Question

You learned in this chapter that ionization generally increases as you move from left to right across the periodic table. However, consider the following data, which shows the ionization energies of the period 2 and 3 elements: Notice that the increase is not uniform. In fact, ionization energy actually decreases a bit in going from elements in group \(2 \mathrm{~A}\) to \(3 \mathrm{~A}\) and then again from \(5 \mathrm{~A}\) to \(6 \mathrm{~A}\). Use what you know about electron configurations to explain why these dips in ionization energy exist.

Short answer

Ionization energy decreases from Group 2A to 3A due to the addition of an electron to a new, higher energy p orbital, and from 5A to 6A because adding an electron to a half-filled p subshell increases repulsions, thus destabilizing it and making it easier to ionize.

Step-by-step solution

Identify Group 2A to 3A Transition

The transition from Group 2A (alkaline earth metals) to Group 3A (group 13 elements) involves moving from an element with a filled s subshell to one where the first p subshell electron is added. Elements in Group 2A have their outermost s subshell completely filled, which is a particularly stable arrangement. Adding one more electron (moving to Group 3A) puts the electron in a p orbital, which is higher in energy, and not as effectively shielded by the inner electrons from the nucleus. This makes it easier to remove the first electron in a Group 3A element compared to a Group 2A element.

Identify Group 5A to 6A Transition

Moving from Group 5A to Group 6A, you are adding an electron to a p subshell that is already half-filled. For Group 5A elements, this means there are three unpaired electrons in the p subshell, which is a stable configuration due to minimized electron repulsion. Adding an additional electron (to form a Group 6A element) pairs with one of the unpaired electrons, increasing electron-electron repulsions and slightly destabilizing the arrangement. Therefore, it requires slightly less energy to remove an electron from a Group 6A element than from a Group 5A element.

Consider Electron Configuration Stability

The stability of an element's electron configuration plays a vital role in its ionization energy. Elements with a completely filled or half-filled subshell are generally more stable, making their ionization energies higher. Conversely, adding an electron to a filled or half-filled subshell makes the electron easier to remove, corresponding to the dips in ionization energy observed when moving from Group 2A to 3A and from Group 5A to Group 6A.

Key concepts

Electron Configurations

Understanding electron configurations is essential for grasping why certain ionization energy trends occur in elements. Electron configurations describe the arrangement of electrons around the nucleus of an atom in definite energy levels and sublevels. The most stable configurations are those where sublevels are either completely full or half-full.

For instance, the noble gases exhibit very high ionization energies as their outermost energy levels are fully filled, creating a stable electron configuration. Similarly, elements with half-filled sublevels can be more stable than those with one more or one less electron. This characteristic stability influences the ionization energy, making it higher for fully or half-filled configurations and lower when such harmony is disrupted by adding or removing electrons.

When an extra electron is introduced to a stable structure, it increases electron–electron repulsion, which destabilizes the atom and makes it easier to remove this additional electron. Understanding this concept helps explain certain quirks in ionization energy trends across the periodic table.

Periodic Table Groups

The periodic table is organized into vertical columns known as 'groups' or 'families,' which contain elements with similar properties due to their uniform valence electron configurations. As we move down a group, the number of energy levels increases, leading to electrons being farther from the nucleus. This larger distance and the additional inner electron shells reduce the effective nuclear charge experienced by the outermost electrons, making them easier to remove.

Contrastingly, across a period (horizontally), elements gain additional protons in the nucleus and electrons in the same energy level, increasing the effective nuclear charge. Typically, this leads to a gradual increase in ionization energy from left to right across a period. However, this trend is occasionally interrupted, particularly when the new electron occupies a sublevel that is less stable than before, such as moving from a filled s sublevel to the first p orbital electron.

Alkaline Earth Metals

Alkaline earth metals make up Group 2A of the periodic table, and include elements such as beryllium, magnesium, and calcium. These metals are characterized by their electron configuration ending in an s₂ sublevel. Having two electrons in their outermost s orbital means this sublevel is filled, which is a particularly stable arrangement.

The energy required to remove one of these electrons (ionization energy) is relatively high because removing an electron disrupts this stability. However, when moving from an alkaline earth metal in Group 2A to the next element in Group 3A, the new electron goes into the p orbital. This addition results in a less stable configuration, and thus, the ionization energy drops. When learning about periodic trends, students must note the distinctive stability of fully-filled orbitals like those found in alkaline earth metals, which directly impacts their ionization energies.

P Orbital Electron Removal

When discussing ionization energy and electron removal, the concept of p orbitals plays an important role. P orbitals begin to fill after the s sublevel is filled, and can hold a total of six electrons, forming three pairs. The removal of a p orbital electron is typically easier than the removal of an s orbital electron because p electrons are farther away from the nucleus and thus experience a lesser effective nuclear charge.

Ionization energy trends show a dip between the removal of the last electron in an s sublevel and the first electron in a p sublevel. This occurs because p orbitals are higher in energy and less shielded from nuclear charge by inner electrons. Moving from Group 5A to Group 6A involves adding an electron to a p orbital that is already half-filled, which means overcoming electron repulsion. Therefore, removing such an electron from a Group 6A element requires slightly less energy than from a Group 5A element. This nuanced understanding of p orbital electron removal is pivotal for students tackling questions about irregularities in ionization energy across the periodic table.