Consider electron configuration as it applies to the periodic table and explain in detail how this accounts for the formation of ions.
Elements are arranged in the periodic table (see figure 1) according to their electronic configuration, which describes the number and arrangement of electrons in an atom, helping to make sense of the chemistry of an element. Electron configuration was first conceived of under the Bohr model of the atom, in which the electron shells were orbits at a fixed distance from the nucleus, and it is still common to speak of atmoic structures in terms of shells and sub-shells. An electron shell is the set of atomic orbitals which share the same prinicpal quantum number n (Krauskopf 1995) – the number before the letter in the oribital label. Orbitals are filled in the order of increasing n+1, where two orbitals have the same value of n+1, they are filled in order of increasing n. This gives the following order for filling the orbitals:
1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d 7p
If we look at this in relation to the periodic table we can see that starting with Hydrogen (which has one electron) going across to Helium having two. These two electrons have no filled up the first ‘shell’, called 1s. The next ‘shell’, 2s, is filled by Lithium and Beryllium. Moving across to the 2p sub shell (1p does not exist) – B, C, N, O, F, Ne. Each whole (one row in the table) ‘shell’ holds 8 electrons (with the S sub-shell holding 2 electrons, and the p sub-shell holding 6). So the second shell, looking at one row of the periodic table, is Li and Be (2s), and B, C, N, O, F, and Ne (2p). The elements of group 2 of the perdioic table have an electron configuration of [E]ns2 (where [E] is an inert gas configuration). Those elements grouped together in the periodic table have notable similarities in their chemical properties (Drever 1997).
Electrons fill energy levels according to the Aufbau principle – the principle that the electron configurations of atoms build up according to a set of rules. The three rules are that:
- Electrons go into the orbital at the lowest available energy level
- Each orbital can only contain at most two electrons (with opposite spins)
- Where there are two or more orbitals at the same energy, they fill singly before the electrons pair up.
Figure 1: Periodic Table of Elements
‘Valence electrons’ are the electrons contained in the outer shell (often referred to as the ‘valence shell’) of an atom, and are important in determining the chemical properties of an element (Krauskopf 1995). As a result of this, elements with the same number of valence electrons are grouped together in the period table. As a general rule, the fewer electrons an atom holds, the less stable it becomes and the more likely it is to react. Conversely the more complete the valence shell is the more inert an atom is and the less likely it is to chemically react. Elements in the same group of the periodic table have similar properties because they have the same outer electron configuration.
There are trends in properties down a group because of the shielding effect of the increasing number of inner full shells (Drever 1997). Electrons are able to move from one energy level to another by emission or absorption of a quantum of energy, in the form of a photon. It is this gain or loss of energy that can trigger an electron to move to another shell or even break free from the atom and it’s valence shell. When an electron absorbs/gains more photons, then it moves to a more outer shell depending on the amount of energy the electron contains and has gained due to absorption. When an election releases/loses photons, then it moves to a more inner shell depending on the amount of energy the electron contains and has lost.
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If we use fluorine as an example, we can see that the full electron configuration of Fluorine is 1s2s2p5 (F is 5th from the left in p-block, one behind Neon so has 5 2p electrons). The valence electrons are 2s2p5 as there are two shells and these electrons are in the outer one. The key point is that atoms like to have a whole shell of 8 electrons, as this makes them more stable. As we can see from figure 1, Fluorine has only 7 electrons (7th from the right on the second row). It really wants to gain an electron (to be like Neon) in order to have 8, and complete its shell. Fluorine is, therefore, very reactive and ‘steals’ and electron off anything it can find. When it does this it gains an electron and becomes a negative ion – F– (1s2s2p6). The reverse of this is Sodium (1s2s2p3s1), where 3s1 are the valence electrons. It really wants to loose this one extra electron to become 1s2s2p8 like Neon. It looses an electron and becomes a positive ion (Na+).
REFERENCES:
Baird, C. (1995) Environmental Geochemistry. USA: W.H. Freeman and Company
Drever, J.I. (1997) The Geochemistry of Natural Waters. London: Prentice-Hall
Krauskopf, K.B, Bird D.K. (1995) Introduction to Geochemistry. USA: McGraw-Hill
Howard A.G. (1998) Aquatic Environmental Chemistry. Oxford: Science Publications
Garrels, R. M., and J. C.Christ. (1965). Solutions, minerals, and equilibria. San Francisco: Freeman, Cooper.
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