Question

In many reactions, the addition of \(\mathrm{AlCl}_{3}\) produces the same effect as the addition of \(\mathrm{H}^{+}\). (a) Draw a Lewis structure for \(\mathrm{AlCl}_{3}\) in which no atoms carry formal charges, and determine its structure using the VSEPR method. (b) What characteristic is notable about the structure in part (a) that helps us understand the acidic character of \(\mathrm{AlCl}_{3}\) ? (c) Predict the result of the reaction between \(\mathrm{AlCl}_{3}\) and \(\mathrm{NH}_{3}\) in a solvent that does not participate as a reactant. (d) Which acid-base theory is most suitable for discussing the similarities between \(\mathrm{AlCl}_{3}\) and \(\mathrm{H}^{+}\) ?

Short answer

The Lewis structure of AlCl3 is trigonal planar with 24 valence electrons. The notable characteristic of AlCl3 is its vacant p-orbital on the central Aluminum atom, which makes it prone to accepting electron pairs and acting as a Lewis acid. When AlCl3 reacts with NH3, it forms a coordinate-covalent bond, resulting in the complex AlCl3NH3. The most suitable acid-base theory to discuss the similarities between AlCl3 and H+ is the Lewis acid-base theory, as both prefer accepting electron pairs.

Step-by-step solution

(a) Lewis structure and VSEPR structure of AlCl3

: First, let's draw the Lewis structure of AlCl3. 1. Count the total number of valence electrons: Aluminum (Al) has 3 valence electrons and each Chlorine (Cl) has 7, so the total number of valence electrons is 3 + 3(7) = 24. 2. Place the least electronegative atom (Al) in the center. 3. Connect the Chlorine atoms to the center atom with single bonds. 4. Complete the octets of the Chlorine atoms by adding lone pairs. Now, we'll determine the structure using the VSEPR method. 1. Aluminum has 3 bond pairs and no lone pairs around it. According to the VSEPR theory, the structure will be trigonal planar with bond angles of approximately 120 degrees.

(b) Notable characteristic of the structure

: The notable characteristic of AlCl3 structure is that Aluminum has a vacant p-orbital as it has only 3 valence electrons. This vacant p-orbital in the central Aluminum atom makes it prone to accept an electron pair, which helps us to understand the acidic character of AlCl3 since it can accept electron pairs and act as a Lewis acid.

(c) Predicting the result of the reaction between AlCl3 and NH3

: AlCl3 can act as a Lewis acid, and ammonia (NH3) can act as a Lewis base since it has a lone pair of electrons on the nitrogen atom. Therefore, when AlCl3 reacts with NH3, it will accept the electron pair from NH3 and form a coordinate-covalent bond, resulting in the formation of the following complex: \[AlCl_3 + NH_3 \rightarrow AlCl_3NH_3\]

(d) Suitable acid-base theory

: The most suitable acid-base theory for discussing the similarities between AlCl3 and H+ is the Lewis acid-base theory. H+ acts as a Lewis acid by accepting electron pairs as it has an empty orbital, while AlCl3 can accept electron pairs due to its vacant p-orbital. Both H+ and AlCl3 prefer accepting electron pairs and act as Lewis acids rather than donating protons like in other acid theories, which explains their similar effect in many reactions.

Key concepts

VSEPR Theory

VSEPR theory stands for Valence Shell Electron Pair Repulsion theory. It helps to predict the shape of molecules by considering the repulsion between electron pairs around a central atom. In the case of \( \text{AlCl}_3 \), the theory suggests that since aluminum (Al) has three bond pairs and no lone pairs surrounding it, the molecular shape of \( \text{AlCl}_3 \) will be trigonal planar. This is due to the aim of minimizing repulsion between electron pairs. A trigonal planar structure has bond angles of approximately 120 degrees.
To visualize, you can think of the chlorine (Cl) atoms evenly distributed around the central aluminum atom, much like placing three items at equal angles on a flat disc. VSEPR theory is a valuable tool in chemistry as it helps us understand why molecules take certain shapes, influencing their properties and reactions.

Lewis Acids and Bases

Lewis acids and bases expand the concept of acid-base reactions beyond just protons. A Lewis acid is a substance that can accept a pair of electrons, while a Lewis base is one that can donate a pair of electrons.
In the context of \( \text{AlCl}_3 \), it acts as a Lewis acid. This is because it can accept an electron pair into its vacant p-orbital, a characteristic feature noticeable in its Lewis structure.
On the other hand, ammonia \( (\text{NH}_3) \) is a common Lewis base, as it has a lone pair of electrons on nitrogen ready to be shared. When \( \text{AlCl}_3 \) is reacted with \( \text{NH}_3 \), a coordinate-covalent bond is formed. The \( \text{NH}_3 \) molecule donates its lone pair to \( \text{AlCl}_3 \), forming the complex \( \text{AlCl}_3\text{NH}_3 \).
This demonstrates the versatility of the Lewis acid-base theory, where electron pair acceptance and donation dictate the reactions, broadening the concept of acidity and basicity.

Acid-Base Theories

Acid-base theories explain the behavior of acids and bases under different frameworks. While the Bronsted-Lowry theory defines acids as proton donors and bases as proton acceptors, the Lewis theory expands these definitions.
The Lewis theory describes acids as electron pair acceptors and bases as electron pair donors. In this context, \( \text{AlCl}_3 \) and \( \text{H}^+ \) both act as Lewis acids. \( \text{H}^+ \) has an empty orbital ready to accept electrons, similar to the vacant p-orbital in \( \text{AlCl}_3 \). This common feature explains their similar reactivity in certain reactions.
Choosing the correct theory allows us to better understand and predict reactions, especially when protons are not directly involved, highlighting the flexibility and wider applicability of the Lewis theory compared to other traditional acid-base concepts.