Question

Hydrazine is somewhat toxic. Use the following half-reactions to explain why household bleach (highly alkaline solution of sodium hypochlorite) should not be mixed with household ammonia or glass cleansers that contain ammonia. \(\mathrm{ClO}^{-}+\mathrm{H}_{2} \mathrm{O}+2 \mathrm{e}^{-} \longrightarrow 2 \mathrm{OH}^{-}+\mathrm{Cl}^{-} \quad \mathscr{E}^{\circ}=0.90 \mathrm{~V}\) \(\mathrm{N}_{2} \mathrm{H}_{4}+2 \mathrm{H}_{2} \mathrm{O}+2 \mathrm{e}^{-} \longrightarrow 2 \mathrm{NH}_{3}+2 \mathrm{OH}^{-} \quad \mathscr{E}^{\circ}=-0.10 \mathrm{~V}\)

Short answer

Combining household bleach (containing \(\mathrm{ClO}^{-}\)) with ammonia or ammonia-containing products results in a redox reaction that produces toxic ammonia gas (\(\mathrm{NH}_{3}\)): \[\mathrm{ClO}^{-}+\mathrm{N}_{2}\mathrm{H}_{4}+3\mathrm{H}_{2}\mathrm{O}\longrightarrow\mathrm{NH}_{3}+\mathrm{Cl}^{-}+4\mathrm{OH}^{-}\] Ammonia poses health risks if inhaled or mixed with other chemicals, hence it is not advisable to mix these substances.

Step-by-step solution

Identify the oxidizing and reducing agents

In order to find out which compound is the oxidizing agent and which is the reducing agent, we should look at the standard reduction potentials \(\mathscr{E}^{\circ}\) given for both half-reactions: - For the first half-reaction, \(\mathrm{ClO}^{-}\) is reduced (gains electrons) to form \(\mathrm{Cl}^{-}\) with \(\mathscr{E}^{\circ}=0.90 \mathrm{~V}\) - For the second half-reaction, \(\mathrm{N}_{2}\mathrm{H}_{4}\) is reduced (gains electrons) to form \(\mathrm{NH}_{3}\) with \(\mathscr{E}^{\circ}=-0.10 \mathrm{~V}\) Since a higher standard reduction potential indicates a greater tendency to be reduced, \(\mathrm{ClO}^{-}\) will act as the oxidizing agent, while \(\mathrm{N}_{2}\mathrm{H}_{4}\) will act as the reducing agent.

Balance the half-reactions

Before combining the half-reactions, make sure both are balanced in terms of elements and charges. The given half-reactions are already balanced.

Combine the half-reactions to form a redox reaction

To form the redox reaction, multiply the half-reactions so that the number of electrons transferred in both reactions is equal, and then add them together. Given half-reactions: \[\mathrm{ClO}^{-}+\mathrm{H}_{2} \mathrm{O}+2 \mathrm{e}^{-} \longrightarrow 2 \mathrm{OH}^{-}+\mathrm{Cl}^{-}\] \[\mathrm{N}_{2} \mathrm{H}_{4}+2 \mathrm{H}_{2} \mathrm{O}+2 \mathrm{e}^{-} \longrightarrow 2 \mathrm{NH}_{3}+2 \mathrm{OH}^{-}\] Since the number of electrons transferred in both half-reactions is already equal (2 e⁻), add them together to form the redox reaction: \[\mathrm{ClO}^{-}+\mathrm{N}_{2}\mathrm{H}_{4}+3\mathrm{H}_{2}\mathrm{O}\longrightarrow\mathrm{NH}_{3}+\mathrm{Cl}^{-}+4\mathrm{OH}^{-}\]

Analyze the reaction products

After combining the half-reactions, we get the following redox reaction: \[\mathrm{ClO}^{-}+\mathrm{N}_{2}\mathrm{H}_{4}+3\mathrm{H}_{2}\mathrm{O}\longrightarrow\mathrm{NH}_{3}+\mathrm{Cl}^{-}+4\mathrm{OH}^{-}\] The redox reaction between household bleach (\(\mathrm{ClO}^{-}\)) and hydrazine (\(\mathrm{N}_{2}\mathrm{H}_{4}\), present in ammonia-containing products) produces ammonia (\(\mathrm{NH}_{3}\)), chloride ions (\(\mathrm{Cl}^{-}\)), and hydroxide ions (\(\mathrm{OH}^{-}\)). Ammonia is a toxic and irritating gas that poses health risks if inhaled or mixed with other chemicals. Therefore, it is not advisable to mix household bleach with ammonia or ammonia-containing products like glass cleansers.

Key concepts

Standard Reduction Potential

The standard reduction potential (SRP), denoted as \(\mathscr{E}^\circ\), is a measure of the tendency of a chemical species to gain electrons and thereby be reduced. In redox reactions, electrons are transferred from one substance to another, and SRP values provide insight into the direction of electron flow. A more positive SRP means a substance is more likely to gain electrons, whereas a more negative SRP indicates a substance is less inclined to be reduced.

For example, in the exercise, chlorate ion \(\mathrm{ClO}^{-}\) has an SRP of +0.90 V, suggesting it readily accepts electrons. In contrast, hydrazine \(\mathrm{N}_{2}\mathrm{H}_{4}\) has an SRP of -0.10 V, implying it's less likely to gain electrons. Understanding SRPs is crucial for predicting redox behavior and determining which reactants will serve as oxidizing or reducing agents in a chemical reaction.

Oxidizing and Reducing Agents

In every redox process, two key players are involved: the oxidizing agent and the reducing agent. The oxidizing agent is the chemical that gains electrons (is reduced) during the reaction, while the reducing agent is the chemical that loses electrons (is oxidized).

The exercise provided a clear demonstration: \(\mathrm{ClO}^{-}\), with its higher SRP, acts as the oxidizing agent because it is more eager to gain electrons. On the other hand, hydrazine \(\mathrm{N}_{2}\mathrm{H}_{4}\), with its lower SRP, serves as the reducing agent by losing electrons. Recognizing these agents allows chemists to predict the flow of electrons and thereby the course of the reaction.

Chemical Safety

Redox reactions can produce unexpected and sometimes hazardous products, emphasizing the significance of chemical safety. Ammonia, for instance, is a common component of many household cleaners, but it can form dangerous compounds when mixed with other chemicals like bleach. The redox reaction in our exercise demonstrates the potential for generating toxic ammonia gas when bleach, a product containing \(\mathrm{ClO}^{-}\), reacts with hydrazine, a component found in some glass cleaners.

Therefore, it's crucial to understand both the chemicals you're working with and their potential interactions. Always consult material safety data sheets (MSDS) and follow established safety protocols to minimize risks associated with hazardous chemical reactions.

Balancing Redox Reactions

Balancing redox reactions can be complex, but it's a vital skill for ensuring that all atoms and charges are accounted for in chemical equations. Start by identifying the number of electrons transferred in the oxidation and reduction half-reactions. The next step is to balance these electrons to maintain electrical neutrality. In our exercise, both half-reactions involve a transfer of two electrons, so they are already balanced and can be combined directly.

When balancing more complicated equations, it might be necessary to multiply the half-reactions by appropriate factors to equalize the number of electrons transferred. After that, add the half-reactions together, ensuring that elements and charges are balanced. This process ensures that the redox reaction is accurately represented and adheres to the law of conservation of mass.