The ionization energy refers to the energy required to remove an electron from an atom or ion in its gaseous state. In the case of aluminum (Al) and magnesium (Mg), both elements belong to Group 2 (or 2A) of the periodic table, also known as the alkaline earth metals. When these metals form ions, they lose electrons to attain a stable electronic configuration.
Aluminum forms the Al3+ ion by losing three electrons, while magnesium forms the Mg2+ ion by losing two electrons. To understand why the ionization energy of Al3+ is greater than Mg2+, we need to consider the following factors:
Effective nuclear charge: The effective nuclear charge is the positive charge experienced by an electron in the outermost energy level. As we move across a period in the periodic table, the atomic number increases, resulting in a greater effective nuclear charge. In the case of aluminum, with an atomic number of 13, the effective nuclear charge experienced by its outermost electrons is higher than that of magnesium (atomic number 12). The stronger attraction from the increased positive charge makes it more difficult to remove electrons from the Al3+ ion, leading to a higher ionization energy.
Shielding effect: The shielding effect is the reduction in the attractive force between the nucleus and the valence electrons due to inner electron shells. As we move across a period, the shielding effect remains relatively constant. Since both aluminum and magnesium are in the same period (Period 3), the shielding effect is similar for both ions. Therefore, the shielding effect does not significantly affect the comparison between their ionization energies.
Considering these factors, the greater effective nuclear charge experienced by the outermost electrons in the Al3+ ion compared to the Mg2+ ion results in a stronger attraction between the electrons and the nucleus. Consequently, it requires more energy to remove an electron from the Al3+ ion, leading to a higher ionization energy for aluminum compared to magnesium.