Question

Describe the change in hybridization (if any) of the \(\mathrm{Al}\) atom in this reaction: $$ \mathrm{AlCl}_{3}+\mathrm{Cl}^{-} \longrightarrow \mathrm{AlCl}_{4}^{-} $$

Short answer

The hybridization of the aluminum atom changes from \(sp^2\) to \(sp^3\) during the reaction.

Step-by-step solution

Identify the original hybridization state

In the beginning, the aluminum atom in the \(AlCl_3\) compound is making three bonds, which correspond to the three chlorine atoms. Since it has 3 sigma bonds and no lone pairs of electrons, this implies that the hybridization state of the aluminum atom is \(sp^2\).

Identify the final hybridization state

After receiving an extra \(Cl^-\) ion, the aluminum atom is now part of a \(AlCl_4^-\) ion and it is making four bonds, which correspond to the four chlorine atoms. In this case, considering 4 sigma bonds and no lone pairs of electrons, this implies that the hybridization state of the aluminum atom is now \(sp^3\).

Describe the change in hybridization

Comparing the initial state of hybridization (\(sp^2\)) with the final state of hybridization (\(sp^3\)), it can be concluded that the hybridization of the aluminum atom has changed from \(sp^2\) to \(sp^3\) as a result of this reaction.

Key concepts

Chemical Bonding

Chemical bonding refers to the connections between atoms that hold molecules together. In the realm of chemistry, there are different types of bonds, including ionic, covalent, and metallic bonds. These bonds are critical for understanding how atoms associate with one another to form compounds.
In the given reaction, \[ \text{AlCl}_3 + \text{Cl}^- \rightarrow \text{AlCl}_4^-,\]aluminum (Al) and chlorine (Cl) exhibit covalent bonding. Covalent bonds are characterized by the sharing of electron pairs between atoms. This type of bond forms when the electron clouds overlap, creating a stable balance of attractive and repulsive forces. As a part of this process, hybridization can occur, which involves the mixing of atomic orbitals to form new hybrid orbitals suitable for pairing electrons during chemical bond formation.

Sigma Bonds

Sigma bonds (\( \sigma \)-bonds) are the primary type of covalent bond, forming when atomic orbitals overlap head-on. They are the strongest type of covalent bond and are characterized by the direct overlap of orbitals along the internuclear axis. In terms of molecular geometry, sigma bonds play a crucial role since they allow the atoms in a molecule to rotate freely around the bond axis.
In the case of \( \text{AlCl}_3 \) and \( \text{AlCl}_4^- \), the aluminum atom forms sigma bonds with chlorine atoms. Initially, \( \text{Al} \) forms three sigma bonds with \( \text{Cl} \) atoms, aligning with its \( \text{sp}^2 \) hybridization state. When an additional chloride ion (\( \text{Cl}^- \)) is introduced, it forms a fourth sigma bond, which reshuffles the hybrid orbitals to an \( \text{sp}^3 \) configuration. Thus, in the transition from \( \text{AlCl}_3 \) to \( \text{AlCl}_4^- \), the aluminum atom's involvement in sigma bonds crucially reflects the change in hybridization.

Electron Pairs

In the context of chemical reactions and bonding, electron pairs play an integral role in determining the structure and behavior of molecules. Electron pairs can be categorized into two types: bonding pairs, which are shared between two atoms to form covalent bonds, and lone pairs, which are not involved in bonding.
For the aluminum atom in \( \text{AlCl}_3 \), it initially forms three bonding pairs with three chlorine atoms, resulting in three sigma bonds. There are no lone pairs on the aluminum atom in this structure. Upon accepting an extra chlorine ion to form \( \text{AlCl}_4^- \), a new bonding pair is established, allowing aluminum to maintain four bonding electron pairs, corresponding to its four sigma bonds.
This shift exemplifies how electron pair distribution dictates changes in molecular geometry and hybridization, transitioning from \( \text{sp}^2 \) (three bonding pairs) to \( \text{sp}^3 \) (four bonding pairs).

Lewis Structures

Lewis structures are diagrammatic representations that demonstrate the bonding between atoms within a molecule and the lone pairs of electrons that may exist. They are fundamental in visualizing the configuration of valence electrons around atoms.
For \( \text{AlCl}_3 \), the Lewis structure would depict the central aluminum atom single-bonded to three chlorine atoms, with no lone pairs of electrons. This layout underscores its \( \text{sp}^2 \) hybridization state.
When an additional \( \text{Cl}^- \) ion bonds with aluminum, the Lewis structure of \( \text{AlCl}_4^- \) comes into play. Here, the aluminum atom is shown forming single bonds with four chlorine atoms, reflecting its \( \text{sp}^3 \) hybridization, while leaving a negative charge on the molecule to indicate the extra electron from the \( \text{Cl}^- \) ion. This graphical detail helps in comprehending the structural change and the transformation of hybridization that results from the reaction.