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# 23.4: Electrolysis

Difficulty Level: At Grade Created by: CK-12

## Lesson Objectives

The student will:

• identify the anode and the cathode given a diagram of an electrolysis apparatus that includes the compound being electrolyzed.
• write oxidation and reduction half-reactions given a diagram of an electrolysis apparatus that includes the compound being electrolyzed.

## Vocabulary

• anode
• cathode
• electric current
• electrolysis
• electroplating
• ionic conduction
• metallic conduction

## Introduction

Electrolysis involves using an electric current to force an otherwise non-spontaneous chemical reaction to occur.

## Electric Current

Any flow of electric charge is an electric current. Electrons are far easier to remove from atoms than protons, so most of the electric current you experience in daily life is electron flow. Electrons move easily through a piece of metal because of the freedom of movement of valence electrons in the metallic bond. The movement of electrons through a piece of metal is called metallic conduction. An electric current also exists when positive or negative ions move along a path. The movement of ions through a solution is called ionic conduction.

Electric current of all types flow because there is a difference in electric potential energy at two positions. A common device for providing this difference in electric potential energy is a battery. A battery has two terminals, and chemical reactions inside the battery cause one terminal to have higher electric potential energy than the other terminal. The higher potential energy terminal is called the negative terminal, and the other terminal is called the positive terminal. If a low resistance path is provided, electrons will flow along the path from high to low potential energy. A metal wire provides such a low resistance path for electrons. The potential energy lost by the electrons as they move from high electric potential energy to low electric potential energy is converted into heat and light by the light bulb filament, as illustrated below.

## Current Through an Electrolyte

If we connect pieces of metal to the terminals of a battery with wires and suspend the pieces of metal (called electrodes) in an ionic solution as shown below, the high potential energy electrons will flow onto the negative electrode. Simultaneously, some electrons will flow off the other electrode, leaving it positive.

The cations (A+)\begin{align*}(A^+)\end{align*} in the solution are attracted to the negative electrode and will migrate through the solution to contact the negative electrode. When a +1\begin{align*}+1\end{align*} cation contacts the negative electrode, it will pick up an electron from the electrode and become a neutral atom. In this way, electrons leave the negative electrode. You should recognize that the cations that have picked up an electron are reduced.

At the same time the positive ions are migrating toward the negative electrode, the anions (B)\begin{align*}(B^-)\end{align*} are migrating toward the positive electrode. When a negative ion touches the positive electrode, it gives up its extra electron to the electrode and becomes a neutral atom. The anions that donate electrons to the positive electrode are oxidized. In this way, electrons are added to the positive electrode. Even though the electrons that leave the negative terminal of the battery and the ones that arrive at the positive terminal are not the same, the circuit is nevertheless complete. As illustrated below, current flows through the circuit just as if a wire had been placed between the two electrodes. The processes that occur at the two electrodes are simultaneous.

You should also realize that the ions in the solution are being consumed by this process, and the current will only flow as long as there are ions to accept and donate electrons. Once the solution runs out of ions, current flow stops. When a current flows through an electrolyte in this manner, the electrolytic solution is called the internal part of the circuit, and the wires and battery are called the external part of the circuit.

## Electrolysis of Liquid NaCl

Solid sodium chloride consists of alternating sodium ions and chloride ions in a tightly packed, three-dimensional, crystal lattice (as seen in the image below). Ionic solids do not conduct electric current because the ions cannot migrate through the solid material. If sodium chloride is heated to a high enough temperature, however, it melts to a liquid. In the liquid form, ions can migrate. Therefore, ionic substances in liquid form will conduct electricity in approximately the same way as ions in solution.

If electrodes connected to battery terminals are placed in liquid sodium chloride, the sodium ions will migrate toward the negative electrode and be reduced while the chloride ions migrate toward the positive electrode and are oxidized. The processes that occur at the electrodes can be represented by what are called half-equations.

At the positive electrode: Na++eNa\begin{align*}\mathrm{Na}^+ + e^- \rightarrow \mathrm{Na}\end{align*} (reduction)

At the negative electrode: 2 ClCl2+2 e\begin{align*}2 \ \mathrm{Cl}^- \rightarrow \mathrm{Cl}_2 + 2 \ e^-\end{align*} (oxidation)

Half-equations are very helpful in discussing and analyzing processes, but half-reactions cannot occur as they appear. Both oxidation and reduction must occur at the same time, so the electrons are donated and absorbed nearly simultaneously.

The two half-reactions may be added together to represent a complete reaction. In order to add the half-reactions, the number of electrons donated and the number of electrons accepted must be equal. In the case of the oxidation of the chloride ion and the reduction of the sodium ion, the sodium half-reaction must be doubled so that each half-reaction involves two electrons.

2 (Na++eNa)=2 Na++2 e2 Cl2 Na++2 e+2 Cl2 NaCl2+2 e2 Na+Cl2+2 e\begin{align*} \begin{array}{rll} 2 \ (\mathrm{Na}^+ + e^- \rightarrow \mathrm{Na}) = 2 \ \mathrm{Na}^+ + 2 \ e^- & \rightarrow & 2 \ \mathrm{Na}\\ 2 \ \mathrm{Cl}^- & \rightarrow & \mathrm{Cl}_2 + 2 \ e^-\\ \hline 2 \ \mathrm{Na}^+ + 2 \ e^- + 2 \ \mathrm{Cl}^- & \rightarrow & 2 \ \mathrm{Na} + \mathrm{Cl}_2 + 2 \ e^- \end{array} \end{align*}

The two electrons on each side of the equation can be canceled, and we are left with the net reaction for the electrolysis of sodium chloride.

2 Na++2 Cl2 Na+Cl2\begin{align*}2 \ \mathrm{Na}^+ + 2 \ \mathrm{Cl}^- \rightarrow 2 \ \mathrm{Na} + \mathrm{Cl}_2\end{align*}

Many chemical reactions occur spontaneously because the products have less potential energy that the reactants. In electrolysis reactions, however, the products have more potential energy than the reactants. Electrolysis reactions will not run unless energy is put into the system from outside. In the case of electrolysis reactions, the energy is provided by the battery.

## Electrolysis of Water

If we place electrodes connected to the terminals of a battery into pure water, no current is conducted and no electrolysis occurs. This is because ions must be present in order for a current to be conducted. Even though water is very slightly ionized to H+\begin{align*}\mathrm{H}^+\end{align*} and OH\begin{align*}\mathrm{OH}^-\end{align*} ions, the concentrations of these ions is too small to produce electrolysis. When a small amount of ion-supplying substance, such as H2SO4\begin{align*}\mathrm{H}_2\mathrm{SO}_4\end{align*} or Na2SO4\begin{align*}\mathrm{Na}_2\mathrm{SO}_4\end{align*}, is added to the water in the electrolysis set-up, hydrogen gas is rapidly produced at the negative electrode and oxygen gas is produced at the positive electrode. We will now call the negative electrode the cathode and the positive terminal the anode. The cathode is the electrode where reduction occurs, and the anode is the electrode where oxidation occurs. Cations were so named because they are attracted to the cathode, and anions were so named because they are attracted to the anode.

In the electrolysis of water, oxidation and reduction half-reactions are:

half-reaction at the cathode:half reaction at the anode:and the net reaction is:4 H2O+4 e2 H2O6 H2O2 H2(g)+4 OHO2(g)+4 H++4 e2 H2(g)+O2(g)+4 H++4 OH\begin{align*} \begin{array}{rrll} \text{half-reaction at the cathode:} & 4 \ \mathrm{H}_2\mathrm{O} + 4 \ e^- & \rightarrow & 2 \ \mathrm{H}_{2(g)} + 4 \ \mathrm{OH}^-\\ \text{half reaction at the anode:} & 2 \ \mathrm{H}_2\mathrm{O} & \rightarrow & \mathrm{O}_{2(g)} + 4 \ \mathrm{H}^+ + 4 \ e^-\\ \hline \text{and the net reaction is:} & 6 \ \mathrm{H}_2\mathrm{O} & \rightarrow & 2 \ \mathrm{H}_{2(g)} + \mathrm{O}_{2(g)} + 4 \ \mathrm{H}^+ + 4 \ \mathrm{OH}^- \end{array} \end{align*}

(This is the net reaction if the two half-reactions occur in separate chambers. If the two half-reactions occur in the same chamber, the proton and hydroxide ions react to form water.)

## Electroplating

With appropriate treatment of the electrode material and appropriate adjustment of the current level from the battery, it is possible to get the metal being reduced to adhere strongly to the electrode during the electrolysis process. The use of electrolysis to coat one material with a layer of metal is called electroplating. Usually, electroplating is used to cover a cheap metal with a layer of more expensive and more attractive metal. Sometimes, electroplating is used to get a surface metal that is a better conductor of electricity. When you wish to have the surface properties of gold (attractive, corrosion resistant, or good conductor), but you don’t want to have the great cost of making the entire object out of solid gold, the answer may be to use cheap metal to make the object and then electroplate a thin layer of gold on the surface.

To silver plate an object like a spoon, the spoon is placed in the position of the cathode in an electrolysis set up with a solution of silver nitrate. When the current is turned on, the silver ions will migrate through the solution, touch the cathode (spoon), and adhere to it. With enough time and care, a layer of silver can be plated over the entire spoon. The anode for this operation would often be a large piece of silver from which silver ions would be oxidized, and these ions would enter the solution. This is a way of ensuring a steady supply of silver ions for the plating process.

Half-reaction at the cathode: Ag++eAg\begin{align*}\mathrm{Ag}^+ + e^- \rightarrow \mathrm{Ag}\end{align*}

Half-reaction at the anode: AgAg++e\begin{align*}\mathrm{Ag} \rightarrow \mathrm{Ag}^+ + e^-\end{align*}

Some percentage of the gold and silver jewelry sold is electroplated. The connection points in electric switches are often gold plated to improve electrical conductivity, and most of the chromium pieces on automobiles are chromium plated.

## Lesson Summary

• An electric current consists of a flow of charged particles.
• When direct current is passed through a solution of an electrolyte, cations are attracted to the negative electrode where they gain electrons, and anions are attracted to the positive electrode where they lose electrons.
• The electrode where oxidation occurs is called the anode, and the electrode where reduction occurs is called the cathode.
• In electroplating, the object to be plated is made the cathode in an electrolysis.

This website provides an interactive animation of electrolysis.

## Review Questions

1. Write the equations for the reactions that occur at the anode and at the cathode in the electrolysis of molten KBr.

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