Question

Among the following, the correct statement(s) is(are) [A] $\mathrm{Al}{\left({\mathrm{CH}}_3\right)}_3$ has the three-centre two- electron bonds in its dimeric structure [B] ${\mathrm{BH}}_3$ has the three-centre two-electron bonds in its dimeric structure [C] ${\mathrm{AlCl}}_3$ has the three-centre two-electron bonds in its dimeric structure [D] The Lewis acidity of ${\mathrm{BCl}}_3$ is greater than that of ${\mathrm{AlCl}}_3$

Short answer

[B], [C], and [D] are correct statements. ${\mathrm{BH}}_3$ and ${\mathrm{AlCl}}_3$ form three-centre two-electron bonds in their dimeric forms, and ${\mathrm{BCl}}_3$ is a stronger Lewis acid than ${\mathrm{AlCl}}_3$.

Step-by-step solution

Examine Option [A]

Analyze the chemical bonding in $\mathrm{Al}({\mathrm{CH}}_3{)}_3$ and consider whether it forms a dimer with three-centre two-electron bonds. The three-center two-electron (3c-2e) bond is a type of chemical bond where three atoms share two electrons. This kind of bonding is observed in electron-deficient compounds, such as the case for ${\mathrm{B}}_2{\mathrm{H}}_6$. However, aluminum alkyls like $\mathrm{Al}({\mathrm{CH}}_3{)}_3$ do not typically form such bonds; instead, they are generally found as monomers due to the sufficient electron count in aluminum alkyl compounds.

Examine Option [B]

Consider the chemical bonding in ${\mathrm{BH}}_3$ and whether it forms a dimer with three-centre two-electron bonds. Diborane, ${\mathrm{B}}_2{\mathrm{H}}_6$, is a well-known molecule where boron forms three-centre two-electron bonds due to boron's electron deficiency. ${\mathrm{BH}}_3$ molecules dimerize to form ${\mathrm{B}}_2{\mathrm{H}}_6$ to complete their octet, thereby forming three-centre two-electron bonds.

Examine Option [C]

Analyze the chemical bonding in ${\mathrm{AlCl}}_3$ to determine if there are three-centre two-electron bonds in its structure. ${\mathrm{AlCl}}_3$ is known to form dimers, such as ${\mathrm{Al}}_2{\mathrm{Cl}}_6$, through chlorine bridging. The Al2Cl6 molecule features Al-Cl-Al bridges, which are a form of three-centre two-electron bond. Thus, option [C] is a correct statement.

Examine Option [D]

Compare the Lewis acidity of ${\mathrm{BCl}}_3$ to ${\mathrm{AlCl}}_3$. Lewis acidity is related to the ability of a compound to accept a pair of electrons. Boron trichloride, ${\mathrm{BCl}}_3$, has an incomplete octet and is more electron deficient than aluminum trichloride, ${\mathrm{AlCl}}_3$, due to the higher electronegativity of Boron compared to Aluminum. Therefore, ${\mathrm{BCl}}_3$ is a stronger Lewis acid than ${\mathrm{AlCl}}_3$ and option [D] is also correct.

Key concepts

Three-Centre Two-Electron Bonds

Three-centre two-electron (3c-2e) bonds are fascinating and slightly unconventional when compared to the typical bonds taught in introductory chemistry. They involve three atoms sharing two electrons, which is different from the more well-known two-centre two-electron bond where a pair of electrons is shared between two atoms.

These 3c-2e bonds are typically formed in compounds where there's not enough electrons to allow for traditional two-electron bonds between all adjacent atoms. An excellent example of a molecule featuring this type of bonding is diborane ((B_2H_6)), where two hydrogen atoms bridge between two boron atoms, each providing one electron to the bonding scheme. This bonding situation allows all three atoms to attain a stable electronic configuration, albeit in an unconventional way.

Dimerization

Dimerization is a process where two identical molecules join to form a larger complex, known as a dimer. This can occur due to various reasons, such as increasing the stability of the system or completing electron octets that can't be fulfilled in the monomer state.

In the context of chemistry, especially when discussing electron-deficient molecules, dimerization is a common strategy for atoms to achieve a full valence shell. The formation of (B_2H_6) from (BH_3) molecules is an archetypal example. Here, each boron atom has an electron deficiency, and by pairing up, they share the available electrons more effectively through the formation of three-centre two-electron bonds.

Lewis Acidity

Lewis acidity is an essential concept in chemistry related to a compound's ability to accept electrons. Lewis acids are electron-pair acceptors, and their strength is often gauged by the stability of the adduct formed when they accept an electron pair. This property is directly related to the electron configuration of a molecule and its resultant electron deficiency.

A high Lewis acidity implies a strong propensity to accept electron pairs, often seen in molecules with incomplete octets or central atoms that are electron deficient. For example, (BCl_3) tends to act as a stronger Lewis acid than (AlCl_3) due to boron having a higher electronegativity and a more significant electron deficiency.

Electron Deficiency

Electron deficiency occurs when an atom or molecule does not have enough valence electrons to form conventional two-electron covalent bonds with each of its neighbors. This scenario often leads to unique chemical behaviors, such as the formation of multi-centered bonds or an increased tendency to act as a Lewis acid.

Electron-deficient compounds are more likely to engage in reactions that allow them to complete their octet of electrons. For instance, typical electron-deficient molecules such as boron hydrides will often form compounds with 3c-2e bonds or react readily with Lewis bases to acquire the necessary electrons for stability. Understanding electron deficiency is crucial for predicting the reactivity and bonding patterns of compounds.