Question

Which substance loses electrons and which substance gains electrons in this reaction? $$ 16 \mathrm{Fe}(\mathrm{s})+3 \mathrm{~S}_{8}(\mathrm{~s}) \rightarrow 8 \mathrm{Fe}_{2} \mathrm{~S}_{3}(\mathrm{~s}) $$

Short answer

Iron loses electrons (reducing agent), sulfur gains electrons (oxidizing agent).

Step-by-step solution

Identify the oxidation numbers in reactants

Start by determining the oxidation states of iron (\( \text{Fe} \)) and sulfur (\( \text{S} \)) in the reactants. In its elemental form, the oxidation state of \( \text{Fe} \) and \( \text{S}_8 \) is 0.

Identify the oxidation numbers in products

Next, identify the oxidation states in \( \text{Fe}_2\text{S}_3 \). Iron (\( \text{Fe} \)) in \( \text{Fe}_2\text{S}_3 \) is in the +3 oxidation state and sulfur (\( \text{S} \)) is in the -2 oxidation state.

Determine changes in oxidation states

Compare the oxidation states of the elements from reactants to products. Iron goes from 0 to +3, indicating it has lost electrons. Sulfur goes from 0 to -2, indicating it has gained electrons.

Identify oxidizing and reducing agent

The substance that loses electrons (iron) is the reducing agent, and the substance that gains electrons (sulfur) is the oxidizing agent.

Key concepts

Oxidation States

Understanding oxidation states is crucial in redox reactions, as they help us track electron transfer between atoms. Oxidation states indicate the degree of oxidation of an element. They are essentially imaginary charges assigned to an atom within a compound, assuming all bonds are 100% ionic, with electron pairs entirely belonging to the more electronegative atom.

For an element in its pure form, such as iron ( Fe ) or sulfur ( S_8 ), the oxidation state is zero. When atoms form compounds, their oxidation states change to reflect gained or lost electrons during bonding. In the reaction given in the problem, iron changes from an oxidation state of 0 to +3, showing that it loses three electrons per atom when it forms Fe_2S_3 . Similarly, each sulfur atom changes its oxidation state from 0 to -2, indicating a gain of two electrons per atom.

It's all about following these charge changes to determine what ends up losing or gaining electrons in the redox process.

Oxidizing Agent

In redox chemistry, the oxidizing agent is the substance that gains electrons and, as a result, causes another substance to be oxidized. During the course of a reaction, the oxidizing agent undergoes reduction.

In the given reaction, sulfur ( S_8 ) acts as the oxidizing agent. Initially, in its elemental form, sulfur has an oxidation state of 0. Upon forming the compound Fe_2S_3 , the sulfur atoms end up with an oxidation state of -2. This change indicates that sulfur gains electrons—specifically, it gains a total of six electrons when eight sulfur atoms combine to form the product.

• Oxidizing agents accept electrons.
• They cause the oxidation of other compounds.
• Their oxidation state decreases as they gain electrons.

Understanding what role an oxidizing agent plays helps in balancing redox reactions and predicting the outcome of chemical processes.

Reducing Agent

The reducing agent in a reaction is the substance that donates electrons to another species, thus causing that species to be reduced. In doing so, the reducing agent itself becomes oxidized.

In this exercise, iron ( Fe ) serves as the reducing agent. Initially in its elemental state, iron has an oxidation state of 0. Upon forming Fe_2S_3 , it acquires an oxidation state of +3. This transformation indicates that iron loses electrons—specifically, each iron atom loses three electrons in this reaction.

• Reducing agents donate electrons.
• They cause the reduction of other substances.
• Their oxidation state increases as they shed electrons.

Recognizing the role of a reducing agent is key to understanding the flow of electrons in a redox reaction and predicting how elements will interact under different conditions.