Part I: Electrochemical Cells

© Silvia Kolchens

Pima Community College

Electrochemistry studies the relationship between a chemical change and electrical work. Unlike many other forms of energy conversion (conversion of fossil fuels, wind energy, hydrothermal power into electrical work), electrochemistry converts chemically stored energy directly into electrical work: no moving parts, no steam engines, turbines etc. As such this is a very efficient way of converting energy. Electrochemical principles have many applications in our daily lives. Examples include:

A battery houses a spontaneous chemical reaction that releases free energy to produce electricity. Non-spontaneous chemical reactions absorb free energy from an external source of electricity. Examples for these non-spontaneous reactions are: recovery of metals from ore, electroplating of surfaces, recharging a battery.

The principles of electrochemistry can be studied through the use of electrochemical cells. The simplest electrochemical cell is a zinc wire submersed in a copper solution. A spontaneous reaction occurs in which the zinc dissolves and copper metal is formed:

Zn-wire in CuSO4solution

This is a spontaneous reaction that incorporates a redox reaction to produce electrical energy. However, the energy in the form of electrical work cannot be harvested in such a setup. One needs to construct an electrochemical cell, which is typically composed of two half cells. There are generally two types of electrochemical cells:

The typical setup of an electrochemical cell is this:

What happens in the elctrochemical cell depicted above?

The charges on the electrodes are determined by the source of electrons and the direction of their flow through the circuit:

Zn 2+ enters the solution and electrons enter the wire

Cu 2+ remove electrons from Cu wire

Electrons are generated at the anode and consumed at the cathode

Anode has excess electrons and is negatively charged relative to cathode.

There are generally two types of Electrochemical Cells:

  1. Voltaic Cell (=galvanic cell)

  2. Spontaneous reaction è DG<0

    Generate electricity: Chemical energy is converted into electrical energy, which then is used to operate a load (light bulb, radio etc.)

    Reacting system does work on its surroundings

  3. Electrolytic Cell:
Uses electrical energy to drive a non-spontaneous reaction: DG >0

Surroundings do work on the reacting system (metal plating and recovery)

The basic design of both is the same. The flow of electrons however is reversed.

The two cell types have certain design features in common:

Electrodes = Objects that conduct electricity between cell and surroundings. Electrodes are involved in the reaction or carry electrical charge.

Electrolyte = Mixture of ions (usually in aqueous solution).

Anode = Electrode at which the oxidation occurs: electrons are given up by the substance. The substance is being oxidized.

Cathode = Electrode at which the reduction occurs: electrons enter the cell and are taken up by the substance being reduced (=oxidizing agent)

Salt bridge = "liquid wire" allows the flow of ions and completes the circuit. A salt bridge can easitly be constructed from filter paper that has been soaked in an electrolyte. A commercial salt bridge constists of non-reacting ions that are immersed in a gel. The solution cannot pour out, but ions can diffuse in and out of the half cell.

A redox reaction is composed of two half reactions: a REDuction reaction and an OXidation reaction

Reduction occurs when an element gains electrons (Cu2+(aq) + 2 e- è Cuo(s))

Oxidation occurs when an element loses electron (Zno(s) - 2 e- è Zn2+(aq))

The relative charge on electrodes is opposite in the two types of cells:

Redox reactions can be balanced using a variety of techniques (Oxidation numbers, half-equations and a combination thereof). The half-equation method works best for electrochemical cells and is outlined in KOTZ and TREICHEL on p. 948-957. Another way to balance redox equations is to use oxidation numbers in combination with the half-equation method. This method is explained in the "How to Balance Redox Equations". (Some of you might have covered Redox equations in CHM 151and it will be sufficient to review the concept. Others might want to work through the module and do appropriate examples in the text.)

Just as a redox reaction is separated into an oxidation and a reduction reaction (two half-equations) an electrochemical cell is divided into two half-cells: one in which the oxidation occurs, and one in which the reduction occurs.


For Voltaic cell and an electrolytic cell:

Anode always Oxidaion (AN OX)

Cathode always Reduction (RED CAT)

Besides the copper and zinc electrodes we discussed in our example, there are other types of electrodes


Draw a complete electrochemical cell composed of magnesium and gold and their corresponding salts. Label all components of the cell, explain what happens at each electrode, and indicate the flow of electrons and ions.

Short-hand Notation for Voltaic Cells- Cell diagram:

Rather than drawing out the entire electrochemical cell, it is common to use a cell diagram to describe electrochemical cells. The Zn/Cu cell depicted in the diagram above can be summarized:

(Zn(s) | Zn 2+ (aq) || Cu 2+ (aq)| Cuo(s) The anode (oxidation) is typically written on the left side, the cathode on the right side, the symbol "|" indicates a phase boundary between metal and electrolyte solution, and the symbol "||" indicates the salt bridge.


Anode (Oxidation) | Electrolyte 1 || Eletrolyte 2 | Cathode (Reduction)


Write a cell diagram for an electrochemical cell composed of aluminum and silver and their corresponding salts

Cell Voltage:

We consider the two half reactions:

Zn (s) ó Zn 2+ + 2 e- equilibrium lies to the right

Cu (s) ó Cu 2+ + 2 e- equilibrium lies to the left

A spontaneous reaction occurs as a result of the different abilities of the metals to give up their electrons. The higher the electronegativity of an element, the less the ability to give up electrons. The voltage or cell potential of an electrochemical cell is then determined by the relative ability of the two electrodes to lose or gain an electron.

When we measure the voltage of the Zn/Cu cell, then we find a value of +1.1 V. The voltage generated by an electrochemical cell depends on the material used for the anode and cathode. A positive cell potential is found for a spontaneous overall reaction:

Ecell > 0 and DG<0

The more positive Ecell, the more the reaction will proceed to the right:

Zno + Cu2+ è Zn2+ + Cuo

A negative cell potential is found for a non-spontaneous reaction.

The cell potential is measured in Volt [V], and the electric charge in Coulomb [C]. One electron has an electric charge of 1.602x10-19 Coulomb.

1 V = 1 Joule/ Coulomb = 1 J/C

Voltages of some voltaic cells:

Common dry cell battery: 1.5 V

Lead/acid battery (car battery): 12 V

Calculator battery: 1.3 V

Electric eel: 720 V

Standard Cell Potentials:

Standard cell potentials are measured under standard conditions. Various electrode materials are measured against the hydrogen electrode. The hydrogen electrode employs the reaction:

2 H+(aq) + 2 e- è H2(g)

Standard conditions are: 1 atm (gases), 1 M (solutions, T=298 K

Deviations from standard potentials occur with

Change in concentrations of reactants

Energy loss due to heating of the cell and external circuit

The standard cell potential is the potential associated with a given half-reaction. For hydrogen, the standard cell potential is zero. Cell potentials for other materials are measured in reference to hydrogen and are summarized in Standard Reduction Potential Tables (p. 970 in Kotz and Treichel).

How can we determine the cell potential for any given electrochemical cell using standard reduction potentials?

Cell potentials for half-reactions are typically listed as reductions (reduction potentials), i.e. a positive value for Ecell tells us the electrode material easily accepts (gains) electrons, a negative Ecell means that the electrode material donates (loses) electrons in reference to hydrogen. Let's find the cell potential for the above Zn/Cu cell from standard reduction potentials:

Reduction reaction is:

Cu 2+ + 2 e- è Cuo and we find a value of +0.337 V

Oxidation reaction is:

Zno è Zn2+ + 2 e- and we find a value of +0.763 V*

Adding up both half equations and both standard potentials yields the overall cell voltage:

Cu 2+(aq) + Zno è Cuo + Zn 2+ (aq) Ecell = +1.1 V

(*since only reduction potentials are listed we need to reverse the sign of the listed cell potential)


As a chemical engineer working for the company Umweltfreundlich, Inc. you are asked to construct an environmentally friendly yet powerful battery. All components of the battery should pose little hazard to the environment. Choose appropriate materials, draw a sketch of your design, list potential hazards when operating or discarding the battery and estimate the cost of your battery. Suppose you have the following electrode materials and their corresponding salts available: Pb, Ag, Au, Zn, and Na.

Materials used:

Design of electrochemical cell:

Calculated Cell Potential:


© copyright Silvia Kolchens, Pima Community College 2000