Question

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

Short answer

(a) When an oxygen electrode serves as the cathode in a voltaic cell in acidic aqueous solution, the reduction half-reaction is: \(O2(g) + 4H^{+}(aq) + 4e^- \rightarrow 2H2O(l)\) (b) When an oxygen electrode serves as the anode in a voltaic cell in acidic aqueous solution, the oxidation half-reaction is: \(2H2O(l) \rightarrow O2(g) + 4H^{+}(aq) + 4e^-\) (c) The standard oxygen electrode, also called the standard hydrogen electrode, is a reference electrode used in electrochemistry under standard conditions (25°C, 1 atm for gases, and 1M concentration for other species). It has a reduction potential of 0 volts and is used to determine the standard electrode potential values of other electrodes in comparison.

Step-by-step solution

(a) Oxygen electrode at the cathode in an acidic solution

When an oxygen electrode serves as the cathode in a voltaic cell, it is the site of reduction half-reaction, meaning that oxygen will gain electrons. The reduction half-reaction in acidic solution is given by: O2(g) + 4H+(aq) + 4e- -> 2H2O(l)

(b) Oxygen electrode at the anode in an acidic solution

When an oxygen electrode serves as the anode in a voltaic cell, it is the site of oxidation half-reaction, meaning that water molecules will lose electrons, producing oxygen gas. The oxidation half-reaction in acidic solution is given by: 2H2O(l) -> O2(g) + 4H+(aq) + 4e-

(c) The standard oxygen electrode

The standard oxygen electrode (also called standard hydrogen electrode) is a reference electrode used in measuring the electrode potentials of half-cells (redox reactions), typically in electrochemistry. What is standard about this electrode is that it is set at standard conditions, like 25 degrees Celsius, 1 atm of pressure for gases, and 1M concentration for all other species. This electrode has a reduction potential of 0 volts. It is often used in comparison with other electrodes to determine their standard electrode potential values.

Key concepts

Understanding the Voltaic Cell

A voltaic cell, also known as a galvanic cell, is an essential device in electrochemistry that converts chemical energy into electrical energy. This cell relies on two separate reactions called redox reactions, where one substance is reduced and the other is oxidized.

In a voltaic cell, there are two electrodes: an anode and a cathode, each submersed in an electrolyte solution. At the anode, oxidation occurs, meaning it loses electrons. Conversely, at the cathode, reduction occurs, where it gains electrons. A wire connecting the electrodes allows the flow of electrons, producing electric current.

Electrons move from the anode to the cathode through the wire, often powering devices or charging batteries. The characteristics of a voltaic cell allow it to be used in many fields, from powering electronic devices to being a fundamental part of numerous industrial processes. Understanding voltaic cells provides an excellent insight into how chemical reactions can be harnessed for practical energy applications.

Exploring Half-Reactions in Electrochemistry

Half-reactions are fundamental in explaining the processes at each electrode of a voltaic cell. They depict the gain or loss of electrons, breaking down complex redox reactions into simpler parts.

In our case, focusing on an oxygen electrode in an acidic solution, we observe two distinct half-reactions. At the cathode, the reduction half-reaction involves oxygen molecules gaining electrons:

• \(O_2(g) + 4H^+(aq) + 4e^- \rightarrow 2H_2O(l)\)

This reaction highlights how oxygen, water, and hydrogen ions are interconnected in reduction.

Conversely, at the anode, oxidation involves water molecules losing electrons:

• \(2H_2O(l) \rightarrow O_2(g) + 4H^+(aq) + 4e^-\)

Here, water splits into oxygen gas and protons, releasing electrons. By individually examining these half-reactions, it's easier to understand the comprehensive redox reaction and predict the behavior of substances in electrochemical cells.

The Role of the Standard Oxygen Electrode

In the realm of electrochemistry, the standard oxygen electrode plays a vital role as a reference point. It provides a benchmark for measuring and comparing electrode potentials of other half-cells. This electrode, often mistaken for the hydrogen electrode due to closely related functions, has defined conditions.

Standard conditions for this electrode generally include a temperature of 25°C, a pressure of 1 atm for gases, and concentrations of 1 molar for all involved species. Under these conditions, the standard oxygen electrode is set at a reduction potential of 0 volts.

This standard potential allows the electrode to serve as a universal reference, greatly simplifying the process of determining the standard electrode potential of other half-reactions. Standard conditions ensure consistency and reliability in electrochemical measurements, making the standard oxygen electrode indispensable for scientists and engineers working with electrochemical cells.