Animated Solution for Chemistry - s and p-Block Elements: The electronegativity of aluminium is similar to

When we look at the periodic table, we observe distinct trends in properties like electronegativity. As we move from left to right across a period, the effective nuclear charge increases while the shielding effect remains relatively constant, causing the electronegativity to increase. Conversely, as we move down a group, the addition of new electron shells significantly increases the atomic radius, which outweighs the increase in nuclear charge, leading to a decrease in electronegativity.

Because of these two opposing trends—increasing across a period and decreasing down a group—an interesting phenomenon emerges. Elements that are situated diagonally to each other (moving one step right and one step down) often exhibit remarkably similar properties. This is famously known as the diagonal relationship.

For instance, if we start at Beryllium ($\text{Be}$) in Group 2, Period 2, and move diagonally to the right and down, we land on Aluminum ($\text{Al}$) in Group 13, Period 3.

Ionic Potential

The root cause of this similarity lies in a property called ionic potential ($\varphi$). Ionic potential is defined as the ratio of the effective nuclear charge ($Z^{\ast }$) to the atomic radius ($r$):

For Beryllium and Aluminum, the increase in $Z^{\ast }$ from Group 2 to Group 13 is almost perfectly counterbalanced by the increase in atomic radius from Period 2 to Period 3. As a result, their ionic potentials are nearly identical:

Because their ionic potentials are so similar, their polarizing power and, consequently, their electronegativity values are almost the same. On the Pauling scale, the electronegativity of Beryllium is exactly $1.5$, and the electronegativity of Aluminum is also $1.5$.

Therefore, the electronegativity of Aluminum is most similar to that of Beryllium.