Electron configurations Today we are going to look at how we fit electrons into orbitals

You can think of an atom as a very bizarre house (like an inverted pyramid!) - with the nucleus living on the ground floor, and then various rooms (orbitals) on the higher floors occupied by the electrons. On the first floor there is only 1 room (the 1s orbital); on the second floor there are 4 rooms (the 2s, 2p x, 2p y and 2p z orbitals); on the third floor there are 9 rooms (one 3s orbital, three 3p orbitals and five 3d orbitals); and so on. But the rooms aren't very big... Each orbital can only hold 2 electrons.

A convenient way of showing the orbitals that the electrons live in is to draw "electrons-in- boxes"

Orbitals can be represented as boxes with the electrons in them shown as arrows. Often an up- arrow and a down-arrow are used to show that the electrons are in some way different.

A 1s orbital holding 2 electrons would be drawn as shown on the right, but it can be written even more quickly as 1s 2. This is read as "one s two" -not as- -"one s squared".

You mustn't confuse the two numbers in this notation:

The order of filling orbitals

Electrons fill low energy orbitals (closer to the nucleus) before they fill higher energy ones. Where there is a choice between orbitals of equal energy, they fill the orbitals singly (with one electron each) as far as possible. Example 3s orbitals

This filling of orbitals singly where possible is known as Hund's rule. It only applies where the orbitals have exactly the same energies (as with p orbitals, for example), and helps to minimise the repulsions between electrons and so makes the atom more stable.

This diagram (not to scale) summarizes the energies of the orbitals up to the 4p level.

Notice that the s orbital always has a slightly lower energy than the p orbitals at the same energy level, so the s orbital always fills with electrons before the corresponding p orbitals.

The real oddity is the position of the 3d orbitals. They are at a slightly higher level than the 4s - and so it is the 4s orbital which will fill first, followed by all the 3d orbitals and then the 4p orbitals.

Similar confusion occurs at higher levels, with so much overlap between the energy levels that the 4f orbitals don't fill until after the 6s, for example.

For exam purposes, you simply have to remember that the 4s orbital fills before the 3d orbitals. The same thing happens at the next level as well - the 5s orbital fills before the 4d orbitals. All the other complications are beyond the scope of this class and my brain power

Electron configurations Knowing the order of filling is central to understanding how to write electron configurations

Relating orbital filling to the Periodic Table

The first row Hydrogen has its only electron in the 1s orbital - 1s 1, and at helium the first level is completely full - 1s 2.

The second row Now we need to start filling the second level, and hence start the second row. Lithium's electron goes into the 2s orbital because that has a lower energy than the 2p orbitals. Lithium has an electronic structure of 1s 2 2s 1. Beryllium adds a second electron to this same level - 1s 2 2s 2.

Now the 2p levels start to fill. These levels all have the same energy, and so the electrons go in singly at first. B=1s 2 2s 2 2p x 1 C=1s 2 2s 2 2p x 1 2p y 1 N=1s 2 2s 2 2p x 1 2p y 1 2p z 1

The next electrons to go in will have to pair up with those already there. O = 1s 2 2s 2 2p x 2 2p y 1 2p z 1 F = 1s 2 2s 2 2p x 2 2p y 2 2p z 1 Ne = 1s 2 2s 2 2p x 2 2p y 2 2p z 2

You can see that it is going to get progressively tedious to write the full electronic structures of atoms as the number of electrons increases. There are two simple shortcut to make our job easier.

Shortcut #1 You can lump all the inner electrons together using, for example, the symbol [Ne]. In this context, [Ne] means the electronic structure of neon - in other words: 1s 2 2s 2 2p x 2 2p y 2 2p z 2 You wouldn't do this with helium because it takes longer to write [He] than it does 1s 2. On this basis the structure of chlorine would be written [Ne]3s 2 3p x 2 3p y 2 3p z 1.

Shortcut 2: All the various p electrons can be lumped together. For example, fluorine could be written as 1s 2 2s 2 2p 5, and neon as 1s 2 2s 2 2p 6

The third row At neon, all the second level orbitals are full, and so after this we have to start the third row with sodium. The pattern of filling is now exactly the same as in the previous row, except that everything is now happening at the 3-level.

example Mg =1s 2 2s 2 2p 6 3s 2 or shortcut version [Ne]3s 2 Ar = 1s 2 2s 2 2p 6 3s 2 3p 6 or [Ne] 3s 2 3p 6

The beginning of the fourth row At this point the 3-level orbitals aren't all full - the 3d levels haven't been used yet. But if you refer back to the energies of the orbitals, you will see that the next lowest energy orbital is the 4s - so that fills next.

examples K = 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 Ca = 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2

There is strong evidence for this filling pattern in the similarities in the chemistry of elements like sodium (1s 2 2s 2 2p 6 3s 1 ) and potassium (1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 The outer electron governs their properties and that electron is in the same sort of orbital in both of the elements. That wouldn't be true if the outer electron in potassium was 3d 1.

The elements in group 1 of the Periodic Table all have an outer electronic structure of ns 1 (where n is a number between 2 and 7). All group 2 elements have an outer electronic structure of ns 2. Elements in groups 1 and 2 are described as s-block elements.Elements from group 3 across to the noble gases all have their outer electrons in p orbitals. These are then described as p-block elements.

Remember that the 4s orbital has a lower energy than the 3d orbitals and so fills first. Once the 3d orbitals have filled up, the next electrons go into the 4p orbitals as you would expect. d-block elements are elements in which the last electron to be added to the atom is in a d orbital. The first series of these contains the elements from scandium to zinc, which are called transition elements or transition metals.

Summary Writing the the electron configuration of an element from hydrogen to krypton Use the Periodic Table to find the atomic number, and hence number of electrons. Fill up orbitals in the order 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p until you run out of electrons. The 3d is the awkward one remember that specially. Fill p and d orbitals singly as far as possible before pairing electrons up.