Question

Fluorine reacts with ice and results in the change: \(\mathrm{H}_{2} \mathrm{O}(\mathrm{s})+\mathrm{F}_{2}(\mathrm{~g}) \rightarrow \mathrm{HF}(\mathrm{g})+\mathrm{HOF}(\mathrm{g})\) Justify that this reaction is a redox reaction.

Short answer

The reaction is a redox reaction because fluorine undergoes both oxidation and reduction.

Step-by-step solution

Identify Changes in Oxidation Numbers

First, we need to identify the oxidation states of each element in both reactants and products. In the reactant, water (H₂O), hydrogen has an oxidation state of +1 and oxygen -2. In fluorine gas (F₂), fluorine is in its elemental state with an oxidation state of 0. In hydrogen fluoride (HF) and hypofluorous acid (HOF), fluorine has an oxidation state of -1 in HF and +1 in HOF, while hydrogen and oxygen retain their usual oxidation states of +1 and -2, respectively.

Determine Oxidation and Reduction

Here, oxygen in HOF changes from an oxidation state of -2 in H₂O to -2 after the reaction, with no change. In HF, fluorine changes its oxidation state from 0 in F₂ to -1 in HF, signifying a reduction. Conversely, in HOF, fluorine moves from 0 in F₂ to +1, indicating it is oxidized. Thus, fluorine is both reduced and oxidized simultaneously in this reaction, which is a characteristic of a redox reaction.

Conclusion: Classify the Reaction

Since fluorine is both oxidized and reduced in this reaction, this process can be classified as a redox reaction, specifically a disproportionation reaction, where the same element undergoes both oxidation and reduction.

Key concepts

Oxidation State

The oxidation state, also known as oxidation number, is a concept that helps chemists keep track of how electrons are distributed among elements in chemical compounds. It essentially represents the electrical charge an atom would have if all bonds were ionic.
Understanding oxidation states is crucial in determining how a chemical reaction proceeds, especially in redox reactions. In the reaction between fluorine and ice, determining the oxidation states of each element helps us figure out what is happening on an electron level.

• Each hydrogen ( H ) in water has an oxidation state of +1.
• Oxygen ( O ) in water normally has an oxidation state of -2.
• Fluorine ( F ), being in its elemental form as F₂, has an oxidation state of 0.
• In hydrogen fluoride ( HF ), the oxidation state of fluorine changes to -1.
• For hypofluorous acid ( HOF ), fluorine takes on an oxidation state of +1.

By examining these states, we see a transfer of electrons, which is central to understanding redox reactions.

Disproportionation Reaction

A disproportionation reaction is a special type of redox reaction. In this process, one element undergoes both oxidation and reduction simultaneously. This might sound a bit puzzling, but some elements are capable of taking on different oxidation states in the same reaction.

In the reaction between fluorine and ice, fluorine plays a unique role. It starts in its elemental state as F₂ with an oxidation state of 0. During the reaction, some fluorine atoms are reduced to an oxidation state of -1 in HF, while others are oxidized to an oxidation state of +1 in HOF.

• The simultaneous oxidation and reduction of fluorine, allowing it to achieve these two different oxidation states in one reaction, is what characterizes this process as a disproportionation reaction.

Appreciating this concept is vital since it explains why certain reactions in chemistry can be quite dynamic and complex.

Oxidation and Reduction

To fully grasp redox reactions, one must first understand what oxidation and reduction mean. These processes involve the transfer of electrons between atoms during a chemical reaction.
The term 'redox' is a portmanteau of reduction and oxidation, representing the linked nature of these processes.

• Oxidation refers to the loss of electrons by an atom, resulting in an increase in oxidation state.
• Reduction is the gain of electrons by an atom, leading to a decrease in oxidation state.

In our specific reaction involving fluorine and ice, we see:

• Fluorine in F₂ is reduced when it forms HF ( from 0 to -1 ), since it gains electrons.
• Conversely, fluorine is oxidized as it becomes HOF ( from 0 to +1 ), because it loses electrons.

Thus, understanding these electron movements is crucial for identifying how and why a reaction is classified as redox, and in this case, also a disproportionation reaction.