Consequently, the most electronegative element is fluorine, and the least electronegative (ignoring the unstable francium) is caesium.
The graph below shows how electronegativity varies with atomic number.
Ionisation Energies
The ionisation energy (ionisation enthalpy) of an element is the amount of energy needed to remove an electron from the element to an infinite distance. This energy depends on how strongly the electron is attracted to the nucleus. Strictly speaking, we should refer to the first ionisation energy, the second ionisation energy, etc., where the first ionisation energy is the energy needed for
E → E+
and the second ionisation energy is the energy needed for
E+ →• E2+
When we refer to "ionisation energy" we usually mean first ionisation energy.
As we move along a row of the Periodic Table, the nuclear charge increases. This holds all the electrons more strongly, and the ionisation energy increases. As we start a new period, the electrons in the filled inner shell are able to shield the outer electrons from the nuclear charge, and so there is a large decrease in ionisation energy. The ionisation energy then starts to increase again as we move along the row. Consequently, the lowest ionisation energies are found in the bottom right portion of the table.
If we plot ionisation energy against atomic number we get the following graph: \\
Atomic Size
All atoms contain a nucleus surrounded by electrons. The nucleus is very small compared with the overall size of the atom (about 10-12 of the overall volume). Adding more protons and neutrons to the nucleus makes very little difference to the size of the nucleus. As more protons are added, more electrons are added to balance the charge. This makes a considerable difference to the overall size of the atom. To understand this difference, it is necessary to consider the effects of charges on the nucleus and the electrons.
As more protons are added, the nuclear charge increases. This pulls the electrons in closer to the nucleus. Consequently, the atomic radius decreases as we move along a period. At the start of each new period there is a filled inner shell of electrons between the nucleus and the outer electrons. These inner electrons shield the outer electrons from the nuclear charge, and so as we go down any group the atomic radius increases.
When we plot atomic radius against atomic number we get the graph below.
It can be seen from the graph that fluorine has the smallest atomic radius and caesium the largest. (Francium is not stable enough to measure its size.)
Ionic radius also varies as we move along the group, although in a less obvious manner. Where elements lose electrons to form cations, the radius decreases along the group. However, there is then an increase in radius when we get to atoms that gain electrons to form anions. This is followed by a decrease in radius as we continue moving along the group, due to the increasing nuclear charge. For example, sulphur has atomic radius 104 pm and the S2- ion has radius 184 pm, while chlorine has radius 99 pm and Cl- has radius 181 pm.
