Lesson Explainer: Galvanic Cells Chemistry

In this explainer, we will learn how to describe a galvanic cell and explain how 𝐸 values are measured.

When metal is placed in a solution of its ions, an equilibrium forms between the metal and its ions. For example, when a piece of copper metal is placed in a copper(II) nitrate solution, an equilibrium is established between the atoms of solid copper metal and the Cu2+ ions in solution.

The equation for the equilibrium half-reaction is Cu()Cu()+2esaq2+

Imagine a second beaker containing zinc metal in equilibrium with zinc ions in solution. The equation for this equilibrium half-reaction is Zn()Zn()+2esaq2+

We could consider these two pieces of metal to be electrodes that are used to transfer charge to and from the solutions. Electrodes are often made of metal but sometimes of a nonmetal, for example, graphite. The electrodes can also transfer electrons to and from a circuit; however, in the above diagrams, there are no complete circuits. We will discuss later how electrons flow to and from electrodes in a circuit.

Definition: Electrode

An electrode is a conductor used to transfer electrical current to and from a nonmetallic part of a circuit.

The two beakers in the diagrams above can be called half-cells. A half-cell consists of an electrode in a solution of ions.

Definition: Half-Cell

A half-cell is a system composed of a conductive electrode in an electrolyte solution.

The solutions are called electrolytes because they can conduct electrical charge due to the presence of freely moving ions.

Example 1: Identifying the Diagram That Correctly Represents an Equilibrium between a Metal and Its Ions in Solution

Which of the following diagrams depicts the equilibrium that exists between a piece of zinc metal in contact with an aqueous solution of zinc ions?


Zinc atoms on the solid piece of zinc are released into the solution as Zn2+ ions, leaving behind their two valence electrons on the solid zinc metal. At the same time, Zn2+ ions in solution are “replated” onto the solid zinc metal by accepting two electrons from the solid metal. Both the forward and reverse reactions occur simultaneously: Zn()Zn()+2esaq2+.

So, C is the correct answer.

We can connect the two metal electrodes, copper and zinc, using wires and a high-resistance voltmeter. Next, we can then complete the circuit, by joining the two half-cell solutions using a salt bridge. The diagram shows the setup of a completed circuit.

The salt bridge connecting the two solutions is usually made of a glass U-tube. The U-tube is filled with an ionic solution, for example, potassium nitrate. These ions are able to move, completing the circuit.

Definition: Salt Bridge

It is a tube containing an electrolyte connecting two half-cell solutions and providing electrical contact, thereby completing the circuit.

The voltmeter will register the maximum voltage possible, also known as the electromotive force, between the two pieces of metal. This is because copper and zinc have a different potential to form ions in a solution, in other words, a different potential to become oxidized. Zinc is more easily oxidized than copper. The converse is also true—copper is more easily reduced than zinc. The high-resistance voltmeter prevents current from flowing, because of its high resistance, but it will measure the potential difference between the two different electrodes.

It is this potential difference that is used to drive reactions and generate electricity in batteries. If the voltmeter is replaced with another circuit component, for example, a light bulb, the two half-cells will act as a cell; multiple cells connected together is what we understand as a battery. A spontaneous redox reaction will occur in the half-cells. The reaction converts chemical energy to electrical energy and causes current to flow in the wire, and the light bulb glows/burns. This is an example of an electrochemical cell.

Definition: Electrochemical Cell

Electrochemical cell is a system or a device that converts chemical energy to electrical energy or electrical energy to chemical energy.

An electrochemical cell that undergoes a spontaneous chemical reaction is called a galvanic cell or a voltaic cell. In a galvanic cell, a redox reaction converts chemical energy to electrical energy.

Definition: Galvanic Cell

It is a type of electrochemical cell where electrons are generated spontaneously through a redox reaction; these electrons pass through an external circuit. Another name for a galvanic cell is a voltaic cell.

Galvanic cells convert chemicalenergyelectricalenergy

There is another type of electrochemical cell called an electrolytic cell. In an electrolytic cell, electrical energy needs to be supplied to drive a nonspontaneous chemical reaction.

Definition: Electrolytic Cell

It is an electrochemical cell that uses electrical energy to drive a nonspontaneous chemical reaction.

Electrolytic cells convert electricalenergychemicalenergy

Electrolytic cells are discussed in more depth in another explainer.

Let us look in depth at how our zinc–copper galvanic cell connected to a bulb generates electrical current in the wire. Remember the potential difference between the electrodes is the driving force, and because zinc is more easily oxidized than copper, zinc atoms in the anode will donate their valence electrons into the conducting wire. So, oxidation occurs at the zinc electrode. We call this electrode the anode. These electrons will travel to the copper cathode. Zn2+ ions at the anode will be released into the solution, increasing the concentration of Zn2+ ions in the zinc half-cell. Electrons entering the cathode will attract Cu2+ ions in solution in the copper half-cell. So, reduction occurs at the copper electrode, and we call this electrode the cathode. Cu2+ ions will gain two electrons each and will be plated or deposited onto the cathode. In this way, negative charge flows through the wire from the Zn half-cell to the Cu half-cell.

Definition: Anode

The anode is the electrode at which oxidation occurs.

Definition: Cathode

The cathode is the electrode at which reduction occurs.

Let us rewrite both half-reactions showing both reduction and oxidation: Cu()+2eCu()(thereductionhalf-reaction)Zn()Zn()+2e(theoxidationhalf-reaction)2+2+aqssaq

The overall reaction for this electrochemical cell is Zn()+Cu()Zn()+Cu()(theoverallredoxequation)saqaqs2+2+

Negative ions in solution (in this case, NO3 ions) complete this flow by moving through the salt bridge. It is important to use the correct solution in the salt bridge. Potassium ions and nitrate ions in the salt bridge do not interact with or precipitate out of solution.

The current will continue to flow as long as there is a reaction occurring. The reaction will stop and the current will cease to flow when all the zinc metal has been converted to ions in solution, or when there are no more copper ions in solution because they have all been deposited on the cathode.

It is time-consuming and inconvenient to draw a large diagram of a full electrochemical cell. A more convenient way to represent information about an electrochemical cell is by using cell notation, sometimes called a cell diagram. Here is the cell notation for the zinc–copper galvanic cell:

Information about the anode is written on the left-hand side. We begin by writing the symbol of the metal electrode followed by a vertical line. This vertical line represents a phase boundary. Next, we write the symbol and charge of the corresponding ion in solution in the anode half-cell. Sometimes, the concentration of the solution is given in parentheses. The double vertical line represents the salt bridge connecting the two half-cells. Information about the cathode is then written on the right-hand side, namely, the metal ion in solution followed by a vertical line for the phase boundary, and then the cathode element symbol. Spectator ions are not included in the cell notation. The solution concentrations are specified, and 1 M is the standard concentration.

Example 2: Writing Cell Notation

What is the correct cell notation for a half-cell consisting of a piece of silver metal placed into a solution of silver ions?

  1. Ag()\Ag()saq+
  2. Ag()|Ag(,1M)saq+
  3. Ag(,1M)|Ag()+aqs
  4. Ag()|Ag()saq+
  5. Ag()/Ag(,1M)saq+


B is the correct answer. When silver metal is placed into a solution of silver ions, an equilibrium forms according to the following equation: Ag()Ag()+esaq+.

When we write the cell notation, we need to first know if the electrode in the half-cell is a cathode or an anode since the cathode and anode information in cell notation are written differently—in the opposite order. In this question, we are not told whether the half-cell is acting as an anode or a cathode, so we can assume that it is not part of an electrochemical cell, and thus, we write the electrode or solid metal information first by default. The solid metal is Ag, and so we write this first as Ag()s

Next, we write a vertical line to indicate that a new phase follows; in other words, a phase boundary exists: Ag()s|

And then, we write the identity of the ions in solution, followed by its concentration: Ag()Ag(Msaq|,1)+

Note that the question does not tell us the concentration of the solution. And so, it could be argued that the correct answer is D. However, typically, the electrolyte solution concentration is stated in the cell notation.

There are many types of half-cells. Some have a poorly conducting element, such as a nonmetal, as the oxidized or reduced species. So, an inert electrode, for example, gold, platinum, or graphite electrodes, can be used to conduct electrons in the circuit. In the following electrochemical cell, magnesium is oxidized at the anode, and the cathode is composed of an inert platinum electrode, but it is hydrogen gas that is reduced in the redox reaction: Mg()Mg(MH(MH()Pt()saqaqgs|,1),1)||2++2

Sometimes, a reduction or oxidation that occurs in a half-cell is not between a metal and its ions, but between ions. For example, Fe3+ can be reduced to Fe2+ using a platinum electrode. The reduction reaction is Fe()+eFe()3+2+aqaq

and the electrode notation for this half-cell is Fe3+, Fe|Pt()2+s

The potential difference between two different half-cells depends on the composition of each half-cell. The more reactive an element is, the larger the potential it has. It is difficult to compare different metals in terms of both oxidation and reduction at the same time. For that reason, we compare the potential of metals to be reduced. In other words, we compare the ease with which a metal is reduced. In this way, we can get a meaningful comparison between the potential of different half-cells. We talk about the reduction potential or electrode potential, 𝐸, of an electrode.

The potential of an electrode is influenced by the temperature, pressure, and concentration of the solution it is immersed in. The reduction potentials of different electrodes are compared under a set of standard conditions to give standard reduction potentials 𝐸 (also known as standard electrode potentials). The standard conditions are 1 atm, 25C, and 1 mol/L of electrolyte concentrations. It is necessary to use the standard conditions so that we can quantitatively compare electrodes.

Definition: Standard Reduction Potential 𝐸

It is the potential difference between an electrode and the standard hydrogen half-cell under a set of standard conditions.

The standard hydrogen electrode, or SHE, is used as a reference electrode against which all other electrode reduction potentials are measured. The half-cell notation for the SHE is H(MH(atmPt()+2aqgs,1)|,1)| and the diagram shows its setup. It consists of a platinum or other inert electrode in a glass tube through which hydrogen gas is bubbled at 1 atm pressure. The SHE is immersed in a solution containing hydrogen ions, H+, of concentration 1 mol/L, with the acid source being either hydrochloric or sulfuric acid. The entire system is kept at 25C. The SHE is assigned a reduction potential, 𝐸, of 0.00 V and all other reduction potentials are measured relative to this value.

When the SHE is being used to determine the 𝐸 of another electrode, the voltmeter that is used is a high-resistance voltmeter. This is necessary to prevent current from flowing through the voltmeter during measurement. If current flowed through the voltmeter, the potential difference would decrease and the reading would be inaccurate.

Example 3: Identifying Conditions for Measuring Standard Electrode Potential

The standard electrode potential, 𝐸, is measured under standard conditions. Which of the following is not a standard condition adhered to when measuring these values?

  1. Temperature of 298 K
  2. Solution concentration of 1 M
  3. KNO3 solution in the salt bridge
  4. Pressure of 1 atmosphere (when gases are involved)
  5. Measuring against the standard hydrogen electrode


Standard conditions are 25C (which is 298 K), 1 M solution concentration, and 1 atm pressure. Standard electrode potentials are measured under these conditions against, or relative to, the reference electrode: the standard hydrogen electrode, whose electrode potential is set at 0.00 V by default. So, we can rule out answer options A, B, D, and E.

It does not matter which salt is used in a salt bridge, as long as the ions do not interfere with or take part in the redox reaction, and as long as their concentrations remain constant and they do not precipitate out of solution. The salt bridge ions are merely present to create electrical contact between the two half-cell solutions. So, answer C is the correct answer.

The SHE can act as a cathode or an anode, depending on the reactivity of the other electrode. When the SHE is acting as a cathode (in circuit with a more reactive electrode), its half reaction is 2H(M2eH(atm+2aqg,1)+,1)

When the SHE is acting as an anode (in circuit with a less reactive electrode), its half reaction is H(atm2H(M2e2+gaq,1),1)+

𝐸 values for many metals and other substances are listed in a handy reference table called the electrochemical series. The electrochemical series and how it is used is covered in more depth in another explainer.

Example 4: Identifying the Position of an Equilibrium and Which Electrode Will Have a Larger Potential

The metal magnesium is known to be more reactive than the metal copper. When each metal is placed separately into a solution of its own ions, creating a half-cell, two separate equilibria are established as shown in the following two equations: Mg()+2eMg()Cu()+2eCu()2+2+aqsaqs

  1. In which of the two equations does the equilibrium lie further to the right-hand side?
    1. The magnesium ions turning into atoms of magnesium metal
    2. The copper ions turning into atoms of copper metal
  2. Which of the half-cells would show the biggest potential difference when compared to the standard hydrogen electrode?
    1. The magnesium half-cell
    2. The copper half-cell


Part 1

We are told that magnesium is more reactive than copper. This means that magnesium is more easily oxidized to the Mg2+ ion than copper is oxidized to Cu2+. So, in the magnesium solution, in the first equation Mg()+2eMg()2+aqs the equilibrium will lie more to the left-hand side of the equation, with a large concentration of Mg2+ ions in solution. In the copper solution, in the second equation Cu()+2eCu()2+aqs the equilibrium will lie more to the right-hand side of the equation, with a smaller concentration of Cu2+ ions in solution. So, the correct answer is B.

Part 2

Since the standard hydrogen electrode has a reduction potential of 0.0 V, we can say that the more reactive a metal is, the larger its potential difference, relative to the SHE, will be. Since magnesium is more reactive than copper, we can conclude that the magnesium half-cell would thus show the biggest potential difference when compared to the standard hydrogen electrode. So, the correct answer is A.

Let us summarize what has been learned in this explainer.

Key Points

  • When a metal electrode is placed in a solution of its ions, an equilibrium forms between the metal and its ions. This is called a half-cell.
  • Different electrodes have a different potential to be reduced. This potential difference can be measured with a voltmeter.
  • When two half-cells are joined by connecting their electrodes using a wire and their solutions using a salt bridge, an electrochemical cell is formed.
  • A spontaneous redox reaction occurs in a galvanic cell and this chemical energy is converted to electrical energy. Current flows in the wire.
  • Standard reduction potentials for many substances have been determined relative to the standard reduction potential of the standard hydrogen electrode, and these values are listed in a table called the electrochemical series according to the reducing ability of the substances.

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