Redox Reactions in Society (VCE SSCE Chemistry): Revision Notes

Redox Reactions in Society

Introduction

Redox reactions play a vital role in modern daily life. Many people rely heavily on mobile phones, laptop computers, calculators, and other portable electronic devices. These devices depend on redox reactions in rechargeable batteries to generate electricity. Without these reactions, we would need to keep our devices permanently plugged into mains electricity. This section explores how redox reactions are used to produce electrical energy in simple galvanic cells.

Simple galvanic cells

When a spontaneous reaction occurs between a metal and a non-metal, heat energy is released. The surrounding solution becomes warm, and sometimes even hot. This type of reaction is called an exothermic reaction, and it is typical of spontaneous redox reactions. Previously, we described redox reactions as electron transfer reactions, but we lacked concrete evidence of electrons moving from the reducing agent to the oxidising agent. Now we can see how the flow of electrons between reactants in a redox reaction can be used to operate electrical devices, providing clear evidence of electron movement.

Evidence for electron transfer

Consider what happens when zinc metal is placed in a copper(II) sulfate solution. The zinc is oxidised, and a brown deposit of copper metal forms on the surface:

Oxidation half-reaction: $\text{Zn}(s) \rightarrow \text{Zn}^{2+}(aq) + 2e^-$

Reduction half-reaction: $\text{Cu}^{2+}(aq) + 2e^- \rightarrow \text{Cu}(s)$

Overall equation: $\text{Zn}(s) + \text{Cu}^{2+}(aq) \rightarrow \text{Zn}^{2+}(aq) + \text{Cu}(s)$

Electrons flow from zinc atoms to copper(II) ions as these particles collide. In the direct reaction, heat is released. However, if we separate the reactants using specialized equipment, we can gain clear evidence for the flow of electrons.

In this setup, the zinc metal and the copper(II) ion solution are in separate beakers. This prevents a spontaneous, heat-releasing reaction from occurring. Instead, electrons are forced to travel through the external circuit (the wire) to reach the oxidising agent. A copper strip, dipping into the copper(II) ion solution, is connected to the zinc strip. A galvanometer—a device for detecting electron flow—is also included in the circuit.

The solutions in the two beakers are connected by a salt bridge. The salt bridge contains a solution of an ionic compound, such as potassium nitrate ($\text{KNO}_3$), which will not react with either solution. The salt bridge may be as simple as a piece of filter paper soaked in the potassium nitrate solution.

This apparatus is known as a galvanic cell, also called a voltaic cell, which is a type of electrochemical cell. A galvanic cell converts chemical energy into electrical energy and is one of the most important devices in modern life. It is a simple form of what we call a battery.

When the galvanic cell operates, the positive reading on the galvanometer indicates that electrons are flowing from the zinc strip to the copper strip. This observation provides evidence that the same oxidation and reduction reactions occur when the two reactants are separated as when they are mixed together in a beaker. It also supports the idea that redox reactions involve electron transfer.

Understanding half-cells

All galvanic cells consist of two half-cells. Each half-cell contains a reducing agent and its conjugate oxidising agent. Oxidation occurs in one half-cell, and reduction occurs in the other.

The two half-cells shown consist of metal electrodes in aqueous solutions of metal ions. For example:

• A silver electrode in a solution of silver ions, represented as $\text{Ag}(s)/\text{Ag}^+(aq)$
• A copper electrode in a solution of copper(II) ions, represented as $\text{Cu}(s)/\text{Cu}^{2+}(aq)$

Any combination of half-cells can be used to make a galvanic cell. When we join half-cells with a salt bridge and connect them with wire (the external circuit), we create a complete galvanic cell.

Oxidation and reduction in galvanic cells

Identifying the anode and cathode

Using the reactivity series, we can predict which reactions will occur in a galvanic cell. The stronger reducing agent will be oxidised. For the copper-silver cell shown above, copper is a stronger reducing agent than silver, so copper will be oxidised in the left-hand half-cell. The silver ions are a stronger oxidising agent than the copper(II) ions, so $\text{Ag}^+(aq)$ is reduced in the right-hand half-cell.

The half-cell reactions are:

$\text{Cu}(s) \rightarrow \text{Cu}^{2+}(aq) + 2e^-$

$\text{Ag}^+(aq) + e^- \rightarrow \text{Ag}(s)$

Because electrons leave the copper electrode to move into the external circuit, the copper electrode is known as the negative electrode—the source of electrons. Electrons then flow through the wire from the negative electrode to the positive electrode (the silver electrode). The silver metal does not gain electrons directly, but it conducts the electrons to the $\text{Ag}^+$ solution, where reduction of $\text{Ag}^+(aq)$ occurs.

Important terminology

The electrode at which oxidation occurs is called the anode, while the electrode at which reduction occurs is called the cathode. It is important to remember that the names anode and cathode relate to the type of reaction occurring at each electrode, rather than the charge on the electrode. However, in a galvanic cell, the anode is negative and the cathode is positive.

Constructing and analysing galvanic cells

When working with galvanic cells, you need to be able to identify several key features. Let's examine how to approach this using a zinc-lead galvanic cell as an example.

Historical development of batteries

The development of the battery is a fascinating story of scientific discovery and innovation that began with an unexpected observation about frogs.

Galvani's discoveries

Luigi Galvani (1737-1798), Professor of Anatomy at the University of Bologna in Italy, had a deep interest in frog anatomy and the effects of electricity on muscles. This led to the discovery of electrical flow between two different metals.

In 1791, when Galvani published his findings about electricity movement through the frog, he concluded that animal tissue contained a new form of electricity called "animal electricity", which activated nerves and muscles when metal probes were introduced.

Volta's contributions

Alessandro Volta (1745-1827), Professor of Physics at the University of Pavia in Italy, disagreed with Galvani's "animal electricity" idea. Instead, he proposed that the frog in Galvani's experiments simply conducted electricity between the two metals.

In 1792, Volta began experimenting with disks of different metals—without frogs! He detected the very weak flow of electricity between the disks by placing them on his tongue.

Volta received great recognition for his work from Napoleon I and the Austrian emperor Francis I. Most notably, the unit of electromotive force, the volt, was named in his honour in 1881.

Modern applications of galvanic cells

The portable power source essential for modern lifestyle—the battery—is based on a simple galvanic cell. While a battery is strictly defined as a series of cells connected together (like Volta's "crown of cups"), many portable power sources we refer to as batteries are actually individual cells. They consist of two electrodes and an electrolyte.

Primary cells

A cell that is disposable and designed not to be recharged is called a primary cell. The simplest primary cell that became widely available during the 20th century was known as a dry cell, based on a cell developed in 1866 by French scientist George Leclanche.

If you use a primary cell in a device such as a wireless computer mouse or a child's toy, you are more likely to be using an improvement on the dry cell called an alkaline cell. The difference between a dry cell and an alkaline cell relates to the electrolyte in the cell. While the electrolyte in a dry cell is acidic, the electrolyte in an alkaline cell is alkaline.

Secondary cells and renewable energy

An improvement on the dry cell is the ability to recharge these cells. A rechargeable cell designed to be reused many times is called a secondary cell. Many of the latest technological developments, such as electric vehicles and state-of-the-art power storage, are based on improvements in battery technology that have occurred in recent years. Most of this progress has focused on secondary cells.

Environmental importance of battery recycling

In July 2019, the Victorian government banned all e-waste from landfill, meaning that used batteries should not be thrown into general waste or household recycling bins.

Summary