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) What is standard about the standard hydrogen electrode? (c) What is the role of the platinum foil in a 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: \[ 2H^+_{(aq)} + 2e^- \rightarrow H_{2(g)} \] (b) The Standard Hydrogen Electrode (SHE) is standard due to its stable and reproducible conditions, such as \(1 M\) hydrogen ion concentration, \(1 atm\) hydrogen gas pressure, and a constant temperature of \(25^\circ C\). Under these conditions, the reduction potential of the SHE is defined as 0 Volts. (c) The platinum foil in a standard hydrogen electrode serves as an inert and highly conductive material, providing a catalytic surface for the reduction of hydrogen ions and oxidation of hydrogen gas while ensuring accurate measurements due to its high surface area.

Step-by-step solution

Writing the half-reaction at the hydrogen electrode as a cathode

To write the half-reaction, let's first understand that the cathode is the electrode where reduction (gain of electrons) occurs. In an acidic aqueous solution, hydrogen ions (protons) are reduced to form hydrogen gas. The half-reaction for this process is: \[ 2H^+_{(aq)} + 2e^- \rightarrow H_{2(g)} \]

Explaining what is standard about the standard hydrogen electrode (SHE)

The Standard Hydrogen Electrode (SHE) is the reference electrode against which the reduction potential of all other electrodes is measured. It is considered standard because the half-cell has stable and reproducible conditions. These conditions include: 1. The hydrogen ion concentration in the solution is \(1 M\), which corresponds to a pH of 0. 2. The pressure of hydrogen gas is maintained at \(1 atm\). 3. The temperature is kept constant at \(25^\circ C\). Under these conditions, the reduction potential of the SHE is defined as 0 Volts.

Discussing the role of the platinum foil in a standard hydrogen electrode

The platinum foil plays an essential role in the standard hydrogen electrode due to its unique properties. The main functions of the platinum foil are: 1. Inertness: Platinum is a noble metal, which means it does not readily participate in any chemical reactions. This ensures that the platinum foil itself will not interfere with the half-reaction occurring at the electrode. 2. High conductivity: Platinum has excellent electrical conductivity, which allows electrons to flow freely between the electrolyte solution and the external circuit. 3. Catalytic activity: Platinum helps to catalyze the reactions occurring on its surface. In the case of the standard hydrogen electrode, it facilitates the reduction of hydrogen ions to form hydrogen gas, as well as the oxidation of hydrogen gas to form hydrogen ions. 4. High surface area: The rough surface of the platinum foil creates a large surface area for both the hydrogen ions in the solution and the hydrogen gas to adsorb, which enhances the electrode reaction rate. This ensures more accurate measurements of reduction potentials.

Key concepts

Electrochemical Cells

Electrochemical cells are the cornerstones of various important technologies, including batteries and sensors. These systems convert chemical energy into electrical energy through redox (reduction-oxidation) reactions. At the heart of this process are two electrodes: the anode, where oxidation occurs, and the cathode, where reduction takes place.

During operation, electrons flow from the anode to the cathode through an external circuit, while ions move through an internal path to maintain charge balance. This electron flow is harnessed as electric current, which powers devices or performs electrical work. Understanding the fundamentals of electrochemical cell operation helps in comprehending various applications, from galvanic cells, which discharge energy, to electrolytic cells, which use electrical energy to drive chemical reactions.

Reduction Potential

Reduction potential, also known as redox potential, is a measure of the tendency of a chemical species to acquire electrons and thereby be reduced. It is a crucial concept in electrochemistry as it determines the direction of electron flow in an electrochemical cell.

Reduction potentials are measured in volts and are relative to the standard hydrogen electrode (SHE)—a universal reference with a reduction potential set to 0 volts. By comparing the reduction potentials of different half-reactions, scientists can predict which substances will act as oxidizing or reducing agents and the feasibility of redox reactions. These comparisons form the basis of the electrochemical series, a list that ranks half-reactions by their standard reduction potentials.

Role of Platinum Electrode

Platinum electrodes are used for their remarkable durability and ability to conduct electrical current without corroding or reacting with the electrochemical cell's contents.

In the context of the standard hydrogen electrode, the platinum electrode serves multiple purposes: it facilitates electron transfer, acts as a catalytic surface for the adsorption and reaction of hydrogen ions and molecules, and maintains the integrity of the reaction without participating chemically. This characteristic is crucial for consistent and accurate measurements of reduction potentials across various half-reactions in electrochemistry. Platinum's properties make it an ideal material for sustaining the SHE's conditions and ensuring reproducible results.

Half-Reactions in Electrochemistry

Electrochemistry revolves around the concept of half-reactions, which are the individual oxidation or reduction processes that occur at the electrodes. These half-reactions are essential for understanding the overall redox reaction in an electrochemical cell.

In a galvanic cell, the oxidation half-reaction releases electrons at the anode, while the reduction half-reaction consumes electrons at the cathode. Writing and balancing these half-reactions make it easier to grasp the electron transfer process and allow for the calculation of the cell's overall electromotive force (EMF). Recognizing the components of a half-reaction, such as the species being oxidized or reduced, the exchange of electrons, and the potential changes, aids students in piecing together the bigger picture of electrochemical processes.