Question

Fill in the blanks. During electrolysis of molten \(\mathrm{PbBr}_{2}\), ___ fumes are liberated at the anode.

Short answer

Answer: Bromine gas fumes.

Step-by-step solution

Understanding Electrolysis

Electrolysis is a process in which an electric current is used to drive a non-spontaneous redox reaction. In this case, we are given a molten \(\mathrm{PbBr}_{2}\) compound undergoing electrolysis. In electrolysis, oxidation events happen at the anode, while reduction events occur at the cathode.

Find the ions present in molten PbBr2

Molten \(\mathrm{PbBr}_{2}\) is a liquid, and consists of ions. The ions present will be the cations \(\mathrm{Pb}^{2+}\) and anions \(\mathrm{Br}^{-}\), since \(\mathrm{PbBr}_{2}\) dissociate into its constituent ions when in molten state.

Determine the reactions happening at the Anode

At the anode during electrolysis, the oxidation reaction will take place. The anode is negatively charged and as a result it will attract the negatively charged ions, that is, the bromide ions (\(\mathrm{Br^{-}}\)). Bromide ions will be oxidized into bromine gas. The oxidation half-reaction can be written as: \(\mathrm{2Br^{-}} \rightarrow \mathrm{Br}_{2} + 2 e^{-}\)

Determine the fumes liberated at the Anode

Now, we can see that the anode produces \(\mathrm{Br}_{2}\) gas in the oxidation half-reaction. So, the fumes liberated in the electrolysis of molten \(\mathrm{PbBr}_{2}\) at the anode are: Bromine gas fumes.

Key concepts

Oxidation reaction

An oxidation reaction is a fundamental process where a substance loses electrons during a chemical reaction. In electrolysis, the oxidation reaction specifically occurs at the anode. Here, negatively charged ions lose electrons. The process can be visualized like this:

• The ion loses electrons and changes its oxidation state.
• For example, in the electrolysis of molten \( \mathrm{PbBr}_{2} \), bromide ions (\( \mathrm{Br}^- \)) lose electrons to form bromine gas (\( \mathrm{Br}_{2} \)).

This transformation is crucial in electrolysis as it drives the overall chemical reaction by releasing electrons that travel through the circuit.

Anode reaction

The anode reaction is central to understanding what happens during electrolysis. The anode is the positive electrode where oxidation occurs. During the electrolysis of molten \( \mathrm{PbBr}_{2} \):

• The negatively charged bromide ions (\( \mathrm{Br}^- \)) move towards the anode.
• These ions lose electrons, a process known as oxidation.
• The reaction can be simplified as \( 2\mathrm{Br}^- \rightarrow \mathrm{Br}_{2} + 2e^- \).

As a result, electrons are released and bromine gas is formed. The electrons are then available to flow through the external circuit, allowing the continuous flow of electricity.

Bromine gas liberation

Bromine gas liberation is the visible product of the anode reaction during electrolysis. When \( \mathrm{PbBr}_{2} \) is subjected to an electric current, bromide ions (\( \mathrm{Br}^- \)) are oxidized at the anode:

• Electrons are removed from the bromide ions.
• This results in the formation of bromine gas (\( \mathrm{Br}_{2} \)), which is released as fumes.
• The reaction is a straightforward transformation: \( 2\mathrm{Br}^- \rightarrow \mathrm{Br}_{2} + 2e^- \).

The liberation of bromine gas is an essential clue in identifying the type of reaction occurring and confirming the composition of the solution being electrolyzed.

PbBr2 electrolysis

The electrolysis of \( \mathrm{PbBr}_{2} \) involves the breakdown of this compound into its elements through the application of electricity. The components naturally dissociate into their ions when molten:

• Lead ions (\( \mathrm{Pb}^{2+} \)) and bromide ions (\( \mathrm{Br}^- \)).
• At the cathode, lead ions are reduced to form lead metal.
• At the anode, bromide ions undergo oxidation to produce bromine gas.

This process not only separates the components into valuable elements but also illustrates the principle of electrolysis whereby an electrical input facilitates a chemical change. Understanding this process is key for applications in metallurgy and chemical manufacturing.