Energy Levels or Shells
The simplest model of electrons has them orbiting in shells around the nucleus. Each successive shell is further from the nucleus and has a greater energy.
Sub Shells and Orbitals
This model can be further refined by the concept of sub shells and orbitals.
Sub shells are known by letters s, p, d, and f. The s sub shell can contain 2 electrons, p 6, d 10 and f 14.
Electrons occupy negative charge clouds called orbitals, each orbital can hold only 2 electrons. Each type of shell has a different type of orbital.
How we write electron configurations
Electrons fill the lowest energy level first this means it is generally easy to predict how the electrons will fill the orbitals (it gets more complicated with the transition metals).
Let’s look at building up the electronic arrangement (electron configuration) from hydrogen (Z = 1) as far as krypton (Z = 36).
The order of filling of orbitals in atoms (it is different for positive ions) is:
1s 2s 2p 3s 3p 4s 3d 4p (notice how the 3d comes before 4p)
By going through the elements in sequence, we will see the various rules come into play.
Hydrogen (Z = 1) It has one electron. It must be placed in the orbital with the lowest energy.
Helium (Z = 2) It has two electrons. There is room in the 1s orbital for a maximum of two electrons but they must have opposite spins. Electron spin is a difficult concept but think of it as being no more than a label of "up-spin" or"down-spin".
Lithium (Z = 3) It has three electrons. Although two can go into the 1s orbital, the third one must be placed in the 2s.
Beryllium (Z = 4) It has four electrons. The fourth can also go in 2s (but remember about the opposing spins)
Boron (Z = 5) It has five electrons. The fifth one must go in 2p. There are three 2p orbitals available and these are identified by their direction (px, py, pz). It doesn’t matter which of these is chosen when there is only one electron in one of these orbitals.
Carbon (Z = 6) It has six electrons. The sixth one must go in a different 2p orbital. We now need to specify the particular 2p orbitals we use. It doesn’t matter which combination of x, y, or z but they must be different.
Nitrogen (Z = 7) It has seven electrons. We now need the remaining 2p orbital.
Oxygen (Z = 8) It has eight electrons. We can now complete the filling of the 2p orbitals in turn. Again it doesn’t matter which of the orbitals is completed first.
Fluorine (Z = 9) It has nine electrons.
Neon (Z = 10) It has ten electrons.
or since all of the 2p orbitals are completed we can simply write: 1s22s22p6
Sodium (Z = 11) We now have no space in the second energy level and so have to start on the third.
This can be simplified (unless you are told to write the complete configuration) to [Ne]3s1
Magnesium (Z = 12) to argon (Z = 18) We continue through the whole of the period using the same concepts as we did from lithium to neon.
Magnesium is [Ne]3s2Argon is [Ne]3s23p6
Potassium (Z = 19) As before, we start on the fourth energy level because that is the next in the order of filling.
Calcium (Z = 20) This is where you stopped for GCSE
Scandium (Z = 21) We start on the set of 3d orbitals...
and continue as expected (without worrying about the opposing spins this time) until we come to…
Chromium (Z = 24)
We might expect [Ar]4s23d4 but instead we get [Ar]4s13d5
This is because of the special stability of the half-filled set of 3d orbitals. An electron promoted from the 4s to the 3d. Although this requires energy, this is paid back by the extra stability.
Manganese (Z = 25) The gap in 4s is filled again
Iron (Z = 26), Cobalt (Z = 27) and Nickel (Z = 28) proceed onwards as expected
[Ar]4s23d6 , [Ar]4s23d7 , [Ar]4s23d8
Copper (Z = 29) adopts the unexpected (like chromium) because of the special stability of the full 3d.
Zinc (Z = 30) The gap in 4s is filled again
Continue by filling the 4p orbitals (opposing spin is important again) until we get to
Krypton (Z = 36)