Question

Which of the following is a disproportionation reaction? \((c): 2 \mathrm{Fe}_{(6)}+3 \mathrm{H}_{2} \mathrm{O}_{(D} \stackrel{\Delta}{\longrightarrow} \mathrm{Fe}_{2} \mathrm{O}_{3(j)}+3 \mathrm{H}_{2(g)}\) d) \(2 \mathrm{H}_{2} \mathrm{O}_{(\bar{D}}+2 \mathrm{~F}_{2(\mathrm{e})} \rightarrow 4 \mathrm{HF}_{\left(a_{q}\right)}+\mathrm{O}_{2(\mathrm{~g})}\)

Short answer

Option (d) is a disproportionation reaction.

Step-by-step solution

Understanding Disproportionation Reactions

A disproportionation reaction is a type of redox reaction in which a single species undergoes both oxidation and reduction, forming two different products. In a disproportionation reaction, one element of a single reactant is simultaneously oxidized and reduced.

Analyzing Option (c)

Examine the reaction: \(2 \mathrm{Fe} + 3 \mathrm{H}_2\mathrm{O} \stackrel{\Delta}{\longrightarrow} \mathrm{Fe}_2\mathrm{O}_3 + 3\mathrm{H}_2\). In this reaction, the oxidation state of iron (Fe) changes from 0 in Fe to +3 in \(\mathrm{Fe}_2\mathrm{O}_3\), while hydrogen (H) remains in the same oxidation state: 0 in \(\mathrm{H}_2\) and +1 in \(\mathrm{H}_2\mathrm{O}\). Since the iron is only getting oxidized and no single species is both reduced and oxidized, this is not a disproportionation reaction.

Analyzing Option (d)

Examine the reaction: \(2 \mathrm{H}_2\mathrm{O} + 2 \mathrm{F}_2 \rightarrow 4 \mathrm{HF} + \mathrm{O}_2\). Looking at the oxidation states: oxygen starts at -2 in \(\mathrm{H}_2\mathrm{O}\) and ends up at 0 in \(\mathrm{O}_2\) (oxidation); fluorine starts at 0 in \(\mathrm{F}_2\) and ends up at -1 in HF (reduction). This reaction shows the same element, oxygen, being both oxidized and reduced. Therefore, this is a disproportionation reaction.

Key concepts

Redox Reaction

One of the compelling concepts in chemistry is the redox reaction, short for reduction-oxidation reaction. It's a process where electrons are transferred between two substances. Because electrons carry energy, redox reactions are often energy-transforming events. One substance gains electrons, undergoing reduction, while the other loses electrons, which is oxidation.

Every redox reaction consists of two parts: the oxidation half and the reduction half. These simultaneous processes are eloquently represented in redox equations, showing how electrons hop from one reactant to another. It's crucial for students to recognize that the number of electrons lost is equal to the number of electrons gained, maintaining the principle of charge conservation.

Real-life Instances of Redox Reactions

In real life, redox reactions are everywhere. From the rusting of iron to the metabolic processes in your body, they are the driving force behind numerous natural and man-made phenomena. Understanding how to balance redox reactions, recognizing the movement of electrons, and predicting the outcome of these reactions is vital in mastering this topic.

Oxidation States

The oxidation state, also known as oxidation number, is a helpful concept that lets you keep track of electron movement in redox reactions. It can be thought of as a bookkeeping tool used to tally how 'electron-rich' or 'electron-poor' an atom is in a compound or a chemical reaction. This helps us determine which atoms are oxidized and which are reduced.

The rules for assigning oxidation states are straightforward:

• For atoms in their elemental form, the oxidation state is always zero.
• For monoatomic ions, it's equal to the charge of the ion.
• For molecules, the oxidation states should sum up to zero; for polyatomic ions, they should sum up to the ion's charge.

Getting comfortable with calculating oxidation states is a bedrock skill in chemistry, which helps students analyze chemical reactions and understand the changes their components undergo.

Chemical Reactions in Chemistry

Chemistry fundamentally deals with chemical reactions, the processes leading to the transformation of one set of substances into new substances. Each reaction is governed by specific laws and falls into categories, such as synthesis, decomposition, single replacement, double replacement, and redox reactions, which include a subset known as disproportionation reactions.

In disproportionation reactions, an unusual twist occurs as a single reactant gets oxidized and reduced simultaneously, producing distinct products with different oxidation states from the original element. Educationally, grasping the nature and the outcomes of these reactions paves the way for students to understand more complex concepts like catalysis, electrochemistry, and metabolic pathways in biological systems.

Catalysts in Chemical Reactions

Catalysts are substances that speed up chemical reactions without being consumed. They work by lowering the activation energy required for the reaction to proceed. Understanding the role of catalysts is essential, as they are indispensable in industrial processes for producing various chemicals, pharmaceuticals, and consumer products.