Question

(a) Write the half-reaction that occurs at a hydrogen electrode in acidic aqueous solution when it serves as the cathode of a voltaic cell (b) Write the half-reaction that occurs at a hydrogen electrode in acidic aqueous solation when it serves as the anode of a voltaic cell. (c) What is standard about the standard hydrogen electrode?

Short answer

(a) The half-reaction that occurs at a hydrogen electrode in acidic aqueous solution when it serves as the cathode of a voltaic cell is: \(2H^+ + 2e^- \rightarrow H_2\). (b) The half-reaction that occurs at a hydrogen electrode in acidic aqueous solation when it serves as the anode of a voltaic cell is: \(H_2 \rightarrow 2H^+ + 2e^-\). (c) The standard hydrogen electrode is a reference electrode with a standard reduction potential of 0 Volts by definition. It consists of a platinum wire coated with platinum black, dipped into a hydrochloric acid solution of unit activity, and exposed to hydrogen gas at 1 atm pressure and 298 K temperature. The standard electrode potential for other electrodes is determined by comparing their potential against the standard hydrogen electrode.

Step-by-step solution

Identify the half-reaction for reduction (cathode)

In a voltaic cell, the cathode is where reduction occurs. For a hydrogen electrode in acidic aqueous solution, we have hydrogen ions (H^+), which will be reduced to hydrogen gas (H2).

Balance the half-reaction

Balance the hydrogen atoms by combining two hydrogen ions (H^+) and adding the required number of electrons to the reactant side of the reaction equation: \(2H^+ + 2e^- \rightarrow H_2\) (b) Half-reaction at the hydrogen electrode as an anode in a voltaic cell

Identify the half-reaction for oxidation (anode)

In a voltaic cell, the anode is where oxidation occurs. For a hydrogen electrode in acidic aqueous solution, hydrogen gas (H2) will lose electrons and be oxidized to hydrogen ions (H^+).

Balance the half-reaction

Balance the hydrogen atoms by dividing the hydrogen gas into two hydrogen ions (H^+) and adding the required number of electrons to the product side of the reaction equation: \(H_2 \rightarrow 2H^+ + 2e^-\) (c) What is standard about the standard hydrogen electrode?

Define the standard hydrogen electrode

The standard hydrogen electrode is a reference electrode used in electrochemical measurements and has a standard reduction potential of 0 Volts by definition.

Describe the features of the standard hydrogen electrode

This electrode consists of a platinum wire coated with platinum black (a form of finely divided platinum) dipped into a hydrochloric acid solution of unit activity. It is exposed to a hydrogen gas of 1 atm pressure, and it is operated at a temperature of 298 K. The standard electrode potential for other electrodes is determined by comparing its potential against the standard hydrogen electrode.

Key concepts

Half-Reaction in Voltaic Cell

A voltaic cell creates electrical energy through chemical reactions, and understanding the half-reactions occurring at each electrode is fundamental to grasping its function. Imagine slicing the cell's reaction into two parts, where the reduction takes place at the cathode and oxidation at the anode. These ‘half-reactions’ collectively drive the cell's operation.

For instance, in the exercise, the half-reaction at the cathode for a hydrogen electrode in acidic solution is expressed as \(2H^+ + 2e^- \rightarrow H_2\). This signifies that two protons (hydrogen ions) are gaining electrons (being reduced) and forming hydrogen gas. Conversely, at the anode, the half-reaction \(H_2 \rightarrow 2H^+ + 2e^-\) indicates that hydrogen gas is losing electrons (being oxidized) and generating protons. The movement of electrons from the anode to the cathode through an external circuit is what generates electrical current.

Cathode and Anode Reactions

Distinguishing between cathode and anode reactions in a voltaic cell is crucial for understanding the direction of electron flow and the cell's overall energy output. At the cathode, reduction, a gain of electrons, transpires, and at the anode, oxidation, a loss of electrons, takes place.

To simplify, consider the reactions as two sides of a coin. Reduction at the cathode entails an input of electrons, akin to gaining wealth, while oxidation at the anode involves giving away electrons, similar to spending. In our exercise example, the cathode gains electrons to reduce hydrogen ions to hydrogen gas, and the anode releases electrons by converting hydrogen gas to hydrogen ions.

Standard Reduction Potential

The standard reduction potential is a quintessential concept in electrochemistry. It is a measure of the tendency of a chemical species to acquire electrons and thereby be reduced. Each half-reaction has an associated potential, and the standard hydrogen electrode, with its potential defined as 0 Volts, provides the baseline for these measurements.

Think of this as the 'zero' mark on a ruler from which other distances are measured. Just as objects can be taller or shorter compared to this mark, other half-reactions can have higher or lower reduction potentials compared to the standard hydrogen electrode. These potentials are crucial in predicting the direction of electron flow and the spontaneity of redox reactions within voltaic cells.

Electrochemistry

Electrochemistry deals with the interplay between electrical energy and chemical reactions. It's a fascinating field that lets us convert chemical energy into electrical energy, which powers countless devices in our lives.

In the context of our exercise, electrochemistry guides us through how the voltaic cell operates. The cell uses the redox reactions (oxidation at the anode and reduction at the cathode) to produce electric current. This field encompasses not only the design of batteries but also processes such as electroplating and the production of materials through electrolysis. Knowledge of electrochemistry is thus integral, not only in understanding how various cells operate but also in the development of technology and industrial processes.