Electron Configurations

The Pauli exclusion principle says that all electrons in an atom have to have a unique set of quantum numbers. NO duplicates! It's like a serial number for electrons, except we use n, ℓ, m, and ms.

The aufbau principle tells us to "build up" from the bottom of the energy well to the top. Pour water in a bucket and it fills from the bottom up - same idea. We will completely fill a lower level of energy before we advance to the next higher level. Here is the order of filling for all the orbitals in the atom.

1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d 7p

Oh my, how are you going to remember all that? Well, the best way to memorize the aufbau filling order is to use a periodic table and know how the orbitals "fit" the table. You can see the filling order below.

Now it's not so bad. Next thing to remember is how many electrons go into each orbital set and also the ordering of spin within a set.

Filling an Electron Energy Level Diagram

Once you memorize and can use the periodic table to help you get the correct order of orbitals, you then need to know how to fill those orbitals. The orbitals themselves are shown on an energy diagram as blanks and we will put in up arrows ↿ and down arrows ⇂ to represent the spin quantum numbers +½ and –½. We do this following some rules.

  1. Aufbau Principle - always take the lowest possible energy level and fill it before going up to the next level
  2. Hund's Rule - on any degenerate levels (same energies), always fill singly with up arrows ↿ (+½) before you then pair with the down arrows ⇂ (–½)
    note: electrons with matching spin states are said to have parallel spin states
  3. Pauli Exclusion Principle - no two electrons can have the same set of four quantum numbers - these means no more that 2 electrons per orbital (blank) and when there are two, they have opposite spins ⥮

Here is a nice pdf for you to use to practice dropping in electrons (arrows) in the right order. Its an Electron Energy Diagram (Aufbau Filling Order).

Below is an animated gif showing the proper filling for the 2nd row of the periodic table. It illustrates the rules of filling.



Example Question (Part 1)

Part 1: What is the electron configuration of antimony (Sb)?

Answer: Find antimony on the periodic table - it is element #51 located in group 15 (p-group) on row 5. So this means we will be filling everything up to the 5p level and then partially fill the 5p level. Sb is the 3rd element in the 5p group or set therefore you will get the following electron configuration.

[Sb] = 1s2 2s22p6 3s23p6 4s23d104p6 5s24d105p3

and, you can also take the shortcut method of starting with the previous noble gas as your starting point which in this case is krypton (Kr).

[Sb] = [Kr]5s24d105p3

Example Question (Part 2)

How many unpaired electrons are in antimony?

Answer: THREE. We know the electron configuration and that all the orbital sets are full (all electrons paired) through 4d. It is the 5p level that is partially full. Only 3 electrons go into the set - so all three go in as single UNpaired electrons with parallel spin states (all +½ or ↿). Specifically... 5p =

Below is the entire electron energy diagram for antimony, Sb.

5p
4d
5s
4p
3d
4s
3p
3s
2p
2s
1s



⬇︎ NEW STUFF HERE ⬇︎... Added this chunk on isoelectronic species on 10/1/20. I talked about this in class on 10/1 and said I'd write something. This is that something. Probably will make it its own section in the future. - Dr. McCord


Borrow, Steal, Ditch, and Share

So we've now got the whole electron configuration thing down, right? Every neutral element has a unique set of electrons and every element has the same number of electrons as there are protons in the nucleus.

Electron Envy

Group 18 on the periodic table (the noble gases) is truly the envy of all the other elements. That last group of the periodic table has a perfectly filled valence shell. To be specific, they all have full s and p orbital sets (s2p6) at every main level n. Yes, yes... row 1 only has to fill the 1s orbital and it is "full" as well (no p electrons there).

The thing is, one of the biggest driving forces in all of chemistry is the desire of the elements to get to this filled valence shell of s2p6... a total of 8 electrons. In the next chapter, we will discuss bonding and we will try to follow the octet rule whereby we transfer/donate/take/share electrons between atoms in order to best get to 8 electrons for everybody.

Me Want Electrons - Anions

The non-metals on the periodic table all tend to be very close to making it to 8 electons or what we call a full octet. They just fall a little short. About 1, 2, or 3 borrowed (or stolen, or found) electrons will fix these guys up in no time. All the group 17 elements (the halogens) just want one electron to get to the promise land of 8 electons total. So they are so close they can taste it - this means that they desire the electrons the most and will have the biggest pull on stray or any available electrons when compared to others in their row (aka period). The result? They all acheive –1 status for charge after they get that extra electron. We write this new anions as F (fluoride), Cl (chloride), Br (bromide), and I (iodide). Please forgive me for ignoring At (astatine) and Ts (tennessine)... It's my book and I get to ignore many of the lesser elements of the last 2 rows on the periodic table. Anyway, all of those *-ides together are known as the halides and all of them have noble gas electron configurations... with the business end of that being ns2np6 where n is 2-5 for our four favorite halides.

Isoelectronic means Same Electron Configuration

So fluoride is isoelectronic with neon. Chloride is isoelectronic with argon. Bromide is isoelectronic with krypton. And, iodide is isoelectronic with xenon. Cool. Just remember though that each those halogens have one LESS proton than the noble gas it is next to and therefore is a –1 anion because of the slight excess of one electon.

The same game now is played with group 16 - except that now you need 2 more electrons to reach noble gas configuration. Picking up 2 extra electrons will lead to a –2 charge for the resulting anion. The naming is the same though. Take the stem (the root word) of the element and change the suffix to -ide. So oxygen becomes oxide (O2–), sulfur becomes sulfide (S2–), selenium becomes selenide (Se2–), and last, tellurium becomes telluride (Te2–)... which I hear the skiing is great there (joke... LOL or better ROTFLMAO). Fun stuff there. Play the game again with group 15 and you'll get a series of –3 anions with the main dudes being nitride (N3–) and phosphide (P3–).

Electrons? We don't need no stinkin' electons - Cations

Yeah, that is pretty much what most metals say. They are the unfortunate lot that "went to far" on the electon count past noble gas configurations. So what do they do? They ditch the extra electrons and "backup" to the previous noble gas configuration. So sodium ditches one electon and now has 10 electons total and is isoelectronic with neon (10 total electons), AND is is a 1+ cation now (Na+). The thing about cations and naming - they have the same name as the metal. So Na+ is called sodium (ion). Calcium, like all the group 2 metals, ditches 2 electrons to make calcium ion (Ca2+) and is isoelectronic with argon (18 total electons). To get +3 cation status, you need to lose 3 electrons... aluminum is great at this and goes to Al3+.

And Finally...

Here is an entire series of isoelectronic species. All of the following ions have the same exact electron configuration which happens to match up with argon.

P3–
S2–
Cl
Ar
K+
Ca2+
Ga3+

1s22s22p63s23p6



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