Question

Stability of \(\mathrm{Ge}^{2+}, \mathrm{Sn}^{2+}\) and \(\mathrm{Pb}^{2+}\) is in order: (a) \(\mathrm{Ge}^{2+}>\mathrm{Sn}^{2+}>\mathrm{Pb}^{2+}\) (b) \(\mathrm{Sn}^{2+}>\mathrm{Ge}^{2+}>\mathrm{Pb}^{2+}\) (c) \(\mathrm{Pb}^{2+}>\mathrm{Sn}^{2+}>\mathrm{Ge}^{2+}\) (d) \(\mathrm{Sn}^{2+}>\mathrm{Pb}^{2+}>\mathrm{Ge}^{2+}\)

Short answer

(c) \( \mathrm{Pb}^{2+} > \mathrm{Sn}^{2+} > \mathrm{Ge}^{2+} \)

Step-by-step solution

Understanding Stability in the Group

Ge, Sn, and Pb are in the same group in the periodic table (Group 14), with increasing atomic numbers and masses from Ge to Pb. In general, for elements in the same group, the stability of lower oxidation states increases down the group due to inert pair effect.

Consider the Inert Pair Effect

The inert pair effect is the reluctance of the S-electrons to participate in bonding as you move down the group. This effect results in the stabilization of the lower oxidation state. Pb shows the most pronounced inert pair effect, making its +2 oxidation state more stable compared to Ge and Sn.

Determine the Stability Order

Based on the inert pair effect, the +2 state stability order increases as you move down the group: Ge < Sn < Pb. Those elements tend to favor their lower oxidation state more as you move downward, following this order: Pb's +2 is more stable than Sn's, which in turn is more stable than Ge's.

Key concepts

Oxidation States

Oxidation states are a concept used to reflect the degree of oxidation (or loss of electrons) of an atom in a chemical compound. Group 14 elements, including Ge, Sn, and Pb, can exhibit multiple oxidation states, most notably +2 and +4.

The +4 oxidation state is typically seen when elements are more electronegative or when there is greater energy available to persuade all valence electrons to participate in bonding. However, the +2 oxidation state becomes more stable down the group due to the inert pair effect, as fewer valence electrons participate in bonding. The reasoning behind this stability shift is that heavier elements like Sn and Pb have their outermost p-electrons readily involved, while s-electrons are less inclined to participate due to shielding and energy factors. The consequence is that for elements like lead, Pb, the +2 state is remarkably more stable compared to the lighter Ge or even Sn.

In practical terms, when predicting reactions or forming compounds, understanding these preferred oxidation states can prevent missteps and optimize chemical synthesis design or interpretation.

Inert Pair Effect

The inert pair effect is a key concept in understanding why certain oxidation states are more stable for heavier elements in the periodic table. In simple terms, this effect describes the tendency of the outermost s-electrons to remain non-bonding or inactive in heavier p-block elements. For Group 14, as you go from germanium (Ge) to lead (Pb), the increased atomic number means there's more nuclear charge and also more inner electrons shielding the valence electrons.

In elements like lead, the 6s electrons, although involved in bonding conceptually, remain largely non-participative. This is due to a combination of factors including increased relativistic effects for heavier elements. The inert pair effect, therefore, stabilizes the +2 oxidation state in Pb to a greater degree than in Ge or Sn.

The importance of the inert pair effect becomes evident when considering reactions and compound formation as it influences not just stability but the preferred oxidation state. This effect essentially dictates the chemical behavior of heavy Group 14 elements, making the understanding of bonding preferences crucial for effective chemical problem-solving.

Periodic Table Trends

Periodic table trends consider how properties change systematically as you move across periods or down groups. In Group 14, one such essential trend is the increase in atomic size and mass from Ge to Pb.

With larger atomic size, there is a reduced attraction between the nucleus and the outermost electrons, making it easier for these atoms to exhibit lower oxidation states as opposed to gaining energy to utilize s-electrons for bonding. Additional factors such as increased shielding also play a role in the observable increase in stability of the +2 oxidation state over the +4 as we descend the group.

Periodic trends, including ionization energy, electronegativity, and atomic radii, all highlight the changing chemical behavior as you move down Group 14. Understanding these trends aids in appreciating why certain chemical behaviors occur, such as Pb showing a stable +2 state while Ge is more commonly found in +4 states. These patterns offer insight crucial to predicting reactivity and properties of elements within the group.