Question

Which would be considered a stronger Lewis acid: (a) \(\mathrm{BF}_{3}\) or \(\mathrm{BCl}_{3},\) (b) \(\mathrm{Fe}^{2+}\) or \(\mathrm{Fe}^{3+}\) ? Explain.

Short answer

BF3 is stronger than BCl3; Fe3+ is stronger than Fe2+.

Step-by-step solution

Understanding Lewis Acids

A Lewis acid is a chemical species that can accept an electron pair. The strength of Lewis acids depends on their ability to accept electron pairs. This can be influenced by factors such as electronegativity and the size of the central atom.

Comparing BF3 and BCl3

Both molecules, \(\text{BF}_3\) and \(\text{BCl}_3\), are triangular planar with empty p orbitals that can accept electron pairs. \(\text{BF}_3\) is a stronger Lewis acid because fluorine is more electronegative than chlorine. This high electronegativity in \(\text{BF}_3\) results in a stronger pull on the electrons toward fluorine, making boron more electron-deficient and more eager to accept an electron pair.

Comparing Fe2+ and Fe3+

The comparison between \(\text{Fe}^{2+}\) and \(\text{Fe}^{3+}\) revolves around the positive charge. \(\text{Fe}^{3+}\) has a higher positive charge and smaller ionic radius compared to \(\text{Fe}^{2+}\). The increased positive charge on \(\text{Fe}^{3+}\) makes it more electron-deficient, thus making it a stronger Lewis acid than \(\text{Fe}^{2+}\).

Key concepts

Electron Pair Acceptance

Lewis acids are unique because of their ability to accept electron pairs. This ability is what defines them in the world of chemical reactions. A Lewis acid is essentially an electron pair acceptor, meaning it seeks out atoms with extra electrons to share. This is crucial in reactions where new bonds are formed. The acceptance of electron pairs allows Lewis acids to participate in various reactions, often acting as a catalyst. The strength of a Lewis acid is measured by how readily it accepts an electron pair. There are several factors that influence this, including electronegativity and the shape of the molecule.

Electronegativity

Electronegativity is a property that describes how strongly an atom attracts electrons toward itself. In the context of Lewis acids, it plays a pivotal role in determining their strength. For instance, in comparing \( ext{BF}_3\) and \( ext{BCl}_3\), the electronegativity of the attached atoms is a key factor. Fluorine, being more electronegative than chlorine, attracts electrons more powerfully. This leaves boron in \( ext{BF}_3\) more electron-deficient and eager to accept electron pairs, making \( ext{BF}_3\) a stronger Lewis acid than \( ext{BCl}_3\). Higher electronegativity can increase the electron-withdrawing ability, thus enhancing the molecule's desire to accept electron pairs.

Ionic Radius

The ionic radius refers to the size of an ion. It can influence how easily a Lewis acid accepts electrons. Smaller ions have a higher charge density, which means they can exert a stronger pull on electrons. In the case of iron ions, \( ext{Fe}^{3+}\) has a smaller ionic radius than \( ext{Fe}^{2+}\). This smaller radius allows \( ext{Fe}^{3+}\) to have a more concentrated positive charge, enhancing its ability to pull in electron pairs, thus making it a stronger Lewis acid. The relationship between ionic radius and charge density is a critical aspect of understanding the function and strength of Lewis acids.

Charge

Charge, especially positive charge, significantly impacts an ion's ability as a Lewis acid. The more positively charged an ion is, the stronger Lewis acid it becomes. This is because a higher positive charge means a greater electron deficiency. For example, \( ext{Fe}^{3+}\) is more electron-deficient than \( ext{Fe}^{2+}\) because it has a higher positive charge. This electron deficiency makes \( ext{Fe}^{3+}\) more aggressive in seeking out electron pairs, increasing its potency as a Lewis acid. A high charge translates to a stronger attraction for negatively charged electron pairs, thus amplifying the Lewis acid behavior.

Triangular Planar Geometry

Certain molecules adopt a shape known as triangular or trigonal planar geometry. This shape is a foundational aspect that influences the behavior of Lewis acids like \( ext{BF}_3\) and \( ext{BCl}_3\). In a triangular planar structure, the atoms are arranged in a flat triangle around a central atom. This configuration allows available p orbitals to remain open and receptive to accepting electron pairs. The planar geometry facilitates the arrangement of shared electrons and contributes to the stability of the molecule after electron pair acceptance. Thus, molecular geometry plays a pivotal role in determining how and why a molecule functions as a Lewis acid.