Rank the following elements in order of increasing first ionization energy: Na, Mg, Al, Si.

• A. Mg < Na < Si < Al
• B. Na < Mg < Al < Si
• C. Al < Si < Na < Mg
• D. Si < Al < Mg < Na

Answer: (B)

Verified step by step guidance

1. Recall that the first ionization energy is the energy required to remove one mole of electrons from one mole of gaseous atoms, and it generally increases across a period from left to right due to increasing nuclear charge and decreasing atomic radius.
2. Identify the elements given: Na (Sodium), Mg (Magnesium), Al (Aluminum), and Si (Silicon), which are all in the same period (Period 3) of the periodic table.
3. Understand the general trend: As you move from left to right across Period 3, the first ionization energy tends to increase because the effective nuclear charge increases, pulling electrons closer and making them harder to remove.
4. Consider any exceptions or subtleties: For example, Aluminum has a slightly lower ionization energy than Magnesium because the electron removed from Aluminum is in a p orbital, which is higher in energy and less tightly held than the s orbital electron removed from Magnesium.
5. Based on these principles, arrange the elements in order of increasing first ionization energy by comparing their positions and electron configurations: start with the element with the lowest ionization energy and proceed to the highest.