Problem 30 - Chemistry: The Central Science

Question

(a) Write the half-reaction that occurs at a hydrogen electrode in acidic aqueous solution when it serves as the anode of a voltaic cell. (b) The platinum electrode in a standard hydrogen electrode is specially prepared to have a large surface area. Why is this important? (c) Sketch a standard hydrogen electrode.

Step-by-Step Solution

Answer

(a) The half-reaction that occurs at a hydrogen electrode in acidic aqueous solution when it serves as the anode is given by: \(H_{2} (g) \rightarrow 2H^{+} (aq) + 2e^{-}\). (b) A large surface area for Platinum in a standard hydrogen electrode is important because it ensures efficient interaction of hydrogen gas with the electrode, facilitating electron transfer and providing a stable potential for an extended period. (c) A standard hydrogen electrode comprises a glass container with acidic electrolyte solution, Platinum wire with a Platinum plate having a large surface area for reaction, and hydrogen gas flowing onto the plate. The half-reaction occurs at the Platinum plate and H+ ions move within the solution.

Step 1: a) Half-reaction for hydrogen electrode serving as anode

When a hydrogen electrode serves as the anode, it means that the hydrogen gas (H2) is releasing electrons, leading to the oxidation process. In acidic aqueous solution, protons (H+) will be present, which can act as a hydrogen ion acceptor. The general oxidation half-reaction can be written as: H2 (g) -> 2H+ (aq) + 2e-

Step 2: b) Importance of large surface area for platinum in a standard hydrogen electrode

In a standard hydrogen electrode, the platinum electrode is specially prepared to have a large surface area. This is important because it ensures that more of the hydrogen gas can interact with the electrode for the redox reaction, increasing the efficiency of the electron transfer process. The large surface area not only enhances the reaction rate, but it also allows the electrode to maintain a stable potential for a more extended period, which is crucial for maintaining the standard state of the half-cell.

Step 3: c) Sketch a standard hydrogen electrode

To sketch a standard hydrogen electrode, follow these steps: 1. First, draw a vertical glass container that will hold the acidic electrolyte solution. 2. Inside the container, draw the Platinum wire, and at the bottom of the wire, draw a Platinum plate with a folded or rugged surface, representing a large surface area. 3. Add arrows to the sketch to indicate the flow of hydrogen gas (H2) onto the platinum plate. 4. Label the acidic electrolyte solution containing H+ (aq). 5. Show two half-reactions, one occurring at the Platinum plate, and the other occurring within the electrolyte solution. For the Platinum plate, the half-reaction should be H2 (g) -> 2H+ (aq) + 2e-, and within the electrolyte solution, show H+ ions with their movement represented by arrows. To summarize, the standard hydrogen electrode is composed of a Platinum wire with a Platinum plate inside a glass container holding an acidic electrolyte solution with hydrogen gas flowing onto the plate. This setup allows the redox reaction to occur efficiently due to the large surface area of the Platinum plate.

Key Concepts

Understanding Half-Reactions

In electrochemistry, half-reactions are essential to understanding how electrons are transferred in a complete reaction. A half-reaction represents either the oxidation or reduction process, occurring separately in an electrochemical cell. For the hydrogen electrode, when it acts as an anode, the half-reaction involves the oxidation of hydrogen gas. This means that hydrogen gas releases electrons in the following way:
\[ H_2 (g) \rightarrow 2H^+ (aq) + 2e^- \]
This reaction shows that each molecule of hydrogen gas splits into two protons and two electrons. Half-reactions are vital for balancing redox reactions in cells by showing the transfer of electrons.

The Oxidation Process Explained

Oxidation is a process where a substance loses electrons. In the context of the standard hydrogen electrode, oxidation occurs when hydrogen gas (\( H_2 \)) is converted into protons (\( H^+ \)) with the release of electrons. The key points are:

• Increases the charge of the substance because of the electron loss.
• Occurs at the anode of the voltaic cell.
• Essential for generating electric current as electrons move through the circuit.

Understanding oxidation is crucial for analyzing how chemical energy is converted into electrical energy in cells.

Role of the Platinum Electrode

Platinum is often chosen for electrodes because it is inert and does not react with other substances. For a standard hydrogen electrode, the platinum electrode is crucial for several reasons:

• Acts as a surface for the half-reaction.
• Because it's inert, it does not interfere with the reaction.
• It conducts electricity, helping electrons flow through the circuit.

The specially prepared large surface area enhances electron exchange, crucial for maintaining efficient electrode function.

Importance of Surface Area

The surface area of an electrode plays a significant role in the reaction's efficiency. A large surface area provides:

• More active sites for the reaction, increasing the overall rate.
• Improved stability of the electrode's potential.
• Enhanced exposure to reactants, which is vital for high reaction efficiency.

In a hydrogen electrode, having a large surface area is particularly important as it allows more hydrogen gas to participate in the reaction, making the electrode highly efficient.

Understanding Voltaic Cells

Voltaic cells, also known as galvanic cells, are used to convert chemical energy into electrical energy. They consist of two electrodes dipped in an electrolyte solution with each electrode undergoing a half-reaction. Key components of a voltaic cell include:

• Anode, where oxidation occurs.
• Cathode, where reduction happens.
• An electrolyte to allow ionic movement.
• A salt bridge to maintain neutrality by allowing ions to move across compartments.

Voltaic cells are essential for generating power in batteries and other electric devices, making understanding them fundamental in chemistry.