Ionization energy is the minimum amount of energy required to remove the most loosely held electron from an isolated atom in its gaseous state. Scientists measure successive ionization energies (\(IE_1, IE_2, IE_3\), and so on) for the same atom. A consistent observation across all elements is that each subsequent ionization energy is always higher than the one before it, meaning the energy needed to remove the second electron (\(IE_2\)) is greater than the energy for the first electron (\(IE_1\)).
Defining Ionization Energy and the First Electron Removal
The first ionization energy (\(IE_1\)) specifically measures the energy to initiate the conversion of a neutral gaseous atom, represented by \(X\), into a univalent positive ion, \(X^+\). This process is chemically represented by the equation \(X(g) + \text{energy} \longrightarrow X^+(g) + e^-\). The initial electron being removed is the valence electron, the one farthest from the nucleus and therefore experiencing the weakest attractive force.
A larger atomic radius means the valence electron is further from the positively charged nucleus, resulting in a lower \(IE_1\). A higher number of internal electron shells also creates a stronger shielding effect, which reduces the net pull on the outer electron, making it easier to remove.
The electron being removed is housed in the highest energy level, making it the most accessible target for ionization. For example, sodium (Na) requires \(496 \text{ kJ/mol}\) to lose its first electron, a relatively low value that reflects the ease of removing a single electron in the outermost shell.
The Primary Cause: Increased Effective Nuclear Charge
The requirement for more energy to remove the second electron stems directly from the change in the atom’s charge and structure after the first ionization. Once the neutral atom \(X\) becomes a positively charged ion \(X^+\), the balance between positive protons and negative electrons is disrupted. The resulting ion still possesses the same number of protons in its nucleus, but it now has one fewer electron.
This change means the constant positive charge of the nucleus is now distributed across a smaller number of remaining electrons. Consequently, the attractive force felt by each of the remaining electrons increases substantially, a concept quantified by the effective nuclear charge (\(Z_{eff}\)). The reduction in electron-electron repulsion among the remaining electrons also contributes to a greater net force pulling them toward the nucleus.
This stronger net attraction causes the electron cloud of the \(X^+\) ion to contract slightly compared to the neutral atom. The second electron must be pulled away from an ion that is both smaller and more positively charged overall. For instance, the first ionization energy of magnesium (Mg) is \(737 \text{ kJ/mol}\), but its second ionization energy (\(IE_2\)) is \(1450 \text{ kJ/mol}\).
The second electron for magnesium is still a valence electron, but its removal is from the already charged \(\text{Mg}^+\) ion. The increased effective nuclear charge (\(Z_{eff}\)) on the remaining electrons makes it substantially more difficult to detach. This consistent, moderate rise in energy explains why successive ionization energies always increase.
The Extreme Case: Ionization Energy Jumps After Core Shell Depletion
While the moderate increase between successive ionization energies is explained by the rising effective nuclear charge, a far more dramatic jump occurs when the electron being removed must be taken from a filled inner electron shell. These massive energy discrepancies occur when the ionization process breaks into the atom’s stable, core electron configuration. The core electrons are located closer to the nucleus and are poorly shielded by one another, experiencing an extremely high \(Z_{eff}\).
Sodium (Na), a Group 1 element, provides the clearest example of this phenomenon. Its first ionization energy (\(IE_1\)) is relatively low at \(496 \text{ kJ/mol}\), as it involves removing the single valence electron in the \(3s\) orbital. This removal leaves the \(\text{Na}^+\) ion with a stable electron configuration identical to the noble gas neon.
The second ionization energy (\(IE_2\)) for sodium, however, is a massive \(4563 \text{ kJ/mol}\). This is nearly a tenfold increase from \(IE_1\). This colossal energy requirement is because the second electron must be extracted from the highly stable, full \(2p\) inner shell, which is much closer to the nucleus and has a complete octet configuration.
In comparison, the difference between \(IE_1\) and \(IE_2\) for magnesium was moderate because both electrons were from the valence shell. The extreme jump for magnesium occurs at its third ionization energy (\(IE_3\)), rising from \(1450 \text{ kJ/mol}\) to \(7731 \text{ kJ/mol}\). This is because the third electron must be pulled from its noble gas-like core. These substantial energy jumps serve as direct evidence of the stability of inner electron shells compared to the more loosely bound valence electrons.