bei48482_FM

Here n2 for the 2selectron, its ionization energy is E 2 5.39 eV, and E 1 13.6 eV is

the ionization energy of the hydrogen atom. Hence

Zn  2 1.26

The effective charge is 1.26eand not ebecause the shielding of 2eof the nuclear charge of 3e

by the two 1selectrons is not complete: as we can see in Fig. 6.11, the 2selectron has a certain

probability of being found inside the 1selectrons.

Ionization Energy

Figure 7.10 shows how the ionization energies of the elements vary with atomic number.

As we expect, the inert gases have the highest ionization energies and the alkali metals

the lowest. The larger an atom, the farther the outer electron is from the nucleus and

the weaker the force is that holds it to the atom. This is why the ionization energy

generally decreases as we go down a group in the periodic table. The increase in ion-

ization energy from left to right across any period is accounted for by the increase in

nuclear charge while the number of inner shielding electrons stays constant. In pe-

riod 2, for instance, the outer electron in a lithium atom is held by an effective charge

of about e, while each outer electron in beryllium, boron, carbon, and so on, is held

by effective charges of about  2 e,  3 e,  4 e, and so on. The ionization energy of

lithium is 5.4 eV whereas that of neon, which ends the period, is 21.6 eV.

At the other extreme from alkali metal atoms, which tend to lose their outermost

electrons, are halogen atoms, whose imperfectly shielded nuclear charges tend to

complete their outer subshells by picking up an additional electron each. Halogen

5.39 eV



13.6 eV

E 2



E 1

0102030405060708090

5

10

15

20

25

30

He

Ne

Ar

Zn

Kr

Cd

Xe

Hg Rn

Li Na KCa

Rb

In

Cs

Tl

Atomic number

Ionization energy, eV

H

Figure 7.10The variation of ionization energy with atomic number.

244 Chapter Seven

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