CBSE Class 11 Chemistry Chapter 11: The p-Block Elements NCERT Solutions

The variation in oxidation states for the given groups is explained as follows:

(i) Boron (B) to Thallium (Tl) - Group 13 Elements

The general electronic configuration for Group 13 elements is $ns^2\; np^1$. Consequently, the most common oxidation state exhibited by these elements is expected to be +3, involving all three valence electrons. However, this trend is not uniform across the group.

• Boron (B) and Aluminium (Al): These elements predominantly exhibit the +3 oxidation state.
• Gallium (Ga), Indium (In), and Thallium (Tl): These heavier elements show both +1 and +3 oxidation states.

A significant trend observed is that as we move down the group from Ga to Tl, the +1 oxidation state becomes increasingly stable, while the +3 oxidation state becomes less stable. This phenomenon is attributed to the **inert pair effect**. The two electrons in the ns orbital are held more tightly by the nucleus and are less likely to participate in bonding. This effect becomes more pronounced in heavier elements.

For example:

• Ga(+1) is relatively unstable.
• In(+1) is fairly stable.
• Tl(+1) is very stable and is the most common oxidation state for Thallium.

Conversely, the stability of the +3 oxidation state decreases as we move down the group.

Summary Table for Group 13:

(ii) Carbon (C) to Lead (Pb) - Group 14 Elements

The general electronic configuration for Group 14 elements is $ns^2\; np^2$. The expected common oxidation state is +4, utilizing all four valence electrons. However, similar to Group 13, the stability of oxidation states changes down the group.

• Carbon (C) and Silicon (Si): These elements primarily exhibit the +4 oxidation state.
• Germanium (Ge), Tin (Sn), and Lead (Pb): These heavier elements display both +2 and +4 oxidation states.

The **inert pair effect** is also responsible for the observed trend in Group 14. As we move down the group, the stability of the +2 oxidation state increases, while the stability of the +4 oxidation state decreases. This is because the two ns electrons become increasingly reluctant to participate in bonding.

Therefore, while C and Si predominantly exist in the +4 state, Ge, Sn, and Pb show a greater tendency to exist in the +2 state. For instance, Pb(+2) is more stable than Pb(+4).

Summary Table for Group 14:

Boron (B) and Thallium (Tl) are both members of Group 13 of the periodic table. A key trend in this group is that the stability of the +1 oxidation state increases, while the stability of the +3 oxidation state decreases as we move down the group, primarily due to the inert pair effect.

Boron, being at the top of the group, readily exhibits its +3 oxidation state, and compounds like BCl₃ are stable in this state. In contrast, Thallium is at the bottom of the group. For Thallium, the +1 oxidation state is significantly more stable than the +3 oxidation state because of the strong inert pair effect.

Consequently, TlCl₃ is less stable than BCl₃. In fact, Tl(+3) compounds are strong oxidizing agents and tend to revert to the more stable Tl(+1) state. Therefore, BCl₃ is considerably more stable than TlCl₃.

Boron trifluoride (BF₃) behaves as a Lewis acid because the central boron atom is electron-deficient. Let's look at the electronic configuration and bonding:

• The electronic configuration of Boron is $1s^2\; 2s^2\; 2p^1$. It has 3 valence electrons.
• In BF₃, Boron forms three covalent bonds with three Fluorine atoms. Each bond involves one electron from Boron and one from Fluorine.
• This means Boron uses all its 3 valence electrons to form bonds.
• As a result, the Boron atom in BF₃ is surrounded by only 6 valence electrons (3 bonds x 2 electrons/bond), instead of the stable octet of 8 electrons.

Since the boron atom has an incomplete octet, it is electron-deficient and has a strong tendency to accept a pair of electrons to achieve a stable octet configuration. According to the Lewis definition, a substance that can accept an electron pair is called a Lewis acid. Therefore, BF₃ acts as a Lewis acid.

The vacant p-orbital on the Boron atom is available to accept the electron pair.