I try to explain clearly the principles of this rather confusing but very useful science.


  1. Introduction
  2. Oxidation and Reduction
  3. Electrolytic Cells
  4. Electrode Potentials
  5. Corrosion and Dry Cells
  6. Electroplating
  7. Dependence of the Electrode Potential on Concentration
  8. References


In 1797 the English physician George Pearson laboriously charged Leyden Jars at his electric machine, then discharged them through water, carefully collecting the gases that appeared. Finally, he mixed the gases in a dry container and made a spark with his machine. Drops of water collected on the walls of the container when it cooled. He had decomposed water into its constituents, and then recombined them again. The world took little notice.

In 1800 Alessandro Volta reported the results of his recent studies to the Royal Society of London, of which he was a member. His momentous achievement was the column, or "pile," of discs of silver, zinc and leather moistened with salt solution, repeated over and over. An alternative was the couronne des tasses, a ring of cups joined by arcs of silver and zinc alternately, filled with dilute salt solution. When the ultimate members of the pile or crown were connected by a conductor, a permanent electric current flowed. Much care was taken to show that it had the same qualities as the electricity from a static machine, principally that it could give a shock, or fuse a fine wire. Electricity was now available in unprecedented amounts with no exertion, but at a much lower pressure.

In 1801 Nicholson and Carlisle caused the current to pass through a dilute salt solution and collected the gases that were evolved at the separate platinum electrodes. Twice as much gas was produced at the electrode connected to the zinc of the pile as at the electrode connected to the copper. They may have used different metals, but copper and tin will do as well as silver and zinc. The main idea seemed to be that the metals had to be different. A glowing splint thrust into the copper gas burst into flame, while a glowing splint approaching the zinc gas caused a loud "pop." These gases were, of course, oxygen and hydrogen, only very recently recognized as elements and the constituents of water in Lavoisier's new chemistry.

Electrolysis became the rage. Cruikshank developed more powerful batteries to produce the current, and many experimenters were absorbed in seeing what the electric current would do. None was more successful than Humphrey Davy, the young Cornishman who was the darling of the ladies who attended his lectures at the Royal Institution. This story, and the later work of his assistant Michael Faraday, are told elsewhere. Electrolysis demonstrated beyond a doubt the central role played by electricity in chemistry. In fact, it is the basis of chemistry, and electricity holds everything together. Just how it all worked was by no means clear at the beginning, but the picture became clearer at the end of the century when the electron was finally unambiguously discovered in 1896. One century exactly holds this entire development.

The reader is encouraged to try Nicholson and Carlisle's experiment. All that is needed are a couple of short lengths of wire, two test tubes, a bowl that can be filled with water, and a DC power supply. The most difficult part is arranging for the test tubes to be held more or less vertically as the experiment proceeds. Prepared apparatus is available, but it is more fun, and much cheaper, to do it yourself. Arrange the wires from the power supply so that they point upwards and are stripped of insulation for about 1 cm. Then fill the bowl with water, and mix in a little salt. The purpose of the salt is to make the water conducting, so things will take place at a reasonable speed. Immerse a test tube in the water so it is filled, then lift it up and put it over the wire without letting the water run out. Repeat for the other test tube. Atmospheric pressure will hold the water up. If you happen to have some litmus paper, put a small piece in the vicinity of each electrode. Now turn on the power, and observe. A few volts will be enough, and the current will be small.

Bubbles will be seen almost immediately at the wire connected to the (-) terminal of the power supply. The higher the voltage, the faster they will form. The gas will collect in the top of the test tube, and it will take only 30 minutes or so to collect a good amount. The litmus paper will turn blue, showing that the solution becomes basic in this area. Meanwhile, not much will be happening over at the (+) side if you have used a piece of copper hookup wire. We will see why not much is happening further on. If you use a piece of stainless steel wire, insulating it except at the end with, say, some paint, bubbles will be seen, and the litmus paper will turn pink. I did not have any platinum around, but it would make the best electrode. It is not too easy to collect considerable oxygen in a simple experiment. When you have most of a test tube of hydrogen, remove the test tube from the water, holding your thumb over the top. Light a match with the other hand, and when it is close to the test tube, raise your thumb and you should hear a definite pop. This is by no means violent with the amount of hydrogen available, so there is no reason to cringe.

Hydrogen is very obliging at exploding over a wide range of concentrations, something that will become obvious in a "hydrogen economy." The amount of energy required to make hydrogen is greater than that recovered when it burns, even with 100% efficiency, so it is certainly not a primary source of energy, only a means of passing it around. Electrolysis is not the most economical way of producing hydrogen--reforming of hydrocarbons or the decomposition of steam by hot coke are. They have been used for many years, and are the source of commercial hydrogen. When hydrogen burns, it produces H2O, an important "greenhouse" gas.

Oxidation and Reduction

Some atoms require no external assistance to "get it on." Two neutral oxygen atoms will stick together strongly if they meet. The electrons of one nestle into the positive parts of the other, something called forming a covalent bond, and the system gives up a little energy, so it becomes bound. It is not easy to separate the two thereafter.

In many cases, however, it is necessary to gain or lose electrons to get to a lower energy (or higher entropy) state. Consider a solution of blue vitriol, cupric sulphate, CuSO4. The water molecules have picked the crystal apart, so the blue solution consists of positively charged Cu++ ions and negatively charged SO4-- ions, each surrounded by a halo of water molecules attracted by the charges. The solution as a whole is neutral, with equal numbers of positive and negative charges, over distances of a few intermolecular spacings, say 1 nm or so.

The word "ion" is Greek for "goer," since it goes to one or the other of the electrodes of an electrolytic cell, driven by the electric field. It is one of the terms based on Greek created by Michael Faraday (1791-1867) with the help of the classical scholar William Whewell for the new science of electrochemistry. Electrode, anode, cathode, anion, cation and electrolyte are further examples. This was quite well done!

Now put a strip of zinc into the solution. The water molecules would like to take this crystal apart, as they did the cupric sulphate. If they do, then Zn++ ions would be formed, and two electrons, written 2e-, would remain in the metal, causing it to have a net negative charge. This charge attracts positive ions, so whatever small tendency the zinc has to dissolve would soon be balanced by its tendency to return to the metal crystal. A thin layer of opposite charges results, like a surface density of electric dipoles pointing outwards, over the strip.

This dissolving of a little zinc would take place with or without the copper and sulphate ions, even in pure water. However, the copper ions are not ignorant of the dipole layer created by the zinc, and are attracted to the oppositely charged metal. When they arrive, it is likely that they will snatch up a couple of electrons each, in the reaction Cu++ + 2e- → Cu. They will then accumulate on the surface of the zinc, as long as they exist in the solution, and as long as bare zinc is available to replace them in the solution.

The reaction Zn → Zn++ + 2e- is called an oxidation in chemistry. No oxygen is actually involved, but this is similar to what happens to the zinc when it burns: 2Zn + O2 → 2ZnO. In ZnO, the zinc is formally considered to be Zn++ and the oxygen to be O--, though the actual bond need not correspond to the complete transfer of the electrons from zinc to oxygen. The reaction Cu++ + 2e- → Cu is called a reduction. When metals are produced from their ores, they are said to have been reduced, and this is formally equivalent to having lost oxygen: 2Cu2O + C → 4Cu + CO2.

When we put the zinc in the vitriol, the zinc is oxidized and the copper is reduced. The net effect is the removal of two electrons from the zinc metal and the addition of the two electrons to the copper ion. Everything must remain electrically neutral, so we must take care to balance electric charge in our descriptions of what is happening. The great opportunity that now comes over the hill is to perform the oxidation in one place and the reduction in another, moving the electrons over a conductor as an electric current. This would allow us not only to control a reaction, as in Nicholson and Carlisle's electrolysis, but also to create an electric current as in Volta's pile.

Electrolytic Cells

To do this, we must not allow the reactants of the two reactions to mix. That is, we must put the zinc in one solution, and the vitriol in another. When we do this, there must also be a way to current to flow from one solution to the other, and this can be done by salt water. The details are a little complicated, but the idea is simple so we will not worry about the details. In electrochemistry, this can be done by something called a "salt bridge" which is just an interval of salt solution in contact with the reaction solutions at its two ends. A convenient salt bridge can be made from agar gel made with KCl solution. It is a wobbly solid, and keeps the solutions separated. A salt bridge acts like a wire, but actually using a wire would involve additional electrodes and reactions, which is not wanted here. The combination of two electrodes and electrolyte is called an electrolytic cell. An electrolyte is any medium, usually fluid, that conducts electricity, normally by ionic conduction.

An electrolytic cell in which this has been done is shown in the diagram. The liquid in the cell, through which a current passes, is the electrolyte. The external circuit is connected to the electrodes, which are usually metal or carbon, and often take part in the cell reaction. Sometimes they merely carry current, and are inert electrodes. Platinum or carbon are inert electrodes. The cathode is the electrode near or at which the reduction occurs, and electrons move into the cathode from the external circuit. The anode is the electrode near or at which the oxidation occurs, and is a source of electrons to the external circuit. The name assigned to an electrode depends on the reaction that occurs there, and can change for different conditions. Electrons always move from anode to cathode in the external circuit. Since electrons are assigned a negative charge, the direction of current flow is opposite to the direction of electron flow. The current always flows from cathode to anode. This can cause confusion, but is perfectly correct and leads to less confusion overall than making special definitions. Current flows in the direction of motion of a positive ion, and opposite to the direction of motion of a negative ion. In electrochemistry, we have charge carriers of both signs, so the current could not always be in the direction of motion, however much we would like it to be.

The terms cathode and anode are used with electronic devices in general, not just electrolytic cells. The conventions for a semiconductor diode and a thermionic triode are shown at the left. The current direction for each is in the direction of "easy conduction" when the diode is forward-biased and current in the thermionic tube, which is carried by free electrons, is flowing. In each case, the anode is made positive with respect to the cathode. Electrons enter the cathode and leave from the anode.

In this cell, zinc dissolves at the anode, releasing electrons that move to the anode, through the external circuit. making it negative so that it attracts positive ions. Here there are copper ions in the solution, which are attracted to the copper electrode and are discharged there. Copper is "plated out" on this electrode. The electrode could just as well be inert, platinum or copper, and the same thing would happen. The electrode takes no part in the cathodic reaction in this case. Sulphate ions, not indicated in the diagram but required for electrical neutrality, migrate toward the salt bridge in the electric field that extends from cathode to anode, repelled from the SO4-- left unbalanced there and attracted to the Zn++ created there. The electrons that would normally provide electrical neutrality have departed the scene over the external circuit. At the salt bridge, they are comforted by the Na+ ions and their current is continued on by the Cl- ions. At the other end of the salt bridge, Zn++ ions have been repelled from the anode, and find the Cl- ions to balance them. The current is continuous, carried by electrons in the external circuit, and by the ions of both signs in the electrolyte. Current never begins or ends in the steady state.

The electrons move like lightning, while the ions move like creeping things. This means that if the reactions take place rapidly, as they will if the cell is short-circuited, positive charge builds up at the anode and negative charge builds up at the cathode. This electric field is opposite to the one created by the reactions, and more or less neutralizes it. This causes the current through the external circuit to decrease. The polarity of the cell is shown in the diagram. The cathode, the copper side, becomes (+) while the anode, the zinc side, becomes (-). This is the polarity required to drive the current in the external circuit in the direction shown (opposite to the electron motion). The voltage is related to the current by E = IR, where I is the current in amperes, R is the resistance in ohm (including the resistance of the electrolyte and other such effects), and E is the terminal voltage of the cell. The buildup of electric charge in the electrolyte is called polarization. In this kind of cell, the internal resistance can be made quite low, and it does not suffer from noticeable polarization. These are, in fact, the reactions that occur in the Daniell Cell, a great improvement on Volta's Pile, which polarized rather quickly. E is about 1.1V, and was the original standard of voltage.

It is not necessarily the same electrons that appear in the reactions at the anode and cathode. The external circuit is like a pipeline, in that when an electron is added at one end, it lowers the potential of the wire. Then one is taken from the other to restore electrical neutrality in the wire. The actual electrons drift rather slowly from end to end, but the effect of adding an electron at one end is felt at the other as quickly as an electromagnetic pulse can get there.

A gravity cell, once very popular in the United States on closed-circuit telegraph circuits, is shown at the left. It is Daniell cell, but the two electrolytes are in direct contact. The blue electrolyte at the bottom is saturated CuSO4 solution, kept saturated by crystals (not shown) packed around the copper cathode. The clear electrolyte at the top begins as a very dilute sulphuric acid solution, and gradually becomes filled with Zn++ ions as the cast zinc anode at the top dissolves. The downward migration of the Cu++ ions and the greater density of the copper sulphate solution keeps the blue copper ions away from the anode. Of course the cell must not be moved or disturbed, and the current through it must be kept flowing to maintain this state, or the copper ions would diffuse througout. The usual current was 20 to 100 mA, and the internal resistance of the cell was about 2 Ω, for the normal size cells. The cell could safely be shorted, since it would draw at most 500 mA. When the zinc anode was eaten away, the upper solution was becoming too concentrated anyway, and the cell was poured out, cleaned and a new anode supplied. The cathodes were sent for reclamation of the copper when they became too heavy.

Now suppose we attach our own power source to the cell, with a polarity that opposes the voltage of the Daniell Cell, as shown at the right. By increasing the applied voltage, the current can be made to decrease, and finally to reverse. When the current is exactly zero, the applied voltage is equal and opposite to the cell voltage. This is, in fact, the way the cell voltage is measured, by an instrument called a potentiometer. Electrode reactions are, at least in principle, reversible if all the reactants are present. We shall assume them to be ideally reversible, which is not true in practice. When we have caused the current to reverse, electrons are now supplied to the zinc and taken from the copper. This means the zinc is being reduced, and the copper oxidized, so the names of the electrode change. The apparent polarity of the cell is the same, with the copper positive and zinc negative. It is very important to keep this straight. The electrode reactions can either go their natural ways, or can be forcefully reversed.

Electrode Potentials

Now we come to the very interesting question of determining in which directions two electrode reactions will proceed spontaneously when they take place at the electrodes of a cell. In this case, we can say that zinc has a greater tendency to lose electrons than copper. When we put them together, the zinc sticks electrons down the throat of the copper. The ionization potential of an isolated zinc atom is 9.391V. This is the energy that has to be supplied to knock an electron out. The ionization potential of copper is only 7.723V, so it would appear that zinc holds on to its electrons even more firmly than copper. In electrochemistry we are not dealing with isolated zinc and copper atoms, but water is always lurking in the background, and the zinc and copper are in crystalline form. How much they love their electrons is determined by very complex interactions, so we must see what happens empirically.

If we had to look at all combinations of electrode reactions and see which way they go by measuring the voltages with a potentiometer, it would be very tedious. However, if we consider connecting cells in series, we realize that the potentials are additive. This means we only have to measure each electrode reaction against some standard electrode reaction. The standard electrode is the normal hydrogen electrode, realized by bubbling hydrogen gas over a platinum surface. The reaction is 2H+[1N] + 2e- → H2[1 atm]. Note that the concentration of the hydrogen ion is stated (1 Normal, which is quite acidic) as well as the pressure of the hydrogen gas (1 atm). Cell voltages depend on the concentration of the reactants, something that we will discuss in more detail below. For now, we will just assume that all the concentrations are standard.

If we now use our zinc electrode with a hydrogen electrode, we find that the zinc becomes the anode and is oxidized, going into solution, while the hydrogen electrode becomes the cathode and is reduced, hydrogen gas being liberated. The hydrogen electrode is the positive pole, and the zinc the negative pole. The zinc is observed to be at a voltage of -0.761V with respect to the hydrogen electrode. This voltage is called the single electrode potential, or reduction potential of the cathode reaction Zn++ + 2e- → Zn. Please do not complain to me about the convention; it seems to be what chemists use. The oxidation potential is the negative of the reduction potential, 0.761 V in this case. The more negative the reduction potential is, the more the reduction reaction would rather be in the other direction. That is, the easier it is to get electrons out of it. The electrode potentials that are shown in tables are, unless otherwise specified, the standard electrode potentials, since they refer to specific concentrations of the reactants and a specific temperature.

Sometimes anode (oxidation) reactions are written instead of cathode (reduction) reactions, and sometimes the signs of the single electrode potentials are changed. Just be careful, and remember that the zinc becomes negative, and the anode, relative to hydrogen. Any reaction above (in the way the table is usually written down) hydrogen can displace hydrogen from water, since hydrogen becomes the cathode. The standard electrode potential is negative for such a reaction.

In any electrolytic cell, the cathode and anode reactions proceed in step as the electrons are transferred. An overall reaction can be written in which the electrons do not appear, since as many are created by oxidation as are absorbed by reduction. The equations can be used to find out how much product will be produced for a given amount of charge transferred. One equivalent weight (the molecular weight divided by the valence of the ion) is produced or consumed for every faraday of charge, 96,480 C. This is approximately the charge transferred by a current of 1 A flowing for 28 hours. The faraday is just Avogadro's number times the absolute value of the electronic charge. The reactions can be stopped, in principle, by opening the external circuit, and even reversed by applying an external potential of the opposite polarity. Of course, in reality reactions may go on locally with no current, and the reaction may not be reversible for some reason (such as the production of gases that escape from the cell).

The single electrode potential for the reaction Cu++ + 2e- → Cu is 0.344V. That is, it is below hydrogen, and will become the cathode, while the hydrogen electrode becomes the anode. It will be 0.344V positive with respect to the hydrogen electrode. Copper will not displace hydrogen from water; in fact, it will oxidize hydrogen instead. If we combine the two reactions, Zn → Zn++ + 2e- will have a voltage of 0.761V, and Cu++ + 2e- → Cu a voltage of 0.344V, for a total of 1.105V for the Daniell Cell. This presumes that the concentrations of copper and zinc ion are each 1N, or one equivalent weight per litre of solution. The zinc will be the anode, and the copper the cathode, when the cell is delivering current spontaneously. The electrode potentials allow us to determine which way the reactions will go, and the voltage.

The least electrode potential is -3.024V for Li+ + e- → Li, and the greatest is +3.03V, for F2 + 2H+ + 2e- → 2HF. A lithium-fluorine battery would be a nasty thing, but would give 6V! The lithium side would be negative, and the fluorine side positive. Electrode potentials can be calculated from the free energy change in a reaction if this is known. Otherwise, they can be used to find free energy changes.

Pure water is slightly ionized. [H+] = [OH-] = 10-7N. The quantities in square brackets are concentrations. At these concentrations, the cathode reaction 2H+ + 2e- → H2 has an electrode potential of -0.414V. This means that if it were connected to a standard hydrogen electrode, it would become the anode and be at this voltage relatively. The cathode reaction O2 + 4H+ + 4e- → 2H2O has an electrode potential of 0.815V. Spontaneously, then, oxygen would become water at the cathode and hydrogen would be oxidized at the anode to supply the hydrogen ions necessary for the water. That is, oxygen and hydrogen would burn to water, establishing a potential of 1.229V that could supply an external circuit. We have, indeed, an oxygen-hydrogen fuel cell. There is a number of problems to be solved before the cell would be practical, but such fuel cells already exist.

When we apply an external voltage of greater than 1.229V, with the oxygen side positive and the hydrogen side negative, then the oxygen side becomes the anode and the hydrogen side the cathode. Oxygen and hydrogen gas are produced at the expense of the water. This is, of course, the electrolysis of water, which you can easily observe at home. Suppose we have added salt to the electrolyte to increase its conductivity. Why do we not get sodium metal (which would form sodium hydroxide) and chlorine instead? The reaction Na+ + e- → Na has an electrode potential of -2.714V, while the reaction Cl2 + 2e- → 2Cl- has an electrode potential of 1.3583V. Even at 1N concentrations, these voltages are far beyond those that result in oxygen and hydrogen, which are much more easily produced. Sodium and chlorine would much rather remain as ions. In very concentrated solutions, chlorine is observed to be given off with the oxygen.

It should be clear now why no oxygen is seen when you electrolyze water with copper wires. Copper can be oxidized with only 0.344V, while discharging oxygen requires 0.815V. Stainless steel is harder to oxidize, so you get a little oxygen. With platinum, the cell would have no option but to produce oxygen. Electrode potentials can be used to figure things like this out.

Corrosion and Dry Cells

Corrosion is mainly due to electrolysis occurring around impurities and other centers that may create a local difference in electrode potentials. The area that becomes the cathode is safe, but the area that becomes the anode is subject to oxidation. In some cases, this can be overcome by deliberately creating a cell with the material to be protected as cathode, with the anode consisting of some material higher in electrode potential, that will react first. This is called cathodic protection. For example, consider iron. The reaction Fe++- → Fe has an electrode potential of -0.441V, meaning that it is fairly reactive and in danger of becoming an anode, with corrosion resulting. However, magnesium is even more reactive, with Mg++ + 2e- → at -2.34V. Magnesium is higher in the electrochemical series than any conceivable anode in this case.

If we are protecting a pipeline, a piece of magnesium is buried near the pipeline, and connected with the pipeline by a length of copper wire. Normally, copper would act as a cathode relative to the iron, and corrode it, but here the magnesium takes care of everything and becomes an anode to iron, copper and all. The magnesium becomes negative and the pipe positive, so that current flows from the pipe to the magnesium. The magnesium sacrifices itself to protect the iron, so it is called a sacrificial anode. It must be replaced periodically when it becomes too wasted. The actual reaction that occurrs at the magnesium probably involves magnesium hydroxide, since the earth is alkaline, but the electrode potential is then even more negative. This will not work if the ground is too dry to carry the ion current back to the iron pipe. Cathodic protection is used not only for pipelines, but also for ships, buildings, bridges and other structures subject to corrosion. It also shows that there is a danger in mixing copper and iron piping where moisture can provide a path for the electrolytic currents.

The electromotive series of the metals is a rough guide in such matters. The series is: Li, Rb, K, Ba, Sr, Ca, Na, Mg, Be, Al, Mn, Zn, Cr, Fe, Cd, Co, Ni, Sn, Pb, H, Sb, As, Bi, Cu, Hg+, Ag, Pd, Hg++, Pt, Au. A metal will protect any metal below it. A metal above H will displace H2 from weakly acidic water.

The carbon-zinc "dry" cell has a cathode of carbon and an anode of amalgamated zinc. The cell is not really dry, of course, but the electrolyte is pasty and must remain moist, so the cell is sealed. Amalgamating the zinc makes the surface uniform, preventing local action. The anode reaction only consumes the zinc, not the mercury, which serves for the life of the cell. Manganese dioxide, MnO2 is packed around the carbon to absorb any oxygen that is formed and prevent it from increasing the internal resistance of the cell. The electrolyte is ammonium chloride, NH4Cl, mixed with zinc chloride, ZnCl2. The open circuit voltage is about 1.6V, usually considered to be 1.5V when any current is drawn. The anode reaction is (the reverse of) Zn(NH3)4++ + 2e- → Zn + 4NH3, -1.03V, or something similar. The manganese dioxide near the cathode may react like MnO2 + 4H+ + 2e- → Mn++ + 2H2O, 1.28V. The voltages do not add up to 1.5 because we have not considered the concentrations.

The alkaline "dry" cell also has a carbon cathode and a zinc anode, but the electrolyte is sodium hydroxide with some sodium zincate. This cell has about twice the life of the carbon-zinc cell and a lower internal impedance, but the electrolyte is corrosive and must be carefully contained in a steel jacket. Here, the anode reaction is something like ZnO2-- + 2H2O + 2e- → Zn + 4OH-, even higher on the scale at -1.216V. The cathode reaction may be related to O2 + 2H2O + 4e- → 4OH-, 0.401V.

The Nickel-Cadmium rechargable cell is a relative of the Edison alkaline storage battery. The cathode is nickel oxide, the anode is cadmium (taking the place of the similar zinc), and the electrolyte is KOH + LiOH, giving 1.35V on open circuit. The cathode reaction may be Ni(OH)2 + 2e- → Ni + 2OH-, -0.66V, while at the anode Cd++ + 2e- → Cd, -.4021V. Again, the voltages do not add up because the concentrations are not standard. Since Cd(OH)2 is insoluble, it stays where it is as the cell discharges, as does the nickel. When the current is reversed, it changes back to Cd as the Ni changes back to NiO.

A rechargable battery is one in which the products of the forward reaction do not escape or diffuse away from the electrodes, so that the electrode reactions will be reversible. In the lead-acid cell, PbSO4 is formed at and clings to both electrodes. In the nickel-cadmium cell, the nickel hydroxide and cadmium hydroxide are insoluble and stay where they are. Even the alkaline cell is reversible to some degree, since the zinc hydroxide does not get far, but the cell is not designed with this in mind.


It happens that the cheap and strong metals, such as steel, are subject to rapid corrosion, while metals that resist corrosion, such as silver, gold, copper, tin, zinc, lead and nickel, are either weak, or expensive. The benefits of both can be realized by coating the strong metal with the unreactive metal. This protection is provided either by the fundamental unreactivity of the coating (gold, silver and copper), the formation of a protective layer (nickel, zinc, tin, lead), or cathodic protection (zinc on steel). A layer that does not provide cathodic protection (like tin on steel) must be sound and continuous, since corrosion may actually be promoted by the coating at such "pinholes."

Methods of applying the coating are: (a) hot-dipping in the molten protective metal; (b) cementation--forming a surface alloy without melting; (c) cladding--fusing the protective layer to the metal before it is rolled; (d) sputtering from a cathode in vacuum, or deposition from a vapor; and (e) electroplating. We will discuss only the last method here. See articles on the individual metals for examples of the other methods.

In electroplating, the object to be coated is made the cathode in an electrolytic cell. The electrolyte contains an ion that is reduced to the metal at the cathode surface in proportion to the current passing through the cell. One equivalent weight (atomic weight divided by the electrovalence) is deposited for each faraday of charge at an electrolytic efficiency of 100%. The anode may be the same metal that is being plated, where the metal ion is oxidized and replenishes the electrolyte automatically. Alternatively, the anode may be inert, and additional ions added as required. This is the case in chromium plating, where chromic acid is added continually. The voltage required is usually low, only a few volts, and the current density adjusted to give the best plating. Current densities of about 5-50 A/sq.ft. are not exceptional.

Reduction of H+ ion to hydrogen gas at the cathode competes with reduction of the metal ion. In acidic solution, only a few metals, such as Cu, Ag and Au, are beneath hydrogen in the electrochemical series, and can be successfully deposited. If the solution is alkaline, the reduction potential of hydrogen is then -0.828V, permitting many more metals to be reduced at the cathode, including Zn, Sn, Cd, Cr and Ni, usually from complex ions, such as cyanide, stannate and chromate. The pH must usually be carefully controlled to ensure the proper electrode reactions. The complex ions supply a controlled small concentration of positive metal ions for reduction, making a better and more uniform deposit.

The first step in any coating process, and especially electroplating, is the preparation of the object to be coated. There are three steps: (1) degreasing; (2) removal of oxides; and (3) surface mechanical preparation. Degreasing can use a solvent such as TCE (trichloroethylene). TCE is noninflammable, but must be used in a closed system since it is hazardous. An alternative or additional method is electrolysis in an alkaline solution of sodium carbonate, with small amounts of sodium hydroxide, trisodium phosphate and borax, at a current density of about 10 A/sq.ft. The object is made the cathode, and the copious emission of hydrogen gas gives a thorough cleaning. Next, any surface oxide must be removed, or it will react with the coating to give a weak area that may separated from the substrate. For steel, wire-brushing followed by pickling in a 4%-5% H2SO4 solution at 65-75°C is satisfactory. Brass is pickled in dilute sulphuric and hydrochloric acids mixed with a little nitric. A dip in 2%-10% H3PO4 produces a beneficial phosphatic film. Finally, the surface is roughened if this helps the adherence of the plating.

The plated surface comes from the bath in a matte form, and must be polished if it is to be shiny. Sometimes protective coatings are added as well. Chromium plate is always porous, so it cannot be used directly over steel because of the danger of pinhole corrosion. The electrolyte is chromic acid (chromium trioxide) in sulphuric acid, with an inert anode. The pH and concentration must be carefully controlled to avoid reduction of hydrogen. Chromium is usually plated over nickel, which does give a dense, impervious surface and protects the steel. Nickel is plated from an electrolyte of NiSO4, NiCl2 and H3BO3 at 50-60°C, with carefully controlled pH.

Aluminium would make a good coating, but it cannot be reduced electrolytically from an aqueous solution because it is above hydrogen for any pH and concentration. If aluminium were not passivated by its coating of Al2O3, it would release hydrogen when dipped in pure water. It is this that prevented the easy winning of aluminium by electrolysis, as in the case of copper.

Dependence of Electrode Potential on Concentration

Let us now see how concentration affects cell voltage. We will take a special case, when we have only a concentration difference, and no reaction. That is, we consider a cell in which the same reaction occurs at the electrodes, but the concentration of an active ingredient is different. Suppose that the concentration c of an ion is maintained at a value a1 at one electrode, and at a value a2 at the other. The ions will diffuse toward regions of lower concentration with a diffusion coefficient D, such that the ion flux is q = -D(dc/dx). In equilibrium, since there is no change of concentration at any point, an electric field E is set up that exerts a force on the ions that tends to make them drift in the direction opposite to the diffusion. The velocity of an ion is v = KE, where K is the mobility of the ion in cm/s per V/cm. The ion drift flux is then q = cv = cKE. The net ion flux is then cKE - D(dc/dx), which must vanish. Then (dc/dx) = cKE/D, or dc/c = (K/D)Edx. This equation integrates to ln(a1/a2) = (K/D)V, where V is the potential difference V2 - V1.

Therefore, V = (D/K)ln(a1/a2). The ratio of the diffusion coefficient to the mobility is a function of the temperature only: D/K = kT/ne (Einstein's relation), where k is Boltzmann's constant, ne the charge on the ion, and T the absolute temperature. This remarkable relation will be proved below. We have V = (kT/ne)ln(a1/a2), which is the desired relation. If k is in erg/K and e is in esu, then V will be in statvolts. Substituting e = 4.803 x 10-10 esu, k = 1.38 x 10-16 erg/K, T = 298K, and converting from natural logarithms to common logarithms, and statvolts to volts, this becomes V = ±(0.0592/n)log(a1/a2). The correct sign can be determined by considering the picture of opposing diffusion and drift fluxes. If we have positive ions, then the electric field must be directed toward the point of greater concentration, so the potential will be higher at the point of smaller concentration.

As an example, consider the standard hydrogen electrode. If instead of 1N hydrogen ions we have 10-7N, as in pure water, then the voltage difference is V = 0.0592 x 7 = 0.414V. If we connect the two electrodes, then current will flow from the standard electrode to the other in an external circuit, and the standard electrode will be the cathode, so the electrode potential will be -0.414V.

To prove Einstein's relation, D/K = kT/ne, we appeal to the familiar Boltzmann factor that is a standard result of statistical mechanics. If the electrostatic potential as a function of position is φ, then the concentration of ions of charge ne at any point is given by c = c(0)exp(-neφ/kT), where c(0) is a constant. From this equation, we find that ln(c/c(0)) = -neφ/kT, or φ = (kT/ne)ln(c(0)/c), which is exactly the equation we found above. If we express the net flux as we did above, in terms of D and K, and set it equal to zero, we get Einstein's relation. A better proof along these lines can be made, but this should be enough to make the principles clear.


N. A. Lange, ed., Handbook of Chemistry 10th ed. (New York: McGraw-Hill, 1961). There is an extensive list of electrode potentials on pp. 1212-1218.

R. C. Weast, ed., Handbook of Chemistry and Physics, 56th ed. (Cleveland: Chemical Rubber Publishing Co., 1975). There is a large list of electrode potentials, arranged alphabetically as well as by voltage, on pp. D-141 to D-146.

L. Pauling, General Chemistry (New York: Dover, 1970). Any general chemistry text will introduced the electrochemical series and electrode potentials.

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Composed by J. B. Calvert
Created 16 November 2002
Last revised 23 November 2002