Question

How would you explain the lower atomic radius of Ga as compared to Al?

Short answer

Ga's 3d electrons have poor shielding, enhancing nuclear pull, making its atomic radius smaller than Al's.

Step-by-step solution

Understand the Basic Concepts

First, let's recall what atomic radius is: it's a measure of the size of an atom from its nucleus to its outer boundary of electron cloud. As we move down a group in the periodic table, the atomic radius generally increases because new energy levels are being added.

Recall Periodic Trends

While Al (aluminum) and Ga (gallium) are in the same group, moving from period 3 (Al) to period 4 (Ga), one would initially expect the atomic radius to increase due to the addition of an electron shell.

Consider the Role of d-Orbitals

In Ga, after the 3p subshell, the 3d subshell starts to fill with electrons. The ineffective shielding effect of these inner d electrons causes Ga to have a smaller atomic radius than expected, as these electrons do not shield the outer electrons effectively from the attraction of the nucleus.

Apply Electron Shielding Concept

Even though Ga has more protons than Al, the ineffective shielding from the electron-filled 3d orbitals in Ga leads to a stronger effective nuclear charge than in Al. This pulls the outer electrons closer to the nucleus in Ga, resulting in a smaller atomic radius than Al.

Key concepts

Periodic Trends

When we talk about periodic trends, we explore the predictable changes in properties of elements as we move across the periodic table. These trends are significant for understanding the behavior of elements, particularly in the context of atomic radius.
One major trend is how the atomic radius changes as you move down a group in the periodic table. Generally, the atomic radius increases as you descend through a group. This increase is due to the addition of more electron shells.
The greater number of shells means electrons are further away from the nucleus, causing a larger atomic size. Another trend is observed when moving across a period, typically from left to right.
In this direction, the atomic radius tends to decrease. This is because more protons and electrons are added, increasing the attraction between the nucleus and the electron cloud, pulling the electron cloud closer to the nucleus. Understanding these trends can help you make sense of why elements do not always behave as expected, such as why gallium is smaller than aluminum, when conventional wisdom would suggest the opposite.

Electron Shielding

Electron shielding is an essential concept to understand atomic size variations, especially when atomic radii seem to defy expected trends. Shielding refers to the process where inner electrons partially block the attraction between the nucleus and the outer electrons.

• As more inner electrons are present, the effective nuclear charge felt by the outermost electrons decreases.
• This decrease in effective nuclear charge can often lead to an increase in atomic size because the outer electrons are less tightly held by the nucleus.

However, not all electrons shield effectively. For instance, in gallium, the 3d-electrons do not shield the nuclear charge as effectively as other s- or p-electrons do.
This poor shielding by d-electrons results in a higher effective nuclear charge experienced by the outer electrons, drawing them closer to the nucleus and resulting in a smaller atomic radius compared to aluminum.

d-Orbitals

d-Orbitals play a peculiar role in altering atomic radii, largely due to their electron shielding properties. The presence of d-orbitals can sometimes lead to unexpected changes in an element's atomic properties. In gallium, the filling of 3d-orbitals is key to understanding its smaller atomic radius relative to aluminum.

• Although both d- and p-orbitals contribute to electron shielding, d-orbitals are less effective.
• This inefficiency causes the outer electrons to feel a stronger pull towards the nucleus because they aren't adequately "shielded" from the positive charge.

This is particularly evident as we move from aluminum to gallium.
Gallium's additional electrons occupy 3d-orbitals, which do not shield as well as 3p-electrons. As a result, despite gallium having more electrons and protons, its outer electrons are pulled closer to the nucleus, making its atomic radius smaller than that of aluminum.