Define `IE_(1)` and `IE_(2)`. Why is `IE_(2) > IE_(1)` for a given atom? Discuss the factors than effect IE of an element.

Explanation

(1) Ionization energy is the amount of energy required to remove the most lossely held electron from isolated a neutral gaseous atom to convert it into gaseous ion. It is also known as first ionization energy because it is the energy required to remove the first electron from the atom. It is denoted as `I_(1)` and is expressed in electron volts atom, kilo per atom, kilo calories (or) kilo joules per mole.
`M_((g)) + I_(1) rarr M_((g)^(+)) + e^(-)`
`I_(1)` is first ionization potential.
(2) The energy required to remove another electron from the unipositive ion is called the second ionization energy. It is denoted as `I_(2)`
`M_((g))^(+) + I_(2) rarr M_((g))^(2+) + e^(-)`
(3) The second ionization potential is greater that the first ionization potential. On removing an electron from an atom, the unipositive ion formed will have more effective nuclear charge than the number of electrons. As a result the effective nuclear charge increases over the outermost electrons. Hence more energy is required to remove the second electron. This shows that the second ionization potential is greater than the first ionization potential For sodium, `I_(1)` is 5.1 eV and `I_(2)` is 47.3 eV
`I_(1) lt I_(2) lt I_(3)....I_(n)`
Factors affecting ionization potential: (1) Atomic radius: As the size of the atom increases the distance between the nucleus and the outermost electrons increases. So the effective nuclear charge on the outermost electrons decreases. In such a case the energy required to remove the electrons also decreases. This shows that with an increase in atomic radius the ionization energy decreases.
(2) Nuclear charge: As the positive charge of the nucleus increases its attraction increases over teh electrons. So it becomes more difficult to remove the electrons. This shows that the ionization energy increases as the nuclear charge increases.
(3) Screening effect or shielding effect: In multielectron atoms, valence electrons are attracted by the nucleus as well as repelled by electrons of inner shells. The electrons present in the inner shells screen the electrons present in the outermost orbit from the nucleus. As the number of electrons in the inner orbits increases, the screening effect increases. This reduces the effective nuclear charge over the outermost electrons. It is called screening or shielding effect. With the increase of screening effect the ionization potential decreases. Screening efficiency of the orbitals falls off in the order `s gt p gt d gt f`.
(Magnitude of screening effect) `prop (I)/(("Ionization enthalpy"))`
Trend In a Group: The ionisation potential decreases in a group, gradually from top to bottom as the size of the elements increases down a group.
Trend in a Period: In a period from left to right I.P. value increases as the size of the elements decreases along the period.