<img src="https://d5nxst8fruw4z.cloudfront.net/atrk.gif?account=iA1Pi1a8Dy00ym" style="display:none" height="1" width="1" alt="" />
Skip Navigation

23.3: Electrolysis

Difficulty Level: At Grade Created by: CK-12
Turn In

Lesson Objectives

  • Distinguish between voltaic and electrolytic cells.
  • Describe a Down’s cell and identify the products of the electrolysis of molten sodium chloride.
  • Describe the reactions that occur during the electrolysis of water.
  • Identify the products that would be generated during the electrolysis of an aqueous solution of sodium chloride.
  • Describe the process of electroplating.

Lesson Vocabulary

  • electrolysis
  • electrolytic cell
  • electroplating

Check Your Understanding

Recalling Prior Knowledge

  • What half-reactions occur at the anode and cathode of an electrochemical cell?
  • How can you tell from a standard cell potential whether an electrochemical reaction is spontaneous or nonspontaneous under standard conditions?

A nonspontaneous reaction is one in which the reactants are favored over the products under a given set of reaction conditions. However, if a chemical system is supplied with energy from an external source, it is possible to drive a reaction in the nonspontaneous direction. In this lesson, you will learn about this type of electrochemical process, which is called electrolysis.

Electrolytic Cells

A voltaic cell uses a spontaneous redox reaction to generate an electric current. It is also possible to do the opposite. When an external source of direct current is applied to an electrochemical cell, a reaction that is normally nonspontaneous can be forced to proceed. Electrolysis is the process in which electrical energy is used to cause a nonspontaneous chemical reaction to occur. Electrolysis is responsible for the metal coatings that appear on many everyday objects, such as gold-plated or silver-plated jewelry and chrome-plated car bumpers.

An electrolytic cell is the apparatus used for carrying out an electrolysis reaction. Figure below shows an electrolytic cell composed of Zn|Zn2+ and Cu|Cu2+ half-cells.

An electrolytic cell uses an external power source (a battery) to drive a nonspontaneous reaction. The copper half-cell undergoes oxidation, while the zinc half-cell undergoes reduction.

Recall that in the last section, this same pair of half-cells was used as an example of a voltaic cell. In the spontaneous direction, Zn metal is oxidized to Zn2+ ions while Cu2+ ions are reduced to Cu metal. In a voltaic cell, the zinc electrode would be the anode and the copper electrode would be the cathode. However, when the same half-cells are connected to a battery via an external wire, the reaction is forced to run in the opposite direction. The zinc electrode is now the cathode and the copper electrode is the anode.

\begin{align*} &\text{Oxidation (anode):} && \text{Cu}(s) \rightarrow \text{Cu}^{2+}(aq) + 2\text{e}^- &&\text{E}^0 = -0.34 \ \text{V} \\ &\text{Reduction (cathode):} && \text{Zn}^{2+}(aq) + 2\text{e}^- \rightarrow \text{Zn}(s) && \text{E}^0 = -0.76 \ \text{V} \\ \hline &\text{Overall reaction:} && \text{Cu}(s) + \text{Zn}^{2+}(aq) \rightarrow \text{Cu}^{2+}(aq) + \text{Zn}(s) &&\text{E}^0_{\text{cell}} = -1.10 \ \text{V} \end{align*}

The standard cell potential is negative, indicating a nonspontaneous reaction. The battery must be capable of delivering at least 1.10 V of direct current in order for the reaction to occur. Another difference between a voltaic cell and an electrolytic cell is the signs that are commonly given to the electrodes. In a voltaic cell, the anode is negative and the cathode is positive. In an electrolytic cell, the anode is positive because it is connected to the positive terminal of the battery. The cathode is negative. Electrons still flow through the cell from the anode to the cathode.

Examples of Electrolysis Reactions

Several electrolysis reactions are commonly performed on a large scale for the commercial production of certain substances. In this section, we will examine three examples of electrolysis.

Electrolysis of Molten Sodium Chloride

Molten (liquid) sodium chloride can be electrolyzed to produce sodium metal and chlorine gas. The electrolytic cell used in this process is called a Down’s cell (Figure below).

A Down’s cell is used for the electrolysis of molten sodium chloride. Liquid sodium metal is produced at the cathode, while chlorine gas is produced at the anode.

In a Down’s cell, the liquid sodium ions are reduced at the cathode to liquid sodium metal. At the anode, liquid chloride ions are oxidized to chlorine gas. The reactions and cell potentials are shown below.

\begin{align*} &\text{Oxidation (anode):} && 2\text{Cl}^-(l) \rightarrow \text{Cl}_2(g) + 2\text{e}^- &&\text{E}^0 = -1.36 \ \text{V} \\ &\text{Reduction (cathode):} && \text{Na}^+(l) + \text{e}^- \rightarrow \text{Na}(l) && \text{E}^0 = -2.71 \ \text{V} \\ \hline &\text{Overall reaction:} && 2\text{Na}^+(l) + 2\text{Cl}^-(l) \rightarrow 2\text{Na}(l) + \text{Cl}_2(g) &&\text{E}^0_{\text{cell}} = -4.07 \ \text{V} \end{align*}

The battery must supply over 4 volts to carry out this electrolysis. This reaction is a major industrial source of chlorine gas, and it is the primary way to obtain pure sodium metal. Chlorine gas is widely used as a disinfectant, such as in swimming pools.

Electrolysis of Water

The electrolysis of water produces hydrogen and oxygen gases. The electrolytic cell consists of a pair of platinum electrodes immersed in water containing a small amount of an electrolyte, such as H2SO4. The electrolyte is necessary because pure water does not contain enough ions to effectively conduct a current. At the anode, water is oxidized to oxygen gas and hydrogen ions. At the cathode, water is reduced to hydrogen gas and hydroxide ions.

\begin{align*} &\text{Oxidation (anode):} && 2\text{H}_2\text{O}(l) \rightarrow \text{O}_2(g) + 4\text{H}^+(aq) + 4\text{e}^- &&\text{E}^0 = -1.23 \ \text{V} \\ &\text{Reduction (cathode):} && 2\text{H}_2\text{O}(l) + 2\text{e}^- \rightarrow \text{H}_2(g) + 2\text{OH}^-(aq) && \text{E}^0 = -0.83 \ \text{V} \\ \hline &\text{Overall reaction:} && 2\text{H}_2\text{O}(l) \rightarrow \text{O}_2(g) + 2\text{H}_2(g) &&\text{E}^0_{\text{cell}} = -2.06 \ \text{V} \end{align*}

In order to obtain the overall reaction, the reduction half-reaction was multiplied by two to equalize the electrons. The hydrogen ion and hydroxide ions produced in each reaction combine to form water. The added electrolyte is not consumed in the reaction.

Apparatus for the production of hydrogen and oxygen gases by the electrolysis of water.

Electrolysis of Aqueous Sodium Chloride

Earlier we examined the electrolysis of molten sodium chloride. It may be logical to assume that the electrolysis of aqueous sodium chloride, called brine, would yield the same result by the same reactions. However, the reduction reaction that occurs at the cathode does not produce sodium metal, because the water is reduced instead. This is because the reduction potential for water is only −0.83 V compared to −2.71 V for the reduction of sodium ions. This makes the reduction of water preferable, because its reduction potential is less negative. Chlorine gas is still produced at the anode, just as in the electrolysis of molten NaCl.

\begin{align*} &\text{Oxidation (anode):} && 2\text{Cl}^-(aq) \rightarrow \text{Cl}_2(g) + 2\text{e}^- &&\text{E}^0 = -1.36 \ \text{V} \\ &\text{Reduction (cathode):} && 2\text{H}_2\text{O}(l) + 2\text{e}^- \rightarrow \text{H}_2(g) + 2\text{OH}^-(aq) && \text{E}^0 = -0.83 \ \text{V} \\ \hline &\text{Overall reaction:} && 2\text{Cl}^-(aq) \rightarrow \text{O}_2(g) + 2\text{H}_2(g) &&\text{E}^0_{\text{cell}} = -2.06 \ \text{V} \end{align*}

Since the hydroxide ion is also a product of the net reaction, the important chemical sodium hydroxide (NaOH) is obtained by evaporating the water after the hydrolysis is complete.


Many decorative objects like jewelry are manufactured with the aid of an electrolytic process. Electroplating is a process in which a metal ion is reduced in an electrolytic cell to deposit the solid metal onto a surface. Figure below shows a cell in which silver metal is to be plated onto a stainless steel spoon.

An electrolytic cell used in the electroplating of silver onto a metal spoon. A silver strip is the anode, while the spoon itself is the cathode.

The cell consists of a solution of silver nitrate and a strip of silver, which acts as the anode. The spoon is the cathode. The anode is connected to the positive electrode of a battery, while the spoon is connected to the negative electrode.

When the circuit is closed, silver metal from the anode is oxidized, allowing silver ions to enter the solution.

\begin{align*}\text{Anode:} \;\;\;\;\;\; \text{Ag}(s) \rightarrow \text{Ag}^+(aq) + \text{e}^-\end{align*}\

Meanwhile, silver ions from the solution are reduced to silver metal on the surface of the cathode, the steel spoon.

\begin{align*}\text{Cathode:} \;\;\;\;\; \text{Ag}^+(aq) + \text{e}^- \rightarrow \text{Ag}(s)\end{align*}

The concentration of silver ions in the solution is effectively constant. The electroplating process transfers metal from the anode to the cathode of the cell. Other metals commonly plated onto objects include chromium, gold, copper, and platinum.

Lesson Summary

  • Electrolysis is a process in which a nonspontaneous redox reaction is driven forward by an external power source, such as a battery. The voltage of the battery must be at least as great as the negative cell potential.
  • Molten sodium chloride can be electrolyzed in a Down’s cell to yield sodium metal and chlorine gas.
  • The electrolysis of water produces hydrogen and oxygen gases.
  • When a concentrated aqueous solution of sodium chloride (brine) is electrolyzed, chlorine and hydrogen gases are produced.
  • Electroplating is a process by which a solution of metal ions is plated out as a neutral metal surface onto the object used as the cathode.

Lesson Review Questions

Reviewing Concepts

  1. Distinguish between a voltaic cell and an electrolytic cell in terms of the nature of the redox reaction.
  2. What sign is assigned to the cathode in a voltaic cell? In an electrolytic cell?
  3. What is produced at the anode of a Down’s cell? At the cathode?
  4. Why can’t sodium metal be manufactured by the electrolysis of brine?
  5. Write the half-reaction that occurs at the cathode during the electroplating of chromium metal from a solution of chromium(III) nitrate.


  1. Aluminum metal is obtained by a process that involves the electrolysis of molten aluminum oxide (Al2O3). Write the half-reaction that occurs at the cathode.
  2. Consider the electrolysis of molten potassium bromide.
    1. Write the equations for the half-reactions that occur at each electrode.
    2. Write the overall redox reaction.
    3. Use the Standard Reduction Potentials at 25°C table (in the lesson “Cell Potentials”) to calculate the minimum voltage of the battery used to electrolyze potassium bromide.
  3. Consider the electrolysis of water.
    1. Write the equation for the overall reaction.
    2. In a certain electrolysis experiment, 2.20 L of oxygen gas is produced. What volume of hydrogen gas is produced?
  4. During the electrolysis of molten sodium chloride in a Down’s cell, 356 L of chlorine gas is produced at a temperature of 850°C and a pressure of 1.00 atm.
    1. How many moles of chlorine gas are produced? (Hint: Use ideal gas law with R = 0.0821 L•atm/K•mol)
    2. How many moles of sodium metal are produced in the same process?
    3. What mass of sodium is produced?
  5. An electrolytic cell is constructed with a Pb|Pb2+ anode and a Cd|Cd2+ cathode.
    1. Write the half reactions for each electrode and the overall reaction.
    2. Calculate the standard cell potential.
    3. The cell is run for a certain amount of time in which the mass of the lead electrode decreases by 1.00 g. By how much does the mass of the cadmium electrode increase in the same time?

Further Reading / Supplemental Links

Points to Consider

Aluminum is widely used in all sorts of modern materials from beverage cans to airplanes.

  • Why was pure aluminum metal so difficult and expensive to obtain prior to the development of electrolysis on an industrial scale?
  • How much energy savings comes from recycling aluminum instead of producing it from aluminum ore?

Notes/Highlights Having trouble? Report an issue.

Color Highlighted Text Notes
Please to create your own Highlights / Notes
Show More

Image Attributions

Show Hide Details
Files can only be attached to the latest version of section
Please wait...
Please wait...
Image Detail
Sizes: Medium | Original