In general, ionization energy increases from left to right across a given period. Aluminum, however, has a lower ionization energy than magnesium. Explain.

The periodic table is a systematic arrangement of elements, ordered by their atomic numbers, electron configurations, and recurring chemical properties. Elements are placed in rows by increasing atomic number in a new row, also known as a period, when the outer electron shell is filled. Within each period, the ionization energy— the energy required to remove an electron from a gaseous atom or ion—typically increases as one moves from left to right. This trend is due to the increasing positive charge of the nucleus, which more strongly attracts the negatively charged electrons.

However, there are exceptions to this rule, such as the case with aluminum and magnesium. While magnesium has a higher ionization energy compared to aluminum, despite being to its left, it can be better understood by considering not just the position on the periodic table, but also the underlying electron configuration and the resulting atomic structure, which affect the force holding electrons in the atom.

Electrons in an atom are found in regions called orbitals, and the pattern in which these orbitals are filled is known as the electron configuration. It significantly influences an atom's chemical behavior and properties, including ionization energy. For instance, magnesium's electron configuration is \[\text{[Ne]}3s^2\], meaning it has two electrons in the 3s subshell following the noble gas neon's configuration. On the other hand, aluminum, with an electron configuration of \[\text{[Ne]}3s^23p1\], has an additional electron in the 3p subshell. Electrons in the p subshell are generally farther from the nucleus and more shielded by electrons in the s subshell, which makes them easier to remove—hence, requiring less ionization energy.

In Al's case, the single 3p electron experiences a higher level of shielding by the filled 3s subshell as well as being naturally slightly further from the nucleus compared to Mg's 3s electrons, which explains Al's lower ionization energy.

The atomic structure denotes the distribution of electrons around the nucleus and the forces between them and the nucleus. It shapes the atom's characteristics, such as ionization energy. The closer an electron is to the nucleus, the stronger the electrostatic forces that bind it, implying a higher ionization energy required to remove the electron.

For magnesium, the 3s electrons are closer to the nucleus and experience a stronger pull from it compared to aluminum's 3p electron. Although both elements are in the same period, the different types of orbitals in which the outermost electrons reside affect their atomic structures. The 3s subshell is more penetrating, meaning its electrons can get closer to the nucleus, compared to the 3p subshell, whose electrons are generally held less tightly. Thus, the differences in the atomic structures of Mg and Al explain why, despite periodic trends, there is a decrease in ionization energy from magnesium to aluminum.