Chemistry Reference
In-Depth Information
Tab l e 7 . 2
The atomic configurations and ionization potentials (eV) for
the first two rows of the periodic table. For boron and heavier elements
it is possible to ionize from either 2s or 2p states, and so this data gives
an estimate of the energy differences between these orbitals.
Atom
Z
Atomic config.
1s
2s
2p
2s
−
2p
H
1
1s
1
13.6
—
—
—
He
2
1s
2
24.6
—
—
—
[He]2s
1
2p
0
Li
3
—5.4
— —
[He]2s
2
2p
0
Be
4
—9.3
— —
B
5
[He]2s
2
2p
1
—
14.0
8.3
5.7
C
6
[He]2s
2
2p
2
—
19.4
10.6
8.8
N
7
[He]2s
2
2p
3
—
25.6
13.2
12.4
O
8
[He]2s
2
2p
4
—
32.3
15.8
16.5
F
9
[He]2s
2
2p
5
—
40.2
18.6
21.6
Ne
10
[He]2s
2
2p
6
48.5
21.6
26.9
The ionization energies for the first two rows of the periodic table are given in Table 7.2.
Figure 7.11 shows that for second-row elements the two electrons that occupy the 1s
energy level remain close to the nucleus, and so the 1s energy level is much lower than
the 2s or 2p states. Electrons in the 2s and 2p states of first-row elements spend most of
their time outside the inner shell volume; so, on average, they experience a nuclear charge
which is shielded by the core electrons. The shielding has an effect on the atomic energy
levels of the outer or valence electrons, so that, for example, the ionization energy of Li
(1s
2
2s
1
) at 5.3917 eV is actually lower than that of H, which is 13.5984 eV, despite the
3
nuclear charge of Li. The core electrons of Li shield the valence orbitals so that the
effective nuclear charge, as far as the 2s and 2p orbitals are concerned, is around the same
as the 1
+
of the H nucleus. The H 1s
1
electron is in a state closer to the positive charge, so
is lower in energy and more difficult to ionize than the Li 2s
1
electron.
The small peak in the 2s probability function shown in Figure 7.11b exposes an electron
in this state to the unshielded nuclear charge slightly more often than a 2p electron, and
so the 2s state is always the lower in energy. Hence, for all elements with both 2s and
2p electrons in Table 7.2 the ionization energy for 2s is greater than for 2p. As we move
from left to right across the periodic table the nuclear charge increases and this difference
becomes larger. Figure 7.13 extends this comparison to the third and fourth rows of the
periodic table. The same trend, of the gap between
n
s and
n
p orbitals becoming wider
across the row, can be seen lower down the table. However, the range of values observed
for rows 3 and 4 is less than for the second row, as the higher principal quantum number
for these levels places the electrons, on average, further from the core of the atoms.
It is also notable from the values in Table 7.2 that the ionization potential of He is not
quite twice that of H. This is due to the electron-electron repulsion between the two 1s
electrons confined to the same spatial orbital in He. In the equations discussed so far,
the electron-electron repulsion has not been explicitly included. In the H
2
example we
simply assume that this effect is outweighed by the attractive interaction between the elec-
trons and nuclei, so that it is favourable to put both electrons in the bonding MO despite
the resulting electron-electron repulsion. The MOs obtained give the correct qualitative
picture, but a more accurate account of the electron-electron interaction is required to
+
