Ionization energies differ significantly, depending on the shell from which the electron is taken. For instance, it takes less energy to remove a p electron than an s electron, even less energy to extract a d electron, and the least energy to extract an f electron. As you can probably guess, this is because s electrons are held closer to the nucleus, while f electrons are far from the nucleus and less tightly held. You’ll need to remember two important facts about ionization energy for the test. The first is that ionization energy increases as we move across a period.
The reason for this, as is the case with periodic trends in atomic radii, is that as the nucleus becomes more positive, the effective nuclear charge increases its pull on the electrons and it becomes more difficult to remove an electron.
The second thing you’ll need to remember is that ionization energy decreases as you move down a group or family. The increased distance between electrons and the nucleus and increased shielding by a full principal energy level means that it requires less energy to remove an electron. Shielding occurs when the inner electrons in an atom shield the outer electrons from the full charge of the nucleus. Keep in mind that this phenomenon is only important as you move down the periodic table! Here are the values for the first ionization energies for some elements:
There are some important exceptions to the above two ionization energy trends in the periodic table, so make sure you study these closely:
- When electron pairing first occurs within an orbital, electron-electron repulsions increase, so that removing an electron takes less energy (it’s easier); thus the IE drops at this time. For example, less energy is required to remove an electron from oxygen’s valence in spite of an increasing Zeff because oxygen’s p4 electron is the first to pair within the orbital. The repulsion created lowers the amount of energy required to remove either electron.
- There is also a drop in ionization energy from s2 to p1—also in spite of an increasing Zeff. This drop is due to the fact that you are removing a p electron rather than an s electron. The p electrons are less tightly held because they do not penetrate the electron cloud toward the nucleus as well as an s electron does.
Example
Which of the following elements has the highest ionization energy: K, Ca, Ga, As, or Se?
Explanation
The answer is arsenic, or As. Since IE increases as we move across a period, you may have chosen Se. However, there is a drop in IE in spite of increasing Zeff due to the increased electron-electron repulsion in the family that contains oxygen, since they are np4.
Electron Affinity
An atom’s electron affinity is the amount of energy released when an electron is added to the atom in its gaseous state—when an electron is added to an atom, the atom forms a negative ion. Most often, energy is released as an electron is added to an atom, and the greater the attraction between the atom and the electron added, the more negative the atom’s electron affinity.
