Animated Solution for Chemistry - s and p-Block Elements: Beryllium and aluminium exhibit many properties which are similar. But the two elements differ in

Have you ever looked at the periodic table and wondered why some elements, despite being in completely different groups, act like long-lost twins? This phenomenon is one of the most beautiful quirks of inorganic chemistry, known as the diagonal relationship.

Imagine you are walking across the periodic table. As you move from left to right across a period, the atomic size decreases and the electronegativity increases. But, as you move down a group, the atomic size increases and the electronegativity decreases.

Now, what happens if you take one step to the right and one step down? The two opposing trends effectively cancel each other out! This creates a diagonal pair of elements that share a remarkably similar ionic potential (the ratio of charge to radius).

Beryllium (Group 2, Period 2) and Aluminium (Group 13, Period 3) are the poster children for this diagonal relationship. Because their ionic potentials are nearly identical, their polarizing power is the same, leading to a striking similarity in their chemical behavior.

The Similarities Let's break down the options provided in the question. The problem states that Beryllium and Aluminium share many properties, and we need to find the exception.

First, consider their hydrides. Both Beryllium and Aluminium form polymeric hydrides, specifically $({\text{BeH}}_2{)}_n$ and $({\text{AlH}}_3{)}_n$. These are fascinating electron-deficient structures where hydrogen atoms act as bridges between the metal centers, forming unique 3-center-2-electron bonds (banana bonds).

Next, let's look at their halides. According to Fajans' Rules, a small cation with a high charge will highly polarize a larger anion, pulling the electron cloud towards itself and introducing significant covalent character into an otherwise ionic bond. Because both ${\text{Be}}^{2+}$ and ${\text{Al}}^{3+}$ have high polarizing power, their halides, such as ${\text{BeCl}}_2$ and ${\text{AlCl}}_3$, are predominantly covalent.

Finally, consider their oxides. While most metal oxides are basic, the oxides of Beryllium ($\text{BeO}$) and Aluminium (${\text{Al}}_2{\text{O}}_3$) are amphoteric. This means they have the unique ability to react with both strong acids and strong bases, acting as chemical chameleons depending on their environment.

So, options (b), (c), and (d) represent their shared similarities. The difference must lie in option (a).

Maximum Covalency To understand why Beryllium and Aluminium differ in their maximum covalency, we must dive into the quantum mechanical architecture of their atoms.

Covalency is simply the number of electron pairs an atom can share with other atoms. The maximum covalency is strictly dictated by the number of available orbitals in the atom's outermost shell, known as the valence shell.

Think of the valence shell as a parking lot for electrons. You can only park as many cars (electron pairs) as there are parking spaces (orbitals).

The Rule of Four Let's examine Beryllium. Its atomic number is $Z=4$, and its electronic configuration is $1s^{2}2s^2$.

The outermost shell for Beryllium is the second principal quantum shell ($n=2$). Quantum mechanics tells us that the $n=2$ shell only contains two types of subshells: the $2s$ subshell (which has 1 orbital) and the $2p$ subshell (which has 3 orbitals: $2p_x$, $2p_y$, and $2p_z$).

If we count the total number of "parking spaces" in the $n=2$ shell, we get: $1\text{ (from 2s)}+3\text{ (from 2p)}=4\text{ orbitals}$

Because Beryllium only has 4 orbitals in its valence shell, it can accommodate a maximum of 4 electron pairs. Therefore, its maximum covalency is strictly limited to 4. It simply does not have the architectural space to form any more bonds. A classic example of this is the tetrafluoroberyllate ion, $[{\text{BeF}}_4{]}^{2-}$.

The Power of d-Orbitals Now, let's shift our focus to Aluminium. Its atomic number is $Z=13$, and its electronic configuration is $1s^{2}2s^{2}2p^{6}3s^{2}3p^1$.

The outermost shell for Aluminium is the third principal quantum shell ($n=3$). This is where the magic happens! The $n=3$ shell contains the $3s$ subshell (1 orbital), the $3p$ subshell (3 orbitals), and crucially, the $3d$ subshell (5 orbitals).

Even though the $3d$ orbitals are empty in a ground-state Aluminium atom, they are still physically present and energetically accessible. They act as a massive expansion to our electron parking lot.

When Aluminium reacts with highly electronegative ligands like Fluorine, it can utilize these vacant $3d$ orbitals to accept additional electron pairs. This phenomenon is known as octet expansion.

Because of the availability of these d-orbitals, Aluminium is not restricted to 4 bonds. It can easily expand its covalency to 6, utilizing one $3s$, three $3p$, and two $3d$ orbitals to undergo $s{p}^3{d}^2$ hybridization. A perfect example of this is the hexafluoroaluminate ion, $[{\text{AlF}}_6{]}^{3-}$.

While the diagonal relationship makes Beryllium and Aluminium behave like twins in many macroscopic chemical reactions, their underlying quantum architecture sets a hard limit on their bonding capabilities.

Beryllium, trapped in the second period, is forever constrained by the lack of d-orbitals, capping its covalency at 4. Aluminium, enjoying the spaciousness of the third period, leverages its vacant d-orbitals to push its covalency up to 6.

Therefore, the property in which they differ is their maximum covalency in compounds, making option (a) the correct answer. Always remember this golden rule of inorganic chemistry: second-period elements can never expand their octet!