Question

Suggest products for the following (which are not balanced): (a) \(\left[\mathrm{ClO}_{3}\right]^{-}+\mathrm{Fe}^{2+}+\mathrm{H}^{+} \rightarrow\) (b) \(\left[\mathrm{IO}_{3}\right]^{-}+\left[\mathrm{SO}_{3}\right]^{2-} \rightarrow\) \((\mathrm{c})\left[\mathrm{IO}_{3}\right]^{-}+\mathrm{Br}^{-}+\mathrm{H}^{+} \rightarrow\)

Short answer

(a) Cl−, Fe3+, H2O; (b) I2, SO4^2−; (c) I2, Br2, H2O.

Step-by-step solution

Identify Each Reaction's Type

For each reaction, we must first understand the type of chemical reaction taking place. Here we see that the species involved commonly undergo redox reactions or are involved in acid reactions. We have chlorate ion ([ClO3]−) and iodate ion ([IO3]−), which often act as oxidizing agents.

Determine Potential Products for Reaction (a)

In the reaction \([\mathrm{ClO}_{3}]^{-}+\mathrm{Fe}^{2+}+\mathrm{H}^{+}\rightarrow\), the chlorate ion \([\mathrm{ClO}_{3}]^{-}\) can be reduced to form Cl− while the Fe2+ is oxidized to Fe3+. The presence of \(\mathrm{H}^+\) suggests that H2O or H+ can act as a product alongside the others, resulting in: \[\mathrm{Cl}^- + \mathrm{Fe}^{3+} + \mathrm{H}_2\mathrm{O}\]

Determine Potential Products for Reaction (b)

In \([\mathrm{IO}_{3}]^{-} + [\mathrm{SO}_{3}]^{2-} \rightarrow\), often iodate ion \([\mathrm{IO}_{3}]^{-}\) is reduced to iodine (I2), and \([\mathrm{SO}_{3}]^{2-}\) can be oxidized to sulfate \([\mathrm{SO}_{4}]^{2-}\). Hence, the possible products are: \[\mathrm{I}_2 + [\mathrm{SO}_{4}]^{2-}\]

Determine Potential Products for Reaction (c)

For \([\mathrm{IO}_{3}]^{-} + \mathrm{Br}^{-} + \mathrm{H}^{+} \rightarrow\), iodate ion \([\mathrm{IO}_{3}]^{-}\) might be reduced to I2 or \(\mathrm{I}^-\), and in this acidic environment, bromide ion \(\mathrm{Br}^{-}\) can be oxidized to Br2. Thus, the potential products include: \[\mathrm{I}_2 + \mathrm{Br}_2 + \mathrm{H}_2\mathrm{O}\]

Summarize Suggested Products

Each reaction has suggested products based on common redox reactions: (a) Cl−, Fe3+, H2O(b) \(\mathrm{I}_2\), \([\mathrm{SO}_{4}]^{2-}\)(c) \(\mathrm{I}_2\), \(\mathrm{Br}_2\), \(\mathrm{H}_2\mathrm{O}\)

Key concepts

Oxidizing Agents

In redox reactions, an oxidizing agent is a key player that facilitates the transfer of electrons by accepting them. This process causes another species to lose electrons, or be "oxidized." The oxidizing agents themselves undergo reduction because they gain electrons. For instance, in the original exercise,

• chlorate ion \([\mathrm{ClO}_{3}]^{-}\) acts as an oxidizing agent by accepting electrons from \(\mathrm{Fe}^{2+} \), leading to its reduction to \( \mathrm{Cl}^- \).
• Likewise, iodate ion \([\mathrm{IO}_{3}]^{-}\) receives electrons to become either \( \mathrm{I}_2 \) or \( \mathrm{I}^- \), depending on the specifics of the reaction.

This electron-accepting behavior effectively drives the redox process, making oxidizing agents indispensable in various biochemical and industrial reactions.

Redox Mechanisms

Redox mechanisms describe the detailed steps through which oxidation and reduction occur. These mechanisms are central to understanding how substances interact at a molecular level. The process involves:

• Reduction, where a species gains electrons. In our examples, chlorate and iodate ions are reduced as they gain electrons.
• Oxidation, where another species loses electrons, such as \( \mathrm{Fe}^{2+} \) becoming \( \mathrm{Fe}^{3+} \) or \( \mathrm{Br}^- \) converting to \( \mathrm{Br}_2 \).

The reactions are based on the principle of electron transfer, wherein the number of electrons lost in oxidation equals the number gained in reduction. This conservation ensures that redox reactions maintain electrical neutrality. Understanding redox mechanisms is crucial for studying processes like metabolism, corrosion, and energy storage.

Reaction Balancing

Balancing reactions is crucial because it ensures that all the atoms present in the reactants are accounted for in the products, maintaining mass and charge conservation. However, getting to a balanced reaction, especially in redox reactions, involves several key steps:

• Identify the oxidation and reduction components in the reaction.
• Calculate the change in oxidation numbers for the atoms involved to understand which atoms undergo oxidation and which ones undergo reduction.
• Add coefficients to ensure the number of atoms of each element is equal on both sides of the reaction equation, considering both atoms and charge.
• Finally, consider the medium (acidic or basic) and add \( \mathrm{H}^+ \) or \( \mathrm{OH}^- \) and water \( \mathrm{H}_2\mathrm{O} \) accordingly to balance hydrogen and oxygen atoms.

Balancing chemical reactions accurately is important for predicting reactants' behavior and products, as well as for designing and optimizing industrial chemical processes.