As one goes across a period on the
periodic table, the slight breaks in the
increasing ionization energies occur
between Groups IIA and IIIA
(electrons first enter the npsubshell)
and again between Groups VA and
VIA (electrons are first paired in the
npsubshell).
244 CHAPTER 6: Chemical Periodicity
Elements with low ionization energies (IE) lose electrons easily to form cations.We see that in each period of Figure 6-2, the noble gases have the highest first ioniza-
tion energies. This should not be surprising, because the noble gases are known to be very
unreactive elements. It requires more energy to remove an electron from a helium atom
(slightly less than 4.0 10 ^18 J/atom, or 2372 kJ/mol) than to remove one from a neutral
atom of any other element.He(g)2372 kJ88nHe(g)eThe Group IA metals (Li, Na, K, Rb, Cs) have very low first ionization energies. Each
of these elements has only one electron in its outermost shell (... ns^1 ), and they are the
largest atoms in their periods. The first electron added to a shell is easily removed to form
a noble gas configuration. As we move down the group, the first ionization energies become
smaller. The force of attraction of the positively charged nucleus for electrons decreases
as the square of the distance between them increases. So as atomic radii increase in a given
group, first ionization energies decrease because the outermost electrons are farther from
the nucleus.
Effective nuclear charge, Zeff, increases going from left to right across a period. The
increase in effective nuclear charge causes the outermost electrons to be held more tightly,
making them harder to remove. The first ionization energies therefore generally increase
from left to right across the periodic table. The reason for the trend in first ionization
energies is the same as that used in Section 6-2 to explain trends in atomic radii. The first
ionization energies of the Group IIA elements (Be, Mg, Ca, Sr, Ba) are significantly higher
than those of the Group IA elements in the same periods. This is because the Group IIA
elements have higher Zeffvalues and smaller atomic radii. Thus, their outermost electrons
are held more tightly than those of the neighboring IA metals. It is harder to remove an
electron from a pair in the filled outermost sorbitals of the Group IIA elements than to
remove the single electron from the half-filled outermost sorbitals of the Group IA
elements.
The first ionization energies for the Group IIIA elements (B, Al, Ga, In, Tl) are excep-
tions to the general horizontal trends. They are lowerthan those of the IIA elements in
the same periods because the IIIA elements have only a single electron in their outermost
porbitals. Less energy is required to remove the first pelectron than the second selec-
tron from the outermost shell, because the porbital is at a higher energy (less stable) than
an sorbital within the same shell (nvalue).
Going from Groups IIIA to VA, electrons are going singly into separate nporbitals,
where they do not shield one another significantly. The general left-to-right increase in
IE 1 for each period is interrupted by a dip between Groups VA (N, P, As, Sb, Bi) and VIA
elements (O, S, Se, Te, Po). Presumably, this behavior is because the fourth npelectron
in the Group VIA elements is paired with another electron in the same orbital, so it expe-
riences greater repulsion than it would in an orbital by itself. This increased repulsion
apparently outweighs the increase in Zeff, so the fourth npelectron in an outer shell (Group
VIA elements) is somewhat easier to remove (lower ionization energy) than is the third
npelectron in an outer shell (Group VA elements). After the dip between Groups VA and
VIA, the importance of the increasing Zeffoutweighs the repulsion of electrons needing
to be paired, and the general left-to-right increases in first ionization energies resume.
Knowledge of the relative values of ionization energies assists us in predicting whether
an element is likely to form ionic or molecular (covalent) compounds. Elements with low
ionization energies form ionic compounds by losing electrons to form cations(positivelyBy Coulomb’s Law, F
(q
d)(
2q)
,the attraction for the outer shell
electrons is directly proportional to
the effectivecharges and inversely
proportional to the squareof the
distance between the charges. Even
though the effective nuclear charge
increases going down a group, the
greatly increased size results in a
weaker net attraction for the outer
electrons and thus in a lower first
ionization energy.
General trends in first ionization
energies of A group elements with
position in the periodic table.
Exceptions occur at Groups IIIA
and VIA.
First IEIncreaseDecrease