Ionization Energy

Ionization energy is the energy needed to remove an electron from an atom. We can formally state the definition of ionization energy in the following equations, which represent the 1st, 2nd, and 3rd ionization energies of aluminum.

Why is the 4th ionization energy of aluminum so much larger than the others? Consider the electron configuration of aluminum, [Ne] 3s23p1, after aluminum has lost three electrons to form Al3+ it has an electron configuration identical to neon. Since the noble gas electron configurations are particularly stable it takes a lot of energy to change the electron configuration once the aluminum atom has made it to the promised land.

If you look at table 7.2 in your text, you find that aluminum is not unique in this regard. We see right away that it takes a lot of energy to remove electrons beyond the nearest noble gas configuration. This is numerical confirmation of what I’ve been telling you Atoms like to attain a noble gas configuration.

 

Periodic Trends in 1st IE

1. The first ionization energy increases as you move left to right across a row (period) of the periodic table (due to increasing effective nuclear charge).

2. The first ionization energy decreases as you move down a column (group) of the periodic table.

3. The s & p block elements show a larger range of 1st IE than the d and f block elements.

These trends are exactly opposite of those observed for atomic radius. This correlation is no coincidence. The outermost electrons in small atoms such as He are located close to the nucleus, which results in a strong Coulombic attraction. Thus it takes a lot of energy to remove an electron from a small atom, so these atoms have high first ionization energies. Conversely, the outermost electrons in large atoms are located farther away from the nucleus, this leads to a decreased attraction with the nucleus and a smaller first ionization energy.

Irregularities

If you look at the 1st ionization energies of the elements in figure 7.7 of your book, you will notice some exceptions to the periodic trends stated above. For example nitrogen has a greater 1st IE than oxygen, and berylium has a greater 1st IE than boron. We can understand these anomalies by considering the electron configurations of these elements.

 

. . . . . . . . .

(2p)

(2s)

(1s)


. . . . . . .

(2p)

(2s)

(1s)


. . .

(2p)

(2s)

(1s)


. .

(2p)

(2s)

(1s)


You can see that berylium has only filled subshells (1s and 2s). That will no longer be true if we remove an electron, whereas, boron will attain an electron configuration with only filled subshells (1s and 2s) if we remove one electron. This is enough to reverse the usual trends in ionization energy.

We can sing the same song with nitrogen and oxygen, except that now we are talking about a half-filled subshell (2p) rather than a filled subshell.

This leads to a general rule regarding the stability of electron configurations:

Completely filled and half-filled subshells are more stable than might otherwise be expected.

You should keep in mind that although electron configurations where the s subshell is completely filled (group 2A), or where the p subshell is half filled (group 5A), are generally more stable than their neighbors, these configurations are not nearly as stable as ones where the p-orbitals are completely filled (group 8A, noble gases). You can clearly see this if you compare the difference in 1st ionization energies of Ne (2081 kJ/mol) and Na (496 kJ/mol), with the above examples of Be, B, N, and O.