Problem 56 - Chemistry and Chemical Reactivity

Question

Which two of the following reactions are oxidationreduction reactions? Explain your answer briefly. Classify the remaining reaction. (a) \(\operatorname{CdCl}_{2}(\mathrm{aq})+\mathrm{Na}_{2} \mathrm{S}(\mathrm{aq}) \rightarrow \mathrm{CdS}(\mathrm{s})+2 \mathrm{NaCl}(\mathrm{aq})\) (b) \(2 \mathrm{Ca}(\mathrm{s})+\mathrm{O}_{2}(\mathrm{g}) \rightarrow 2 \mathrm{CaO}(\mathrm{s})\) (c) \(4 \mathrm{Fe}(\mathrm{OH})_{2}(\mathrm{s})+2 \mathrm{H}_{2} \mathrm{O}(\ell)+\mathrm{O}_{2}(\mathrm{g}) \rightarrow 4 \mathrm{Fe}(\mathrm{OH})_{3}(\mathrm{s})\)

Step-by-Step Solution

Answer

Reactions (b) and (c) are oxidation-reduction reactions; (a) is a double displacement reaction.

Step 1: Understanding Oxidation-Reduction Reactions

Oxidation-reduction (redox) reactions involve the transfer of electrons between two species. Typically, one species undergoes oxidation (loses electrons) and another undergoes reduction (gains electrons). To identify redox reactions, check for changes in oxidation states of elements from reactants to products.

Step 2: Analyzing Reaction (a)

For the reaction \(\operatorname{CdCl}_{2}(\mathrm{aq})+\mathrm{Na}_{2} \mathrm{S}(\mathrm{aq}) \rightarrow \mathrm{CdS}(\mathrm{s})+2 \mathrm{NaCl}(\mathrm{aq})\), analyze the oxidation states. Cd remains in the same oxidation state from \(Cd^{2+}\) to \(Cd^{2+}\) in CdS, and similarly, neither Na nor S undergoes change in oxidation state. Therefore, this reaction is not an oxidation-reduction reaction.

Step 3: Analyzing Reaction (b)

For the reaction \(2 \mathrm{Ca}(\mathrm{s})+\mathrm{O}_{2}(\mathrm{g}) \rightarrow 2 \mathrm{CaO}(\mathrm{s})\), examine the oxidation states. Ca starts with an oxidation state of 0 and ends with \(+2\), indicating oxidation. O starts with an oxidation state of 0 and ends with \(-2\), indicating reduction. Therefore, this reaction is an oxidation-reduction reaction.

Step 4: Analyzing Reaction (c)

For the reaction \(4 \mathrm{Fe}(\mathrm{OH})_{2}(\mathrm{s})+2 \mathrm{H}_{2} \mathrm{O}(\ell)+\mathrm{O}_{2}(\mathrm{g}) \rightarrow 4 \mathrm{Fe}(\mathrm{OH})_{3}(\mathrm{s})\), find the changes in oxidation states. Fe changes from +2 in \(\mathrm{Fe}(\mathrm{OH})_{2}\) to +3 in \(\mathrm{Fe}(\mathrm{OH})_{3}\), showing oxidation, while O is reduced from \(\mathrm{O}_{2}\) (0) to -2 in \(\mathrm{OH}^-\). Thus, this is also an oxidation-reduction reaction.

Step 5: Classifying the Remaining Reaction

Reaction (a) is not an oxidation-reduction reaction as there are no changes in oxidation states. It is instead a double displacement (or metathesis) reaction.

Key Concepts

Electron Transfer

In the world of chemistry, electron transfer forms the bedrock of redox reactions. When electrons are transferred between atoms or molecules, one species donates electrons, while another accepts them. This transfer is what fundamentally distinguishes redox reactions from other types of chemical reactions.

To identify such processes, always look for changes in the electron balance. The donor species, which loses electrons, undergoes oxidation. Meanwhile, the receiver, which gains electrons, experiences reduction.

This electron transfer results in a change of oxidation states for the involved species. Understanding who gains and loses electrons can help you pinpoint which reactions entail redox processes.

Oxidation States

Oxidation states, also known as oxidation numbers, indicate the degree of oxidation of an element in a chemical compound. These states are typically expressed in terms of integer values.

Each atom's oxidation state helps chemists determine whether electron transfer—indicative of a redox reaction—occurs during a chemical process.

• An increase in the oxidation state signifies oxidation (loss of electrons).
• A decrease indicates reduction (gain of electrons).

In general, elements in their elemental form have an oxidation state of zero. For instance, in reaction (b), calcium is oxidized from 0 to +2, and oxygen is reduced from 0 to -2. Calculating these changes helps track electron shifts during reactions.

Redox Reactions

Redox, or oxidation-reduction reactions, are characterized by the simultaneous occurrence of oxidation and reduction processes.

In any redox reaction, two half-reactions occur:

• Oxidation Half: The process where a species loses electrons.
• Reduction Half: The process where another species gains electrons.

In reaction (c), iron is oxidized, and oxygen is reduced. This dual nature of redox reactions is crucial for understanding the flow of electrons and the changes in oxidation states.

Balancing these half-reactions in terms of both mass and charge is key to accurately determining and predicting the course of redox reactions in chemical systems.

Double Displacement Reaction

A double displacement reaction, unlike a redox reaction, involves the exchange of ions between two compounds, resulting in the formation of two new compounds. In this process, no electrons are transferred.

In reaction (a), for instance, cadmium and sodium ions swap partners with chloride and sulfide ions without any change in oxidation states.

• This kind of reaction often occurs in aqueous solutions where two soluble salts react to form one soluble and one insoluble product.
• Typically results in the formation of a precipitate as seen with CdS forming from CdCl₂ and Na₂S.

Recognizing these reactions involves checking for ion exchanges without changes in oxidation numbers. They stand apart from redox reactions due to the absence of electron transfer.