Question

Ga is below \(\mathrm{Al}\) in the periodic table, but atomic radius of Ga is less than Al. It is because of (a) Lanthanoid contraction (b) Greater screening effect (c) Inert pair effect (d) None of these

Short answer

The atomic radius of Ga is less than Al due to lanthanoid contraction (Option a).

Step-by-step solution

Identifying the Trend

In this problem, we need to understand why the atomic radius of Gallium (Ga) is smaller than that of Aluminum (Al) even though Ga appears below Al in the periodic table. Typically, atomic radii increase down a group due to added electron shells.

Examining the Lanthanoid Contraction

Lanthanoid contraction refers to the steady decrease in the size of the lanthanoid ions as the atomic number increases. This phenomenon affects elements that are after the lanthanoids, like Ga, causing them to have a smaller atomic radius than expected.

Understanding the Other Options

The greater screening effect would typically lead to a larger atomic radius due to additional layers of electrons. The inert pair effect generally refers to the reluctance of 's' electrons to participate in bonding, which is not directly related to atomic radius changes between Al and Ga.

Concluding with the Correct Reason

Given that lanthanoid contraction affects elements following the lanthanoids and leads to a smaller atomic radius than their predecessors, it is the likely cause for Ga having a smaller atomic radius than Al.

Key concepts

Lanthanoid Contraction

Lanthanoid contraction is a fascinating concept that explains why certain elements don't follow the typical trends in the periodic table. The lanthanoids consist of 14 elements often filling the f-orbitals. As we move across the series, the atomic number increases, and a unique phenomenon occurs: the atomic and ionic radii gradually decrease. This contraction results because the added protons increase the effective nuclear charge, which pulls the electrons closer to the nucleus, making the atoms smaller. This shrinking effect continues beyond the lanthanoids themselves and affects subsequent elements in the periodic table. For example, Gallium, which is the element following the lanthanoids, exhibits a smaller than expected atomic radius compared to Aluminum. Thus, lanthanoid contraction plays a crucial role in explaining why Gallium has a smaller atomic radius than expected when compared to Aluminum.

Atomic Radius

The atomic radius is a fundamental concept in chemistry that refers to the size of an atom. It is considered the distance from the center of an atom’s nucleus to the outermost electron shell. As you move across a period in the periodic table, the atomic radius tends to decrease. This decrease is because the increase in the number of protons increases the effective nuclear charge, pulling electrons closer to the nucleus. Conversely, as you move down the groups, the atomic radius typically increases. This increase is due to the addition of electron shells. However, due to lanthanoid contraction, the expected trend sometimes shifts, as seen in the case of Gallium and Aluminum, where Gallium has a smaller radius despite being lower in the group.

Screening Effect

The screening effect is a crucial concept for understanding atomic structure. It describes how inner-shell electrons can shield outer electrons from the full attractive force of the nucleus. This shielding reduces the effective nuclear charge experienced by the outer electrons, resulting in a lessened pull of the nucleus on these electrons. Consequently, the atomic radius tends to increase with additional electron shells because the electrons don't feel the full nuclear charge. In the context of comparing Gallium and Aluminum, the screening effect was less significant because the lanthanoid contraction overrode the expected increase in atomic radius due to screening. Understanding these nuances highlights the complex interplay of factors that determine atomic size.

Inert Pair Effect

The inert pair effect is an important trend observed primarily in heavier elements of the p-block in the periodic table. It refers to the tendency of the outermost s-electrons to resist participating in bonding. This phenomenon often results in the lower oxidation states of these elements and occurs due to the poor shielding of the nuclear charge by f and d electrons. Although the inert pair effect doesn't directly affect atomic size, it is essential to know this effect exists as it influences the chemical reactivity and bonding of elements. Understanding this effect combined with other concepts like lanthanoid contraction and screening effect gives us a clearer picture of why certain elements behave counterintuitively in the periodic table.