CR Scientific: Electrolysis Experiments

Anions move toward the anode;

Cations move toward the cathode.

When something is oxidized, it loses electrons. Think of iron rusting. The iron (zero charge) is oxidized to Fe³⁺. Electrons are stripped away from the iron to give it a positive net charge.

Reducing the rust back to iron involves giving back electrons so Fe³⁺ can become neutral Fe. One way of looking at this is that the magnitude of the positive charge is being reduced.

Electrolysis Experiments

Part II: Rational Design

CAUTION: Electrolysis of certain compounds could produce poisonous or unstable products.

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One can design simple electrolysis demonstrations starting with a table of standard reduction potentials. These are available in almost any inorganic chemistry reference, such as the CRC Handbook of Chemistry and Physics. However, the reduction potentials themselves don't entirely answer the complexity of what goes on in an electrolysis cell. Often there are side reactions and other variables which can be so complex that it's difficult to predict them with theory.

Let's go over some things one should at least consider when designing an electrolysis cell.

1. Potentials

Suppose our reference book gives only the standard reduction potentials. Why not oxidation, too? It's not necessary. To get the oxidation potentials, just switch the sign. For example, the reaction

Pb²⁺ + 2e^- <----> Pb

has a reduction potential of -0.126 volt relative to the Standard Hydrogen Electrode. The reverse reaction (in other words, the oxidation of Pb to Pb²⁺) therefore has an oxidation potential of +0.126 volt relative to the Standard Hydrogen Electrode. The positive value means the oxidation of lead is actually favored. Metallic lead tends to oxidize to Pb²⁺ ions, although spontaneous transformation is slow.

Oxidation / reduction potentials indicate energetic favorability, but they tell us nothing about how long a reaction could take.

A positive potential means the products are energetically favored over the reactants. In other words, the products have a lower energy than the reactants.

A negative potential means reactants are energetically favored over products; in other words, energy would have to be added to the reactants in order to turn them into the product(s).

2. Overvoltage

Suppose one reaction has a higher (i.e., less negative) potential than another.
Let's say Half-Reaction A has a potential of -1.455 volts, while Half-Reaction B has a potential of -0.815 volts. If both half-reactions could theoretically occur in solution, we might expect Half-Reaction B to occur preferentially because its potential is less negative. Sometimes this doesn't happen!

The half-reaction
2H₂O <---> O₂ + 4H⁺ + 2e^- (-0.815 volt)
is energetically less difficult than
Pb²⁺ + 2H₂O <---> PbO₂ + 4H⁺ + 2e^- (-1.455 volt),
at least according to the oxidation potentials.

In real life, however, the formation of oxygen from water generally requires an overvoltage in order to occur; the actual overvoltage depends on the type of anode material. The overvoltage is the excess voltage required to drive the reaction under real circumstances (that is, over and above the theoretical voltage). Oxygen overvoltages can be as much as 1 volt, depending on the anode composition. There are also hydrogen overvoltages, chlorine overvoltages, etc.; in any case, they depend greatly on the electrode material.

Overvoltage is one more reason why much of electrolysis involves experience with what works and what doesn't. Don't mistake this for random trial-and-error; search the literature instead.

Textbook reduction potentials are given in terms of variation from the standard hydrogen potential, which itself represents an arbitarily-assigned value (i.e., 0.00000; electrochemistry's answer to the Prime Meridian). One cannot expect these values to translate readily to actual applied voltages.

3. Logistics

Suppose we form PbO₂ in an electrolytic cell. The dissociation of PbO₂ back into Pb²⁺ and water is much more energetically favorable than the formation of the lead dioxide. If this compound could find its way back to the cathode, reduction would indeed occur; our PbO₂ would cease to exist. This doesn't happen, and the reason is simple: PbO₂ is not all that soluble; it forms a heavy precipitate that either clings to the anode or falls to the bottom of the cell. Physically, it never makes it back over to the cathode.

Consider another example. Let's imagine we form hydrogen from an aqueous solution via the half-reaction:
2H⁺ + 2e^- <---> H₂,
which is actually the Standard Hydrogen Electrode half-reaction (E⁰ = 0.00000 volts).

Regardless of its oxidation or reduction potential, this one is no different from any other electrolytic half-reaction; in other words, it could go in either direction (depending on applied voltage). The reason this doesn't happen in your electrolysis cell, however, is that H₂ is a gas with poor solubility in water. As soon as it forms, it vacates the cell. The equilibrium essentially becomes a one-way reaction.

Suppose, on the other hand, that we made sodium persulfate in an electrolytic cell. S₂O₈^2- ions don't form an insoluble precipitate with anything in our cell. They don't form a gas that escapes the cell. As a result, they stay in solution and can meander over to the cathode where they are not going to last very long (reduction potential of persulfate ion is around +2.0 volts relative to the H electrode, slightly higher in acidic conditions). Therefore, we'd better use a salt bridge or some other way to keep S₂O₈^2- away from the cathode.

Another example where "logistics" come into play is the reduction of Fe³⁺ to Fe²⁺ in an electrolytic cell. Supposing the voltage and current combination we're using are working smoothly, we still have to remember that atmospheric oxygen will oxidize Fe²⁺ back to Fe³⁺. This will happen most readily at the interface of electrolyte and air. Obviously, you don't want to cap a cell that produces gas bubbles, since pressure buildup could burst it. If the gases form an explosive mixture, the cap could make them build up to dangerous levels. However, a loose-fitting cover can greatly reduce the amount of atmospheric oxygen that can act on the cell.

4. Proven Systems

By now there are many well-characterized systems. It is possible to run them with some set of parameters to yield known products in known proportion. Electrolysis of water, mentioned before, is probably the simplest example. There are also others, especially ones that are or have been used for industrial processes. One of the best-known (also mentioned previously) is electrolytic production of chlorine bleach; indeed, it seems there are few general chem textbooks that don't give it at least passing mention.

Back to Part I of Electrolysis Experiments

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