Hodder CIE revision guide 2010 Chemistry fig 2.1 29 July 2010 Eleanor Jones
2 Atomic structure
As more electrons are added, they go into orbitals of increasing energy:
4p 3d 4s
4s is lower than 3d because electrons are, on average, closer to the nucleus
3p 3s
Put one in each, then pair up
2p 2s 1s
Each orbital holds two electrons Spin
Figure 2.1 Sequence of filling orbitals with electrons
Figure 2.1 illustrates some key points in the arrangement of electrons in atoms. These are things you should remember: ●●
The electrons are arranged in energy levels (or shells) from level 1, closest to the nucleus. On moving outwards from the nucleus, the shells gradually increase in energy.
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Most energy levels (except the first) contain sub-levels (or sub-shells) denoted by s, p and d.
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Different sub-levels contain different numbers of orbitals, with each orbital holding a maximum of two electrons.
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When filling up the energy levels in an atom, electrons go into the lowest energy level first.
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In sub-levels containing more than one orbital, each of the orbitals is filled singly before any are doubly-filled.
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If you study Figure 2.1 there is one strange entry — the 4s-orbital has lower energy than the 3d-orbital.
In an examination, you may be asked to deduce the electron configuration of an atom (or ion) given its proton number (and any charge). The following examples show how to do this.
Example 1 An atom, X, has a proton number of 16. Deduce the electron configuration of this atom. Answer The proton number is 16, so the atom must also contain 16 electrons. Referring back to Figure 2.1, we can count upwards until 16 electrons have been used. This means the 1s-orbital, the 2s- and 2p-orbitals and the 3s-orbital
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