Question

For the complete redox reactions given here, break down each reaction into its half-reactions, identify the oxidizing agent, and identify the reducing agent. (a) \(2 \mathrm{Sr}+\mathrm{O}_{2} \longrightarrow 2 \mathrm{Sr} \mathrm{O}\) (b) \(2 \mathrm{Li}+\mathrm{H}_{2} \longrightarrow 2 \mathrm{LiH}\) (c) \(2 \mathrm{Cs}+\mathrm{Br}_{2} \longrightarrow 2 \mathrm{CsBr}\) (d) \(3 \mathrm{Mg}+\mathrm{N}_{2} \longrightarrow \mathrm{Mg}_{3} \mathrm{~N}_{2}\)

Short answer

(a) Ox: Sr, Red: O₂; (b) Ox: Li, Red: H₂; (c) Ox: Cs, Red: Br₂; (d) Ox: Mg, Red: N₂.

Step-by-step solution

Break down (a) into Half-Reactions

For the reaction \(2 \mathrm{Sr}+\mathrm{O}_{2} \longrightarrow 2 \mathrm{SrO}\): - The oxidation half-reaction is: \(\mathrm{Sr} \rightarrow \mathrm{Sr}^{2+} + 2e^-\).- The reduction half-reaction is: \(\mathrm{O}_{2} + 4e^- \rightarrow 2\mathrm{O}^{2-}\).In the oxidation half-reaction, Sr loses electrons, and in the reduction half-reaction, \(\mathrm{O}_{2}\) gains electrons.

Identify Agents in (a)

In reaction (a), the oxidizing agent is \(\mathrm{O}_2\) because it gains electrons, while the reducing agent is \(\mathrm{Sr}\) as it loses electrons.

Break down (b) into Half-Reactions

For the reaction \(2 \mathrm{Li}+\mathrm{H}_{2} \longrightarrow 2 \mathrm{LiH}\): - The oxidation half-reaction is: \(\mathrm{Li} \rightarrow \mathrm{Li}^+ + e^-\).- The reduction half-reaction is: \(\mathrm{H}_{2} + 2e^- \rightarrow 2\mathrm{H}^-\).\(\mathrm{Li}\) loses electrons (oxidized), \(\mathrm{H}_{2}\) gains electrons (reduced).

Identify Agents in (b)

In reaction (b), \(\mathrm{H}_{2}\) is the oxidizing agent since it gains electrons, and \(\mathrm{Li}\) is the reducing agent as it loses electrons.

Break down (c) into Half-Reactions

For the reaction \(2 \mathrm{Cs}+\mathrm{Br}_{2} \longrightarrow 2 \mathrm{CsBr}\): - The oxidation half-reaction is: \(\mathrm{Cs} \rightarrow \mathrm{Cs}^+ + e^-\).- The reduction half-reaction is: \(\mathrm{Br}_{2} + 2e^- \rightarrow 2\mathrm{Br}^-\).\(\mathrm{Cs}\) is oxidized (loses electrons), \(\mathrm{Br}_{2}\) is reduced (gains electrons).

Identify Agents in (c)

In reaction (c), \(\mathrm{Br}_{2}\) is the oxidizing agent since it gains electrons, and \(\mathrm{Cs}\) is the reducing agent as it loses electrons.

Break down (d) into Half-Reactions

For the reaction \(3 \mathrm{Mg}+\mathrm{N}_{2} \longrightarrow \mathrm{Mg}_{3}\mathrm{~N}_{2}\): - The oxidation half-reaction is: \(\mathrm{Mg} \rightarrow \mathrm{Mg}^{2+} + 2e^-\).- The reduction half-reaction is: \(\mathrm{N}_{2} + 6e^- \rightarrow 2\mathrm{N}^{3-}\).\(\mathrm{Mg}\) is oxidized, \(\mathrm{N}_{2}\) is reduced.

Identify Agents in (d)

In reaction (d), \(\mathrm{N}_2\) is the oxidizing agent as it gains electrons, and \(\mathrm{Mg}\) is the reducing agent because it loses electrons.

Key concepts

Half-Reactions

In redox chemistry, reactions can be broken down into two simpler processes known as half-reactions. These half-reactions make it easier to analyze what happens to each reactant during a redoxic reaction.
Each redox reaction consists of an oxidation half-reaction and a reduction half-reaction.

• Oxidation involves the loss of electrons, and thus increases the oxidation state of the element.
• Reduction involves the gain of electrons, decreasing the oxidation state of the element.

By separating a redox reaction into its half-reactions, we can clearly see the transfer of electrons and determine which species is oxidized and which is reduced. This separation is crucial in balancing redox reactions and understanding their overall stoichiometry.

Oxidizing Agent

An oxidizing agent, or oxidant, plays a crucial role in redox reactions by accepting electrons from another substance. This process causes the oxidizing agent to be reduced, gaining electrons and typically reducing its oxidation state.
To identify the oxidizing agent in a reaction, look for the substance that gains electrons.
Here are a few key points related to oxidizing agents:

• The oxidizing agent undergoes a chemical reduction by gaining electrons.
• It usually contains an element with a high electronegativity or an element in a high oxidation state.
• Common oxidizing agents include oxygen ( ext{O}_2), halogens like chlorine ( ext{Cl}_2) and bromine ( ext{Br}_2), and compounds with metals in high oxidation states like permanganate ext{(MnO}_4^-).

Identifying the oxidizing agent can help determine the direction the reaction will proceed and the nature of the reactants involved.

Reducing Agent

The reducing agent, also known as the reductant, is the species in a redox reaction that donates electrons to another substance. In doing so, the reducing agent itself gets oxidized. This means it loses electrons and its oxidation state increases.
To spot the reducing agent in a redox reaction, look for the species that loses electrons.

• The reducing agent is oxidized as it loses electrons.
• Typically, reducing agents are elements or compounds that have low electronegativity.
• Examples of common reducing agents include alkali metals like lithium (Li) and sodium (Na), hydrogen gas ( ext{H}_2), and carbon monoxide (CO).

The reducing agent effectively powers the reduction of another substance by providing the necessary electrons for that process.
This process is pivotal in various chemical processes such as metabolism, battery functions, and any industrial chemical synthesis involving redox reactions.

Electron Transfer

Electron transfer is at the heart of redox reactions, acting as the driving force behind the chemical changes that occur. During a redox reaction, electrons are transferred from one reactant to another, leading to oxidation and reduction processes.
This electron transfer determines the flow of the reaction and is the key to understanding how energy is transformed during the process.

• Electron transfer occurs through oxidation (loss of electrons) and reduction (gain of electrons).
• Transfer of electrons results in changes to the oxidation states of the atoms involved.
• Balancing redox reactions requires ensuring that the number of electrons lost equals the number of electrons gained.

Understanding electron transfer in detail allows chemists to predict reaction outcomes, measure reaction kinetics, and design processes in fields ranging from biochemistry to electrochemistry. Recognizing this concept is vital for students to master the fundamentals of redox chemistry.