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- Ionization Energy (IE):- also known as the ionization potential , Is the energy required to remove an electron from a gaseous atom or ion. The losing of an electron is an Endothermic process (require energy).
*Where n = 0 (first ionization energy), n = 1 (second ionization energy), and so on. As would be expected from the effects of shielding, the ionization energy varies with different nuclei and different numbers of electrons.
a- Size of the atom - IE decreases as the size of the atom increases,
thus IE decreases down the group.
b- IE increase from left to right in every period because of the
increasing Zeff.
The following table shows the first, second and (some higher) ionization energies of the elements from Na to Ar in KJ/mol unit.
First ionization energies vary systematically through the periodic table (Table below), being smallest at the lower left (near Cs ) and greatest near
the upper right (near He ). The variation follows the pattern of effective nuclear charge, and (as Z eff itself shows) there are some subtle modulations arising from the effect of electron–electron repulsions within the same subshell. A useful approximation is that for an electron from a shell with principal quantum number n
I.E= 13.6 e.v [(Z)^2 / n^2 ]
Z* = Effective nuclear charge
n = principle quantum number
13.6 e.v= the ionization energy of the H atom is 13.6 eV, so to remove an electron from an H atom is equivalent to dragging the electron through a potential difference of 13.6 V.
Ex: Calculate the first ionization energy or (potential) of (^) 3 Li
Ans :
First we must calculate Effective nuclear charge (Zeff or Z*)
3 Li 1S^2 2S^1
σ = (0 * .35) + (2 * 0.85) = 1.
Zeff = Z – σ = 3 – 1.7 = 1.
So I.E = 13.6 e.v [(Z) 2 / n^2 ]
= 13.6 e.v * [(1.3)^2 / (2)^2 ] = 13.6 e.v * [0.4225] = 5.746 e.v
So IE of 3 Li in KJ/ mol = [5.746] * [96.485]
= 554.4 KJ/ mol
On moving across a period , the atomic size decreases and hence the force of attraction exerted by the nucleus on the electrons increases. Consequently, the atom has a greater tendency to attract additional electron i.e., its electron affinity increases.
EA values of metals are low while those of non-metals are high. *Halogens have high electron affinities. This is due to their strong tendency to gain an additional electron to change into the stable ns^2 np^6 configuration.
On moving down a group , the atomic size increases and therefore, the effective nuclear attraction decreases and thus electron affinity decreases.
The next table represents the affinities values of main group elements in KJ/mol.
* Electronegativity ;- Pauling defined electronegativity as the power of
an atom in a molecule to attract shared electrons to itself.
*Others definitions;
** Electronegativity is a measure of an atom’s ability to attract electrons from a neighboring atom to which it is bonded.
** Electronegativity is the ability of an atom to win the competition to attract shared electrons.
- This relative attraction for bonding electron pairs really reflects the comparative Zeff of the two atoms on the shared electrons. Thus, the values increase from left to right across a period and decrease down a group in the same way as ionization energies do.
*Some useful electronegativity values are shown in Figure below:
*In most cases the different methods give similar electronegativity values, sometimes with the exception of the transition metals. We choose to use the values reported by Linus Pauling, Mann, Meek, Allen, and others.
