How would you explain the fact that the first ionization enthalpy of sodium is lower than that of magnesium but its second ionization enthalpy is higher than that of magnesium?
The first ionization enthalpy of sodium is lower than that of magnesium because magnesium has a greater nuclear charge than sodium. Both elements have their outermost electrons in the 3s orbital, but the higher nuclear charge of magnesium holds its electrons more tightly, making them harder to remove compared to the single 3s electron in sodium.
However, the second ionization enthalpy of sodium is higher than that of magnesium. After losing one electron, sodium acquires a stable noble gas configuration (). Removing a second electron from this positively charged sodium ion means pulling an electron from a completely filled inner shell, which is extremely difficult. In contrast, magnesium loses its second electron from the 3s orbital to achieve a stable noble gas configuration, which requires comparatively less energy.
Explanation
The first ionization enthalpy depends on the nuclear charge and the shielding effect. Magnesium has a higher nuclear charge than sodium, resulting in a higher first ionization enthalpy. For the second ionization enthalpy, the electronic configurations of the resulting +1 ions play a crucial role. Na+ has a stable noble gas core, and removing an electron from a positively charged ion with a filled shell requires significantly more energy than removing the remaining 3s electron from Mg+, which also aims to achieve a stable noble gas configuration.