Question

Which of the following is not a Lewis acid? (a) \(\mathrm{AlCl}_{3} \cdot 6 \mathrm{H}_{2} \mathrm{O}\) (b) \(\mathrm{SnCl}_{4}\) (c) \(\mathrm{FeCl}_{3}\) (d) \(\mathrm{AlCl}_{3}\)

Short answer

(a) \( \text{AlCl}_3 \cdot 6 \text{H}_2\text{O} \) is not a Lewis acid.

Step-by-step solution

Understand Lewis Acids

A Lewis acid is a substance that can accept a pair of electrons to form a coordinate covalent bond. These acids usually have an empty orbital or are positively charged.

Analyze Each Option for Lewis Acidity

Evaluate each given compound to determine if it can be considered a Lewis acid. - (a) \( \text{AlCl}_3 \cdot 6 \text{H}_2\text{O} \) is a hydrated form of \( \text{AlCl}_3 \), which can accept electron pairs due to empty orbitals on the aluminum.- (b) \( \text{SnCl}_4 \) is known to act as a Lewis acid because tin can accept electron pairs via its empty \( p \) orbital.- (c) \( \text{FeCl}_3 \) is a Lewis acid because the central iron atom, being positively charged, can accept electron pairs.- (d) \( \text{AlCl}_3 \), similar to option (a), can accept electron pairs due to aluminum's empty orbitals.

Identify the Compound that is Not a Lewis Acid

Since all options (a), (b), and (c) show characteristics of being Lewis acids because they can accept electron pairs, the compound that is least likely to act as a Lewis acid based on independent analysis of potential external factors such as structural change due to hydration is (a) \( \text{AlCl}_3 \cdot 6 \text{H}_2\text{O} \).

Key concepts

Coordinate Covalent Bond

In chemistry, a coordinate covalent bond is a special kind of covalent bond. Unlike a typical covalent bond where each atom donates one electron to the electron pair, in a coordinate covalent bond, both electrons are donated by a single atom. This unique feature distinguishes it from regular covalent bonds.
The atom that contributes the electron pair is often called the "donor", while the atom that receives the pair is the "acceptor". This bond is crucial in the concept of Lewis acids and bases. When a Lewis base donates a pair of electrons to form a coordinate covalent bond with a Lewis acid, a stable product is formed.

• The bond itself is chemically identical to a normal covalent bond once it's formed.
• Coordinate covalent bonds often involve metal ions as the electron pair acceptors.
• This type of bond helps explain the reactivity and binding properties of certain compounds.

Understanding coordinate covalent bonds is fundamental in grasping how substances like metals interact dynamically with ligands through electron pair donation.

Electron Pair Acceptor

At the heart of the Lewis acid concept is the ability to accept an electron pair. A Lewis acid is defined as an electron pair acceptor. This ability to accept electron pairs makes these substances highly reactive and essential in many chemical processes.

• Electron pair acceptors often lack electrons, meaning they either have empty orbitals or a positive charge that creates a strong pull on electrons.
• Metals, such as those in groups 13 and 14 of the periodic table, are common electron pair acceptors.
• Sometimes, even positively charged non-metals can act as electron pair acceptors.

This concept is significant because it underlies many chemical reactions and interactions, allowing for the formation of complex structures through coordinate covalent bonding. Understanding this can help explain the behavior of acids like \( \text{AlCl}_3 \) and \( \text{SnCl}_4 \), which are classic examples of electron pair acceptors.

Empty Orbitals

Empty orbitals play a key role in the ability of a substance to act as a Lewis acid. An empty orbital is simply a vacant space where an electron pair can be accommodated. This characteristic enables the formation of a coordinate covalent bond with a Lewis base.

• Metals, especially in their ionized forms, often have empty orbitals readily available for bonding.
• Presence of empty orbitals makes the substance hungry for electrons, increasing its reactivity.
• Empty orbitals can be found in various energy levels, such as \( p \), \( d \), or \( f \), depending on the atom.

The role of empty orbitals is fundamental in processes like catalysis and complex formation. For instance, the empty \( p \) orbitals of \( \text{AlCl}_3 \) allow it to form strong coordinate covalent bonds with Lewis bases. This ability to accept electrons into empty orbitals is a defining characteristic of many Lewis acids.

Positively Charged Species

Positively charged species are almost inherently positioned as electron pair acceptors because of their lack of electrons. The positive charge creates an attraction for negative particles, such as electrons, which these species desperately seek to stabilize themselves.

• Many metal ions are positively charged and can serve as Lewis acids because of this electron deficiency.
• The positive charge enhances the tendency of the ion or molecule to capture electrons from donors.
• Depending on their nature, these species may have varying strengths in accommodating electron pairs.

For example, \( \text{FeCl}_3 \) is a positively charged species where iron can accept electron pairs because of its partial positive character due to its oxidation state. The presence of positive charge in these species often heightens their reactivity and the potential to engage in electron-pair interactions, allowing them to act effectively as Lewis acids in many chemical reactions.