Question

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

Short answer

In the given species, \(\mathrm{Fe}^{3+}\) has a higher positive charge and electron configuration of \([Ar]~3d^5\) compared to the \(\mathrm{Fe}^{2+}\) ion with an electron configuration of \([Ar]~3d^6\). Since both ions have empty 4s orbitals, the higher positive charge of \(\mathrm{Fe}^{3+}\) results in a stronger attraction to electron pairs. Therefore, \(\mathrm{Fe}^{3+}\) would be the stronger Lewis acid.

Step-by-step solution

Understanding Lewis acids

A Lewis acid is a chemical species that can accept an electron pair from a Lewis base in a chemical reaction. In other words, Lewis acids have an empty orbital that can accommodate a pair of electrons donated by a Lewis base.

Analyzing \(\mathrm{Fe}^{3+}\) and \(\mathrm{Fe}^{2+}\) ions

To determine which species is the stronger Lewis acid, we need to compare their charge and electron configurations. The electron configuration for a neutral Fe atom is \([Ar] 4s^2~3d^6\). When Fe loses 2 electrons and becomes the \(\mathrm{Fe}^{2+}\) ion, its electron configuration becomes \([Ar]~3d^6\). When it loses 1 more electron, forming the \(\mathrm{Fe}^{3+}\) ion, its electron configuration is \([Ar]~3d^5\).

Comparing charge and electron configurations

The \(\mathrm{Fe}^{3+}\) ion has one more positive charge than the \(\mathrm{Fe}^{2+}\) ion. The increased positive charge attracts the negatively charged electron pairs of Lewis bases more strongly. Additionally, both \(\mathrm{Fe}^{3+}\) and \(\mathrm{Fe}^{2+}\) have empty 4s orbitals that can accommodate electron pairs. Their 3d orbitals are partially filled, but the higher positive charge on the \(\mathrm{Fe}^{3+}\) ion results in a stronger attraction to electron pairs.

Conclusion

Based on the higher positive charge and similarity in electron configurations, we would expect the \(\mathrm{Fe}^{3+}\) ion to be the stronger Lewis acid, as it has a greater capacity to attract and accept electron pairs from Lewis bases.

Key concepts

Electron Configuration

Electron configuration is like a molecular fingerprint, showing how electrons are arranged around the nucleus of an atom. For iron (Fe), this arrangement starts from the lower energy levels, filling up to the higher levels. The basic electron configuration follows the pattern based on the Aufbau principle, which prioritizes filling lower energy orbitals first. The electron configuration for a neutral iron atom, Fe, is

• \([Ar] 4s^2 3d^6\),

where '[Ar]' stands for the configuration of argon, a noble gas that comes before iron on the periodic table. When iron loses electrons to form ions like

• \( ext{Fe}^{2+}\)
• \( ext{Fe}^{3+}\),

its electron configuration changes. The order of electron removal starts with the highest energy electrons, which are the ones in the 4s orbital. Therefore,

• \( ext{Fe}^{2+}\) has the configuration \([Ar] 3d^6\),
• while \( ext{Fe}^{3+}\) has the configuration \([Ar] 3d^5\).

The difference in these configurations, especially the electron count in the 3d orbital, can influence their chemical behaviors.

Fe Ions

Iron can form different cations by losing electrons, commonly resulting in

• \( ext{Fe}^{2+}\) and
• \( ext{Fe}^{3+}\) ions.

When iron loses electrons, it essentially gives away some of its outermost electrons in the 4s and 3d orbitals. Each lost electron corresponds to a positive charge, resulting in a cation. The

• \( ext{Fe}^{2+}\) ion, with a loss of two electrons,
• does not alter the fundamental structure as much as the
• \( ext{Fe}^{3+}\), which loses three electrons.

This impacts its overall charge and, consequentially, its chemical properties, such as interaction with other molecules. Notably,

• \( ext{Fe}^{3+}\)'s higher charge leads it to function as a stronger Lewis acid compared to
• \( ext{Fe}^{2+}\) because it can pull in electron pairs more strongly.

The ion type influences how iron behaves in various chemical reactions and its ability to form complex structures, illustrating the diversity of its chemistry.

Charge Attraction

The concept of charge attraction is crucial when analyzing Lewis acids, like

• \( ext{Fe}^{2+}\) and
• \( ext{Fe}^{3+}\) ions.

A positive ion will attract negatively charged particles, such as electron pairs from a Lewis base. This forms a key principle in the functioning of Lewis acids. In this context,

• \( ext{Fe}^{3+}\), with a higher positive charge than
• \( ext{Fe}^{2+}\),

exhibits a stronger force of attraction towards electron pairs. The magnitude of charge directly affects this pull; the greater the charge, the stronger the attraction. This makes

• \( ext{Fe}^{3+}\) more effective as a Lewis acid because it can more robustly engage with and accept electron pairs.
• These interactions also determine how easily these ions can engage in various acid-base reactions, impacting the overall reactivity and stability of iron complexes in the environment or industrial processes.

Charge attraction, therefore, plays a significant role in defining the efficiency and strength of metal ions like iron in dynamic chemical environments.