Question

When an \(\mathrm{Fe}^{3+}\) salt is dissolved in water, the solution becomes acidic due to formation of \(\mathrm{Fe}\left(\mathrm{H}_{2} \mathrm{O}\right)_{5} \mathrm{OH}^{2+}\) and \(\mathrm{H}_{3} \mathrm{O}^{+} .\) The overall process involves both Lewis and Brónsted-Lowry acidbase reactions. Write the equations for the process.

Short answer

The overall process: \[ \text{Fe}^{3+} + 6\text{H}_2\text{O} \rightleftharpoons \text{Fe}(\text{H}_2\text{O})_5\text{OH}^{2+} + \text{H}_3\text{O}^+ \]

Step-by-step solution

- Dissolve \(\text{Fe}^{3+}\) in Water

When \(\text{Fe}^{3+}\) ions are dissolved in water, they hydrate to form \(\text{Fe}(\text{H}_2\text{O})_6^{3+}\). This is the hydration process of \(\text{Fe}^{3+}\) ions. \[ \text{Fe}^{3+} + 6\text{H}_2\text{O} \rightarrow \text{Fe}(\text{H}_2\text{O})_6^{3+} \]

- Lewis Acid-Base Reaction

Next, \(\text{Fe}(\text{H}_2\text{O})_6^{3+}\) acts as a Lewis acid because it can accept an electron pair, and the water molecules surrounding it can act as Lewis bases. This complex can facilitate the release of a proton (H⁺) from a water molecule within the complex. \[ \text{Fe}(\text{H}_2\text{O})_6^{3+} \rightarrow \text{Fe}(\text{H}_2\text{O})_5\text{OH}^{2+} + \text{H}^+ \]

- Brönsted-Lowry Acid-Base Reaction

The released proton (H⁺) then interacts with another water molecule in the solvent, forming a hydronium ion (\text{H}_3\text{O}^+). This step demonstrates the Brönsted-Lowry acid-base reaction where \(\text{H}_2\text{O}\) acts as a Brönsted-Lowry base by accepting a proton. \[ \text{H}^+ + \text{H}_2\text{O} \rightarrow \text{H}_3\text{O}^+ \]

- Overall Reaction

Combining the reactions from the previous steps, the overall process can be written as: \[ \text{Fe}^{3+} + 6\text{H}_2\text{O} \rightleftharpoons \text{Fe}(\text{H}_2\text{O})_5\text{OH}^{2+} + \text{H}_3\text{O}^+ \] This balanced equation shows both the Lewis acid-base interaction (\text{Fe}^{3+} + 6\text{H}_2\text{O} \rightarrow \text{Fe}(\text{H}_2\text{O})_6^{3+}) and the Brönsted-Lowry acid-base reaction (\text{H}^+ + \text{H}_2\text{O} \rightarrow \text{H}_3\text{O}^+).

Key concepts

Lewis Acid-Base Reaction

In a Lewis acid-base reaction, acids are defined as electron pair acceptors, while bases are electron pair donors. When \(\text{Fe}(\text{H}_2\text{O})_6^{3+}\) forms after dissolving iron(III) ion (\(\text{Fe}^{3+}\)) in water, it acts as a Lewis acid. This is because \(\text{Fe}^{3+}\) can accept electron pairs from the water molecules surrounding it. The equation representing this process is: \[ \text{Fe}^{3+} + 6\text{H}_2\text{O} \rightarrow \text{Fe}(\text{H}_2\text{O})_6^{3+} \] Here:

• \text{Fe}^{3+}\ is the Lewis acid as it accepts electron pairs.
• Water molecules act as Lewis bases donating electron pairs to \text{Fe}^{3+}\.

Understanding Lewis acid-base reactions helps in analyzing complex formation in solution chemistry.

Brönsted-Lowry Acid-Base Reaction

Brönsted-Lowry theory defines acids as proton (\(\text{H}^+\)) donors and bases as proton acceptors. After \(\text{Fe}(\text{H}_2\text{O})_6^{3+}\) forms, it can lose a proton to produce \(\text{Fe}(\text{H}_2\text{O})_5\text{OH}^{2+}\). In this step, \(\text{H}_2\text{O}\) actually participates by donating a proton: \[ \text{Fe}(\text{H}_2\text{O})_6^{3+} \rightarrow \text{Fe}(\text{H}_2\text{O})_5\text{OH}^{2+} + \text{H}^+ \] Next, this proton (\(\text{H}^+\)) combines with another water molecule to create a hydronium ion (\(\text{H}_3\text{O}^+\)): \[ \text{H}^+ + \text{H}_2\text{O} \rightarrow \text{H}_3\text{O}^+ \]
Important points to note:

• \(\text{H}_2\text{O}\) acts as a Brönsted-Lowry base by accepting a proton to form \(\text{H}_3\text{O}^+\).
• This interaction illustrates proton transfer, crucial for acid-base reactions.

Hydration Process

The hydration process involves the attachment of water molecules to ions in solution. For iron(III) ion (\text{Fe}^{3+}), this process is critical. When \text{Fe}^{3+} is dissolved in water, it becomes surrounded by six water molecules to form \text{Fe}(\text{H}_2\text{O})_6^{3+}. This process is essential for converting individual ions into a hydrated complex. The equation is: \[ \text{Fe}^{3+} + 6\text{H}_2\text{O} \rightarrow \text{Fe}(\text{H}_2\text{O})_6^{3+} \]

• \text{Fe}^{3+} attracts water molecules due to its high charge density.
• Hydration helps stabilize the ion in the solution.

This hydration is vital for understanding how ions behave in aqueous environments and is a precursor to further chemical interactions.

Iron(III) Ion Chemistry

Iron(III) ion (\text{Fe}^{3+}) has significant implications in chemistry due to its high charge and small size. When dissolved in water, \text{Fe}^{3+} undergoes hydration to form the complex \text{Fe}(\text{H}_2\text{O})_6^{3+}, which further reacts by losing protons. These reactions make solutions containing \text{Fe}^{3+} acidic. The overall chemical reaction can be written as: \[ \text{Fe}^{3+} + 6\text{H}_2\text{O} \rightleftharpoons \text{Fe}(\text{H}_2\text{O})_5\text{OH}^{2+} + \text{H}_3\text{O}^+ \] Key points:

• \text{Fe}^{3+} forms a hydrated complex.
• Releases protons, increasing the acidity of the solution.
• This makes \text{Fe}^{3+} salts important for studying acid-base reactions in aqueous solutions.

Understanding iron(III) ion chemistry helps explain the behavior of many iron-containing compounds in water.

Hydronium Ion

The hydronium ion (\text{H}_3\text{O}^+) is crucial in acid-base chemistry. It forms when a proton (\text{H}^+) from an acid combines with a water molecule (\text{H}_2\text{O}). For example, in the reaction: \[ \text{H}^+ + \text{H}_2\text{O} \rightarrow \text{H}_3\text{O}^+ \] Important facts about the hydronium ion include:

• Represents the actual form of ions in acidic solutions.
• Plays a key role in pH calculations and acid strength.
• Illustrates the Brönsted-Lowry acid concept.

In the context of iron(III) ion reactions, \(\text{H}_3\text{O}^+\) formation signifies how solutions turn acidic. Understanding hydronium ions helps in grasping how acids and bases behave in different environments.