Question

Would you expect \(\mathrm{Fe}^{3+}\) or \(\mathrm{Fe}^{2+}\) to be the stronger Lewis acid? Explain.

Short answer

We can conclude that \(\mathrm{Fe}^{3+}\) is the stronger Lewis acid compared to \(\mathrm{Fe}^{2+}\) as it has a higher positive charge and smaller size. These factors contribute to a stronger attraction toward electron pairs, making \(\mathrm{Fe}^{3+}\) a better electron pair acceptor and therefore a stronger Lewis acid.

Step-by-step solution

List the ions and their charge

We have two ions: \(\mathrm{Fe}^{3+}\) and \(\mathrm{Fe}^{2+}\). The \(\mathrm{Fe}^{3+}\) ion has a +3 charge, and the \(\mathrm{Fe}^{2+}\) ion has a +2 charge.

Understand the concept of Lewis acids

Lewis acids are species that can accept electron pairs from a Lewis base, thus forming a coordinate covalent bond. Stronger Lewis acids have a higher affinity for accepting these electron pairs.

Compare the charges of the ions and their effect in Lewis acidity

Higher positive charge indicates a greater capacity for attracting electron pairs and may increase the Lewis acidity. In our case, \(\mathrm{Fe}^{3+}\) has a higher charge than \(\mathrm{Fe}^{2+}\) (+3 versus +2), which suggests that \(\mathrm{Fe}^{3+}\) could be the stronger Lewis acid.

Compare the size of the ions and their effect on Lewis acidity

When comparing ions of the same element, the ions with higher positive charges will generally be smaller due to stronger effective nuclear charge. This stronger charge pulls the electrons in the ion closer to the nucleus, making it easier for the ion to accept electron pairs. Therefore, \(\mathrm{Fe}^{3+}\) will be smaller in size than \(\mathrm{Fe}^{2+}\), which also favors \(\mathrm{Fe}^{3+}\) being the stronger Lewis acid.

Conclude which ion is the stronger Lewis acid

Taking into account both charge and size, we can conclude that \(\mathrm{Fe}^{3+}\) is the stronger Lewis acid compared to \(\mathrm{Fe}^{2+}\). The higher charge and smaller size of \(\mathrm{Fe}^{3+}\) will allow it to have a stronger attraction toward electron pairs, making it a better electron pair acceptor and therefore a stronger Lewis acid.