Electronic Structure

1. In the currently accepted model of the atom, electrons have fixed energies. They move around the nucleus in certain regions of the atom called shells or energy levels.

2. Each shell is given a number called the principal quantum number. The further a shell is from the nucleus, the higher its energy and the larger its principal quantum number.

3. This model helps to explain why electrons are attracted to the nucleus, but are not drawn into it and destroyed.

4. Experiments show that not all the electrons in a shell have exactly the same energy. The atomic model explains this- shells are divided up into sub-shells that have slightly different energies. The sub-shells have different numbers of orbitals which can hold up to 2 electrons.

1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p.... (s=2, p=6, d=10)

5. The two electrons in each orbital spin in opposite directions.

You can figure out most electronic configurations pretty easily, so long as you know a few simple rules-

1. Electrons fill up the lowest energy sub-shells first.

E.g. The electronic configuration of calcium is:

1s2 2s2 2p6 3s2 3p6 4s2

2. Electrons fill orbitals singly before they start sharing.

3. For the configuration of ions from the s and p block periodic table, just remove or add the electrons to or from the highest energy occupies sub-shell.

E.g. Mg2+= 1s2 2s2 2p6, Cl- = 1s2 2s2 2p6 3s2 3p6

Watch out! noble gas symbols, like that of argon (Ar), are sometimes used in electron configurations. For example, calcium (1s2 2s2 2p6 3s2 3p6 4s2) can be written as [Ar] 4s2 , where [Ar] = 1s2 2s2 2p6 3s2 3p6.

1. Chromium (Cr) and copper (Cu) are badly behaved. They donate one of their 4s electrons to the 3d sub-shell. It's because they're happier with a more stable full or half-full d sub-shell.

Cr atom (24e-): 1s2 2s2 2p6 3s2 3p6 3d5 4s1        Cu atom (29 e-): 1s2 2s2 2p6 3s2 3p6 4s1 3d10

2.And here's another weird thing about transition metals- when they become ions, they lose their 4s electrons before their 3d electrons.

Fe atom (26e-): 1s2 2s2 2p6 3s2 3p6 3d5 4s2 --> Fe3- ion (23 e-): 1s2 2s2 2p6 3s2 3p6 3d5

The number of outer electrons decides the chemical properties of an element.

1. The s block element (Groups 1 and 2) have 1 or 2 outer shell electrons. These are easily lost to form positive ions with an inert gas configuration

E.g Na- 1s2 2s2 2p6 3s1--> Na+ - 1s2 2s2 3p6 (the electronic configuration of neon)

2. The elements in Groups 5, 6 and 7 (in the p block) can gain 1, 2 or 3 electrons to form negative ions with an inert gas configuration. E.g. O- 1s2 2s2 2p4 --> O2- - 1s2 2s2 2p6

3. Group (the inert gases) have completely filled s and p sub-shells and don't need to bother gaining, losing or sharing electrons- their full sub-shells make them inert.

4. The d block elements (transition metals) tend to lose s and d electrons to form positive ions.