Question

A solution of chlorine gas in water is used in chemical analysis to test for the presence of \(\mathrm{Br}^{-}\) and \(\mathrm{I}^{-}\) ions in solution. \(\mathrm{Cl}_{2}\) is able to oxidize easily these two ions to the elemental forms \(\mathrm{Br}_{2}\) and \(\mathrm{I}_{2},\) which may then be identified by their colors. Write a balanced oxidation-reduction equation for each of these processes.

Short answer

The balanced oxidation-reduction equations for the processes involving \(\mathrm{Br}^-\) and \(\mathrm{I}^-\) ions are: 1) 2 \(\mathrm{Br}^- + \mathrm{Cl}_{2} \rightarrow \mathrm{Br}_{2} +\) 2 \(\mathrm{Cl}^-\) 2) 2 \(\mathrm{I}^- + \mathrm{Cl}_{2} \rightarrow \mathrm{I}_{2} +\) 2 \(\mathrm{Cl}^-\)

Step-by-step solution

Write the unbalanced half-reactions

For the oxidation of \(\mathrm{Br}^-\) ions by \(\mathrm{Cl}_{2}\), we have: Oxidation: \(\mathrm{Br}^- \rightarrow \mathrm{Br}_{2}\) Reduction: \(\mathrm{Cl}_{2} \rightarrow \mathrm{Cl}^-\) For the oxidation of \(\mathrm{I}^-\) ions by \(\mathrm{Cl}_{2}\), we have: Oxidation: \(\mathrm{I}^- \rightarrow \mathrm{I}_{2}\) Reduction: \(\mathrm{Cl}_{2} \rightarrow \mathrm{Cl}^-\)

Balance the atoms in the half-reactions

For the \(\mathrm{Br}^-\) oxidation reaction, add a coefficient of 2 in front of the \(\mathrm{Br}^-\) to balance the atoms: 2 \(\mathrm{Br}^- \rightarrow \mathrm{Br}_{2}\) For the \(\mathrm{Cl}_{2}\) reduction reaction, add a coefficient of 2 in front of the \(\mathrm{Cl}^-\) to balance the atoms: \(\mathrm{Cl}_{2} \rightarrow\) 2 \(\mathrm{Cl}^-\) Similarly, for the \(\mathrm{I}^-\) oxidation reaction, add a coefficient of 2 in front of the \(\mathrm{I}^-\) to balance the atoms: 2 \(\mathrm{I}^- \rightarrow \mathrm{I}_{2}\) The reduction reaction for \(\mathrm{Cl}_{2}\) remains the same as for the \(\mathrm{Br}^-\) reaction.

Balance the charges in the half-reactions

For both the \(\mathrm{Br}^-\) and \(\mathrm{I}^-\) oxidation reactions, the charges are already balanced, so no changes are needed. For the \(\mathrm{Cl}_{2}\) reduction reaction, add a coefficient of 2 in front of the \(\mathrm{e}^-\) on the left side of the reaction to balance the charges: \(\mathrm{Cl}_{2} + 2 \mathrm{e}^- \rightarrow\) 2 \(\mathrm{Cl}^-\)

Combine the half-reactions

To form the overall balanced redox equation for the \(\mathrm{Br}^-\) oxidation process, add the balanced half-reactions together: 2 \(\mathrm{Br}^- + \mathrm{Cl}_{2} + 2 \mathrm{e}^- \rightarrow \mathrm{Br}_{2} +\) 2 \(\mathrm{Cl}^-\) + 2 \(\mathrm{e}^-\) The balanced oxidation-reduction equation for the process involving \(\mathrm{Br}^-\) ions is: 2 \(\mathrm{Br}^- + \mathrm{Cl}_{2} \rightarrow \mathrm{Br}_{2} +\) 2 \(\mathrm{Cl}^-\) Similarly, combine the half-reactions to form the overall balanced redox equation for the \(\mathrm{I}^-\) oxidation process: 2 \(\mathrm{I}^- + \mathrm{Cl}_{2} + 2 \mathrm{e}^- \rightarrow \mathrm{I}_{2} +\) 2 \(\mathrm{Cl}^-\) + 2 \(\mathrm{e}^-\) The balanced oxidation-reduction equation for the process involving \(\mathrm{I}^-\) ions is: 2 \(\mathrm{I}^- + \mathrm{Cl}_{2} \rightarrow \mathrm{I}_{2} +\) 2 \(\mathrm{Cl}^-\)

Key concepts

Balancing Redox Reactions

Understanding how to balance redox reactions is essential for students studying chemistry. Redox reactions involve the transfer of electrons between substances, causing changes in oxidation states. To begin balancing these reactions, split them into two half-reactions: oxidation and reduction. The oxidation half-reaction entails the loss of electrons, and the reduction half-reaction involves their gain. For instance, when balancing the reaction where chlorine gas oxidizes bromide ions, the reduction half-reaction is Cl₂ becoming Cl^-, and the oxidation half-reaction is Br^- becoming Br₂.

Once the half-reactions are written, balance the mass by ensuring that the same number of each atom is present on both sides of the equation. Then, balance the charge by adding electrons to one side of each half-reaction. For example, Cl₂ gains two electrons to form 2 Cl^-. After both half-reactions are balanced for mass and charge, they are combined to form the overall balanced redox equation. Throughout the process, pay close attention to maintaining the conservation of mass and charge, as every atom and electron must be accounted for.

Halogen Displacement Reactions

Halogen displacement reactions are a type of redox reaction where a more reactive halogen displaces a less reactive halogen from its compound. Reactivity in halogens decreases down the group in the periodic table, meaning chlorine can displace bromine and iodine, but bromine can't displace chlorine. In our exercise example, chlorine gas in water reacts with bromide (Br^-) and iodide (I^-) ions, showing chlorine's ability to oxidize these ions to form bromine (Br₂) and iodine (I₂), respectively.

Observing the color change that occurs during the reaction helps with qualitative analysis to confirm the presence of Br^- or I^- ions. Bromine water turns from colorless to orange, while iodine solution changes from colorless to brown. Recognizing these visible changes is an integral part of using halogen displacement reactions in chemical analysis.

Identification of Ions in Solution

The identification of ions in a solution is a crucial process in analytical chemistry. By applying chemical tests, we can detect the presence or absence of specific ions. The halogen displacement reaction shown in the exercise is a perfect illustration, where chlorine water is used as a reagent to test for bromide or iodide ions. If these ions are present, the solution will change color as elemental bromine or iodine forms due to the redox reaction.

In laboratory practice, other tests are often performed to confirm the identity of ions, such as the silver nitrate test for halides or the flame test for metal cations. These tests are designed to exploit unique chemical or physical properties of ions, such as distinctive colors or precipitates that form under certain conditions, providing evidence for the ions' presence in a sample.

Electron Transfer in Redox Reactions

Electron transfer is the key feature of redox reactions and dictates how the reaction proceeds. It involves the moving of electrons from one species to another, altering the oxidation states of the reactants. During a redox reaction, the species that loses electrons is oxidized, raising its oxidation number, whereas the species that gains electrons is reduced, lowering its oxidation number. In our example with bromide and iodide ions being oxidized by chlorine, the electrons are transferred from Br^- or I^- to chlorine atoms. Chlorine, gaining the electrons, is reduced, while the halide ions, losing electrons, are oxidized.

Understanding and identifying these electron transfers is fundamental when writing balanced redox equations. It ensures a grasp of how different reactants interact and the transformation they undergo during a chemical reaction. Clearly tracking electron movement also provides insight into the energetic aspects of reactions and how they may be harnessed in practical applications like batteries and electrolysis.