Periodic Properties: Ionization Potential

Definition of ionization potential:

The amount of energy needed to remove the most loosely bound electron from isolated neutral gaseous atom from its gaseous state to convert it into a unipositive gaseous cation.

Unit:

1. eV/atom ( 1 eV = 96.4705 KJ/mole = 23.06 Kcal/mole)
2. Kcal/mole ( 1 Kcal/mole = 4.185 Kj/mole)
3. KJ/mole

Factors effecting the magnitude of Ionization Potential

1. Effective nuclear charge: Greater is the magnitude of effective nuclear charge, greater is the electrostatic force of attraction exerted by the nucleus on the outer electrons. Thus with the increase of the magnitude of effective nuclear charge, the magnitude of ionization potential also increases.
2. Size of the atom (r): Greater is the magnitude of r of an atom, farther is the outer-most shell electron from the nucleus. Thus with the increase of r, it would be more and more easy to remove an electron from the outermost shell.
3. Principal quantum number (n): Grater is the value of n for the valence-shell electron of an atom, farther away this electron will be from the nucleus, and hence lesser will be the force of attraction exerted by the nucleus on it. This means that a lesser amount of energy will be required to remove the valence shell electron.
4. Half-filled and full-filled orbitals: According to Hund’s rule, half-filled ( ns¹, np³, nd⁵) or full-filled (ns², np⁶, nd¹⁰) orbitals are more stable, and hence more energy is required to remove an electron from such orbitals.
5. Shielding effect: Greater is the magnitude of shielding effect working on the valence-shell electron, smaller is the magnitude of the force of attraction between the nucleus and valence shell electron, and hence lower will be the energy to remove the outermost electron.

Trend in a Period:

In general as we move from left to right in a period, the ionization potential of the elements increases due to the successive increase in the nuclear charge (i.e. atomic number)and decrease in atomic size ( some elements show irregular trends).

First period:

Value of 1st ionization potential of H = 1320.0 KJ/mole but this value for He is 2372.3 KJ/mole. In case of He both the electrons are present in the same orbital (1s2) , the first electron is not able to shield the second. This increase the ionization potential of He atom.

2nd period:

On passing from Li to Ne, generally, the ionization potential increases with the increase of effective nuclear charge and decrease of atomic size. Be and N show irregular trends in this period.

In the case of Be (1s2 2s2), it is more difficult to remove an electron from the completely filled 2s-orbital while in the case of B (1s²2s²2p¹) it is easier to remove the same from a partially filled 2p orbital. To remove an electron from a 2s orbital of Be atom requires more energy than to remove the same from a 2p orbital of B atom. Therefore, the ionization potential of Be is higher than that of B.

In the case of N (1s²2s²2p³), it is more difficult to remove an electron from the half-filled 2p-orbital while in the case of O (1s²2s²2p⁴) it is easier to remove the same from a partially -filled 2p-orbital. Thus the ionization potential of N is higher than that of O atom.

In a group:

As we move from top to bottom in a group of s and p-block elements, the ionization potential values of the elements go on decreasing. The decrease can be explained on the basis of the concept of atomic size, the value of principal quantum number (n) of the orbital and shielding effect. It is already proved that all the mentioned parameters increase on descending a group and their increase on moving from top to bottom in a group decreases the ionization potential values .

From the above discussion, we can explain the following :

1. The first ionization potential value of Na is greater than that of K
2. The first ionization potential of B is less than Be
3. The first ionization potential of N is greater than that of O
4. First ionization energy of Al is lowest than that of Mg