stable, and if this is attainable by one electron jump, most atoms are ready to do so. Example are Chromium Group (Group 6) and Copper Series (Group 11).

2) It is OK if they clubbed the orbitals on the basis of their respective Principle Quantum number or by Aufbau's Principle, as long as it do not look absurd. But conventionally, people fill the orbitals with increasing Principle Quantum number, rather than by the diagram. Diagram is useful for allotting electrons, but writing configuration that way may seem confusing.

 

[eurekameh]

One more thing: Why is filling or half-filling a p or d orbital more stable than filling or half-filling an s orbital?

 

[AGNuke]

Probably owing to the fact that s orbital is a single orbital, one can half-fill it and second can fully. Moreover, p and d orbitals are clusters of 3 and 5 orbitals respectively, so they are stable if each orbital possesses equal amount of electron.

 

[JohnRC]

This question (original post) is uncovering a lot of grey areas and textbook misinformation and misunderstandings in the area of atomic structure. I will address several of these:

First, the convention of writing electron configurations is always done principal quantum number first, regardless of relative energies.

Thus the electron configuration of neodymium is written either

Nd: 1s²2s²2p⁶3s²3p⁶3d¹⁰4s²4p⁶4d¹⁰4f⁴5s²5p⁶6s²

or Nd: {Xe}4f⁴6s²

The electronic structure of silver is written

Ag: {Kr}4d¹⁰5s¹, or its longer-winded equivalent.

Second, "Aufbau" is not the name of a German physicist. It is actually a German word, that does not have an exact English equivalent. It describes the process of building up something piece by piece on the foundation of the simpler structure that has been arrived at by placement of the previous piece. Like the process of building up a brick wall by placing bricks in the structure one at a time. The Aufbau principle (never "Aufbau's principle") refers to obtaining the electron configuration for nitrogen by looking at the previous element C: 1s²2s²2p², and saying "nitrogen needs an extra electron. Where can we put it?" to arrive at N: 1s²2s²2p³.

Third, it is not the case that the periodic table groups elements with similar valence electron configuration. The classic example is in group 10, which contains
Ni: {Ar} 3d⁸4s²
Pd: {Kr} 4d¹⁰
Pt: {Xe} 5d⁹6s¹

Fourth, it is not the case that an electron configuration with a half-filled d or f subshell has a lower energy than one that is one electron away from such a state. There is a second factor involved: Spin-Orbit coupling. Russell Saunders coupling is a model for this factor. It shows that electron configurations like those above can split into several different spectroscopic terms (electronic energy levels), with quite widely separated energies. The "electron configuration" that is quoted for each element in a periodic table is the one associated with the spectroscopic term that happens to be associated with the ground electronic state. The lowest energy electron configuration averaged over all of the spectroscopic terms associated with it can often come out quite differently from the actual electron configuration associated with the lowest energy spectroscopic term.