Question

In group \(4 \mathrm{~A}\) or 14 group of the extended form of the periodic table with increase in atomic number, the oxidizing power of tetravalent species increases in the order: (a) \(\mathrm{Ge}>\mathrm{Sn}>\mathrm{Pb}\) (b) \(\mathrm{Pb}>\mathrm{Sn}>\mathrm{Ge}\) (c) \(\mathrm{Pb}>\mathrm{Ge}>\mathrm{Sn}\) (d) \(\mathrm{Ge}>\mathrm{Pb}>\mathrm{Sn}\)

Short answer

The correct order is (b) Pb > Sn > Ge.

Step-by-step solution

Understanding Periodic Trends

The elements in Group 14 of the periodic table include Carbon (C), Silicon (Si), Germanium (Ge), Tin (Sn), and Lead (Pb). As you go down the group, the atomic number increases. The oxidizing power of tetravalent species refers to their ability to gain electrons in a +4 oxidation state.

Analyzing Oxidizing Power

In general, as you go down group 14, the stability of the +4 oxidation state decreases due to the inert pair effect, which makes the heavier elements like Lead (Pb) more stable in a +2 state than in a +4 state. Hence, Lead in the +4 state tends to be a stronger oxidizing agent as it prefers to gain electrons to reach the stable +2 state.

Ordering the Elements by Oxidizing Power

Based on the inert pair effect and the tendency of elements to prefer their lower oxidation states down the group, the order of increasing oxidizing power for the tetravalent state is Lead (Pb), Tin (Sn), and then Germanium (Ge). Thus, the order is Pb > Sn > Ge.

Selecting the Correct Option

From the given options, the sequence Pb > Sn > Ge corresponds to option (b). Therefore, the correct increasing order of the oxidizing power of the tetravalent species is Pb > Sn > Ge.

Key concepts

Oxidizing Power

In chemistry, oxidizing power refers to the ability of a chemical species to accept electrons. When talking about the oxidizing power in Group 14 elements in their tetravalent state, we focus on how well these elements can gain electrons when they are in a +4 oxidation state. As we move down Group 14 in the periodic table—covering elements such as Carbon (C), Silicon (Si), Germanium (Ge), Tin (Sn), and Lead (Pb)—we notice a change in their oxidizing capabilities. These elements can accept electrons to become oxidizing agents, especially when they prefer to be in a lower oxidation state due to other chemical effects. Understanding this concept helps predict the behavior of elements during chemical reactions, which is crucial in both laboratory and industrial chemical processes.

Group 14 Elements

Group 14 of the periodic table, also known as the carbon group, encompasses Carbon (C), Silicon (Si), Germanium (Ge), Tin (Sn), and Lead (Pb). These elements are characterized by having four electrons in their outer electron shell. This electron configuration allows them to form various chemical bonds and exhibit multiple oxidation states.

• As we move down the group from Carbon to Lead, there is an increase in atomic size and a decrease in electronegativity.
• This alteration also leads to a gradual weakening of bonds due to the larger atomic radii and lesser effective nuclear charge experienced by the outer electrons.

These trends play a pivotal role in determining the chemical reactivity and stability of different oxidation states of these elements, especially affecting their role in oxidation and reduction reactions.

Tetravalent Species

A tetravalent species is an atom or element that forms four chemical bonds, corresponding to a +4 oxidation state. In the context of Group 14, tetravalent compounds are particularly common, as these elements naturally have four electrons available for bonding. The stability of their tetravalent state is influenced by several factors:

• As we go from Carbon down to Lead, the stability of the +4 oxidation state diminishes.
• This is due to the "inert pair effect," where the s electrons become more resistant to participate in bonding for heavier elements.

This trend highlights why heavier elements in the group like Tin and Lead lean towards lower oxidation states more readily, leaving their tetravalent state with a more significant oxidizing power.

Inert Pair Effect

The inert pair effect is a significant concept explaining the chemical behavior of heavier p-block elements such as those found in Group 14. It describes the reluctance of the s-electron pair to participate in bonding as you move down a group in the periodic table. Essentially, as the atomic number increases, the s-electrons (part of the outer electron shell) become more stable and less reactive due to their penetration closer to the nucleus and increased shielding effect. This effect is particularly observable in heavier Group 14 elements:

• It causes Lead to prefer a +2 oxidation state over a +4 oxidation state because the s-electron pair remains "inert," or unreactive.
• The inert pair thus influences the oxidizing power of the element, making the tetravalent states of the heavier elements stronger oxidizers when they gain electrons to achieve the more stable +2 oxidation state.

Grasping the inert pair effect is essential for understanding why Lead is stronger as an oxidizing agent in its tetravalent form compared to its lighter counterparts in the group.