Ionization energy (IE) generally increases from left to right across the periodic table. This fundamental periodic property shows a clear pattern of increasing as one moves across a row, or period. This trend means that elements on the far right hold onto their electrons much more tightly than those on the far left. Understanding this pattern requires looking at how the nucleus interacts with its surrounding electrons.
What Ionization Energy Measures
Ionization energy is the energy required to remove the most loosely bound electron from an isolated atom in its gaseous state. This measurement indicates how strongly an atom resists losing an electron to form a positive ion. A higher ionization energy value signifies that more energy is needed to overcome the attractive force of the nucleus, meaning the electron is held more securely.
The first ionization energy (\(\text{IE}_1\)) refers to the energy necessary to remove the first electron from a neutral gaseous atom. Although atoms have successive ionization energies (\(\text{IE}_2\), \(\text{IE}_3\), and so on), \(\text{IE}_1\) is the standard measure used to define the general trend. Values are typically measured in kilojoules per mole (\(\text{kJ/mol}\)).
The General Trend Across a Period
The consistent observation across every row of the periodic table is a rise in first ionization energy as the atomic number increases. Alkali metals, such as Lithium and Sodium (Group 1), exhibit the lowest ionization energies within their periods. Their electrons are relatively easy to remove, reflecting their tendency to form positive ions.
In contrast, the noble gases (Group 18) possess the highest ionization energies for their periods. Their exceptionally stable electron configurations require significantly greater energy to remove an electron. This overall upward slope in energy from left to right is a defining characteristic of the periodic table.
The Role of Effective Nuclear Charge
The general increase in ionization energy across a period is primarily explained by the effective nuclear charge (\(Z_{eff}\)). \(Z_{eff}\) is the net positive charge a valence electron experiences from the nucleus, accounting for the shielding effect of inner electrons. As one moves across a period, the number of protons in the nucleus steadily increases, adding one unit of positive charge with each step.
The added electrons are placed into the same principal energy level or shell. Since the number of inner, core electrons providing shielding remains constant, the increasing number of protons causes the \(Z_{eff}\) felt by the outermost electrons to increase significantly.
A higher effective nuclear charge results in a stronger electrostatic attraction between the nucleus and the valence electrons. This stronger pull draws the electron cloud closer, decreasing the atomic radius. Consequently, more energy is required to overcome this intensified attraction and remove an electron, leading to the observed increase in ionization energy.
Key Deviations from the Smooth Trend
While the overall pattern is an increase, the rise is not perfectly smooth, and two predictable deviations occur due to specific electron configurations.
Deviation 1: Group 2 to Group 13
The first deviation is the drop in ionization energy from Group 2 (Alkaline Earth Metals) to Group 13 (Boron Group). Group 2 elements, such as Beryllium (\(\text{Be}\)), have a stable, full \(\text{s}\)-orbital configuration (\(\text{ns}^2\)).
When moving to Group 13 elements, like Boron (\(\text{B}\)), the next electron enters a new, higher-energy \(\text{p}\)-orbital (\(\text{ns}^2\text{np}^1\)). This \(\text{p}\)-orbital electron is slightly further from the nucleus and is shielded by the filled \(\text{ns}^2\) subshell. Because it is in a higher-energy, more shielded orbital, it is easier to remove than an electron from the stable \(\text{ns}^2\) subshell, causing the ionization energy to dip.
Deviation 2: Group 15 to Group 16
The second deviation is the drop from Group 15 (Nitrogen Group) to Group 16 (Oxygen Group). Group 15 elements, such as Nitrogen (\(\text{N}\)), have a stable, half-filled \(\text{p}\)-orbital configuration (\(\text{np}^3\)) with three unpaired electrons.
When the next electron is added to form a Group 16 element, such as Oxygen (\(\text{O}\)), it must pair up with an existing electron in a \(\text{p}\)-orbital (\(\text{np}^4\)). This pairing introduces electron-electron repulsion within that specific orbital. This repulsion slightly destabilizes the configuration and makes the paired electron easier to remove, resulting in a consistent decrease in the first ionization energy.