CR Scientific: Electrolysis Experiments

Anions move toward the anode;

Cations move toward the cathode.

Simple, yes?

What about the charges...?

Memory tip: Just remember that "cathode rays" (as in a TV or computer monitor) are in fact electrons, which of course carry a negative charge; therefore, something that's attracted to a cathode would have to be positively charged. Hence, "cations" carry a net positive charge.

This is all quite simple for profs to remember, but students tend to get confused on the "anode-cathode" convention.

More Mnemonics:

Some people like to remember "REDCAT", as in "Reduction occurs at the Cathode". Electrical wires are color-coded black and red, but unfortunately the cathode is actually the black one, not the red one (even though RED in this mnemonic doesn't mean the color code).

To avoid that confusion, there's another mnemonic. "OXAN", or "Oxidation occurs at the Anode". The terms with vowels go together.

"OXAN, you don't have to wear that dress tonight."

You could flip the mnemonic around to make ANOX if you really wanted, as in "anoxia"... This might be easier to remember for some.

Electrolysis Experiments

Part I: Introduction - How Electrolysis Works

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The immense variety of reactions that one can perform via electrolysis makes it a useful tool not only for classroom demonstrations, but also in countless types of industry.

The Basics of Electrolysis

When a salt is dissolved in water it yields ions. For example, sodium chloride (NaCl) dissolves in water to give a solution of Na⁺ and Cl^- ions. The presence of ions in solution is what allows water to conduct electricity. Pure water is, in fact, not a conductor; however, unless the water has been specifically de-gassed and de-ionized it will contain enough dissolved ions to make it conductive. (Even after deionization, water can become conductive again as it dissolves CO₂ from the atmosphere).

By passing an electrical current through a solution that contains dissolved ions, we can make the ions participate in chemical reactions. Depending on the ions, the voltage, and other factors, this can sometimes allow the formation of new compounds. Dissolved sodium chloride, for example, can be converted to an entirely different compound, sodium hypochlorite (bleach).

In electrolysis, the anode is the positive electrode, meaning it has a deficit of electrons; species in contact with the anode can be stripped of electrons (i.e., they are oxidized). The cathode is the negative electrode, meaning it has a surplus of electrons. Species in contact with the cathode tend to gain electrons (i.e., they are reduced).*

A higher current flow (amperage) through the cell means it will be passing more electrons through it at any given time. This means a faster rate of reduction at the cathode and a faster rate of oxidation at the anode. This corresponds to a greater number of moles of product. The amount of current that passes depends on the conductance of the electrodes and electrolyte, though it also depends on how much current the power source itself can generate.

Current also makes a difference in that it can shift chemical equilibria by sheer mass action*. The processes in an electrolytic cell with just two or three reactants can become very, very complex. Most of the time it's best to search the literature to see what current density works best for a desired process. For instance, metals plated at a certain current density might form a durable and shiny coating on the substrate, while some other current density might form an excessively grainy, dull coating. In some cases the "plated" metal can even fall away from the electrode as mossy clumps.

A higher potential difference (voltage) applied to the cell means the cathode will have more energy to bring about reduction, and the anode will have more energy to bring about oxidation. Higher potential difference enables the electrolytic cell to oxidize or reduce energetically more "difficult" compounds. This can drastically change what products will form in a given experiment. On a practical level, both current and voltage determine what will form in a cell.

So, what happens in an electrochemical cell? In practice, things aren't as simple as one might plan. There can be side reactions. These can become intractably complex unless one has studied electrochemistry for many years. In many cases it becomes a matter of experience -- or exploring the existing literature. Industrial electroplaters, for example, generally know what conditions work best for a given type of plating: current, voltage, electrode material, electrode shape, and so forth.

When determining what will happen when current passes through the electrolytic cell, five of the most important variables are:

1. Electrode composition

2. Electrolyte composition (including pH, ionic strength, etc.)

3. Voltage and Current levels

4. Temperature of the system

5. Partition, if any (i.e., do the anode and cathode solutions mix freely, or are they separated by a membrane or salt bridge?)

If the electrolyte contains chemical species that will be reduced at the cathode or oxidized at the anode (or both), that means chemical change. The electrolyte can form one or more compounds or ionic species that weren't there before.

A given species could form at one electrode but diffuse back over to the other electrode where it promptly breaks back up into its reactants.

If the electrolyte does not participate in any reaction(s), it will just act as a conductor of electricity. The same goes for the electrodes. Some materials will react but will not produce anything obvious (gases, precipitates, or color changes). Others will not react at all under the specific combinations of voltage and current.

With appropriate metals as the electrodes, electrolysis brings about a very useful process:
The anode is consumed as its metal is oxidized, turning the metal into positive ions which go into solution.
The cathode is plated with freshly-reduced metal ions that come out of solution.

Some metals, such as platinum, don't usually participate in electrolysis reactions, at least not directly. Among other things, this means that the anode will not disintegrate over time, which in turn means the set of electrodes can last years (just don't put them in aqua regia!).

Safety Considerations

Electricity and water don't mix, as the saying goes. We all know it's asking for trouble to go swimming during a lightning storm, for example.

Electrolysis involves water and electricity, but it takes place under carefully-controlled circumstances. These include making sure the current cannot rise to dangerous levels. The first and most important caution is this: DO NOT use "wall current", "house current", or any type of mains current for electrolysis under any circumstances. Attempting to do so could kill you. Remember, electrolysis depends on controlling the voltage and current.

Even if it were current-limited, alternating current (AC) would not be appropriate for electrolysis. Because the "cathode" and "anode" are constantly switching places, AC produces explosive mixtures of hydrogen and oxygen. Many instructors like to perform combustion experiments with the H₂ and O₂ in their respective tubes, but a mixture of both gases isn't something to ignite in a glass container...

Direct current (DC) is used for electrolysis. A laboratory DC power supply provides built-in voltage control and current limiting. This kind of power supply is the recommended means of powering your experiments.

Electrolysis can also be performed with small batteries (e.g., a 9V alkaline battery or a few 1.5V "D" cells in series). Batteries produce DC; small ones can be thought of as current-limited. A small, alkaline-type battery simply can't put out that much current. Even though it's possible, for example, to build a device that steps a flashlight battery's voltage up into the thousands, the current becomes extremely small, even smaller than it would have been at 1.5, 6.0, or 9.0 volts. However, be very careful with large storage batteries such as car batteries; these house a great deal of power. Understand that current, not voltage, is the primary danger of electricity.

Though it is possible to conduct electrolysis using a car battery without incident, one must have a thorough understanding of what could go wrong and how to avoid it. There is no point in having uncontrolled current, even if everything goes right and one avoids fire or other catastrophe.

Let's start with the simple equation V = IR. Current in amperes (I) therefore equals voltage (V) divided by resistance in ohms (R). 6 volts flowing across a resistance of 2 ohms would give a current of 3 amperes. A strong solution of electrolyte will have a very low R, nowhere near a full ohm and usually much less than 1/10 of an ohm. Thus, the value of I will be very high unless one introduces some more resistance to the circuit.

Let's suppose one wanted to limit the current to no more than 3 amperes. This is a realistic value, given that many laboratory power supplies have 3 amps as their maximum output.

Neglecting the resistance of the electrolyte solution and striving for a setup that allows no more than 3 amps to flow with a 12 volt storage battery, we resort once again to V=IR. Since V = 12 volts and I = 3 amps, that means R will have to equal 4 ohms. Recall, though, that a current flowing across a resistance can generate quite a bit of heat. A power supply based on a storage battery will need a sand-block or ceramic resistor rated for the amount of heat involved. Here we use the equation P=I²R ("twinkle, twinkle, little star... power equals I squared R"). P is power in watts; I is current in amps; R is resistance in ohms. If 3 amps were flowing across a resistance of 4 ohms, the power would be 36 watts. Those little sandblock resistors available at the local electronics retailer are usually rated for only 10 watts. Ordinary resistors can handle even less! Some electronics stores have surplus resistors made to handle many tens of watts; these can be several inches long and half an inch in diameter.

Our final recommendation is... forget the car battery; invest in a laboratory power supply.

We've talked about the electrical aspects of safety for electrolysis, but one must also mind the chemical aspects. As long as you follow a well-established procedure given by your lab instructor or textbook, there is no cause for alarm.

DON'T subject ammonium chloride or ammonium perchlorate solutions to electrolysis. Avoid electrolysis on any solution where aqueous Cl₂ or Cl^- might be present with aqueous NH₃ or NH₄⁺, unless you know PRECISELY what you are doing. That means NO ammonia when there is any type of chlorine or chloride present. There are a couple different things that could form and cause an explosion, poison you, or both.

There are other combinations to avoid as well; it is entirely up to the reader to research any possible dangers ahead of time.

Electrolysis of Water

To accomplish this, one must make the water conductive by using an electrolyte that does not react under the conditions of the experiment. Sodium sulfate (Na₂SO₄) or dilute sulfuric acid (H₂SO₄) can be used; keep the concentrations low, especially when using sulfuric acid. Magnesium sulfate ("Epsom salt", MgSO₄) is also suitable.

Electrolysis of water will begin around a minimum of 1.2 volts and will increase in rate as the voltage is increased. Typically, the electrolysis is carried out around 6 volts. Remember that higher voltage (V) across a given resistance (R) equates to lower current (I). That may seem counter-intuitive, but V=IR really works.

The anode is where oxidation takes place. Oxidation involves the loss of electrons. The cathode is where reduction takes place; reduction involves gain of electrons. Now, which gas will be produced at which electrode? If we think of water as hydrogen oxide (which it is), we can consider each hydrogen portion of water as H⁺. These would have to gain electrons to become neutral hydrogen; therefore, this must happen at the cathode. We can consider the oxygen part of water as O^2- (the oxide ion, which always has a -2 charge); this would have to lose electrons to become neutral. Thus, oxygen gas should form at the anode.

Finally, remember that there are three components that could theoretically participate in the electrolysis: the water, the dissolved electrolyte, and the electrodes. In this experiment our choice of electrodes (platinum) and electrolyte (sodium sulfate) has ensured that these don't participate in the chemical reaction. Remember the five variables we listed in the beginning of this article.

Electrolysis of Sodium Chloride Solution

Cautious electrolysis of NaCl solution with the Brownlee apparatus will produce hydrogen plus aqueous NaOCl if the experiment is carried out in a single, unpartitioned jar with stirring. What happens first is that hydrogen, chlorine, and NaOH are produced, and the aqueous chlorine and NaOH then react to form NaOCl ("bleach"). Higher temperature, however, tends to produce ClO₃^- (aq.) instead of OCl^- (aq.). Running the electrodes in two separate partitions will produce gaseous chlorine (poisonous!) and hydrogen, plus aqueous NaOH.

Some chlorine will undoubtedly escape the hypochlorite cell and go into the air without being consumed, so this whole operation should be done in a fume hood or outdoors. The odor of chlorine is sharp enough that, if it escapes into the room, you should have sufficient warning to shut off the power supply and get out of the area before there's any injury.

Another, similar experiment is the electrolysis of a concentrated NaBr or KBr solution. This will produce hydrogen at the cathode and bromine at the anode. The latter is corrosive and poisonous, like chlorine. (This experiment must be performed in a fume hood and with an emergency shower nearby in case of accidents.) Bromine in any appreciable amount will also begin to attack the platinum electrodes (!), so a modified apparatus using graphite electrodes would have to be used.

Electrolysis can force many chemical reactions to go against their "normal" (spontaneous) direction; this will occur when a certain electric potential (i.e., applied voltage) is reached. For example, an electric current can cause Pb⁺⁺ ions in water to form PbO₂ (solid) and H⁺ (aq.), a reaction which normally proceeds the other way.

Here we'll mention a basic caution once more: the Brownlee electrolysis apparatus is designed for Direct Current (DC) only! DO NOT attempt to use AC (alternating current) for electrolysis.

Onward to Part II of Electrolysis Experiments

Notes

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Split water into hydrogen and oxygen. Perform this classic electrolysis experiment in your classroom!

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Voltage and Energy as it applies to electrolysis.....

Voltage expresses the energy per unit of charge.

12 Volts equals 12 Joules per Coulomb.

6 Volts is only 6 Joules per Coulomb.

Power, as in P = I²R, would appear to depend strictly on current flow across a resistance, with no mention of voltage.

However, the current I that is able to flow across a given resistance R is equal to V / R (because V = IR). Thus, voltage still plays a role. In fact, power can be written as P = VI.

If you put 12 volts across a 1 ohm resistor, 12 amps will flow. It will dissipate 144 watts of heat.

If you put 6 volts across 1 ohm, 6 amps will flow. It will dissipate 36 watts of heat.

The confusion arises because power sources and batteries cannot put out infinite power. In real life, a 9-volt transistor battery houses less power than a 6-volt storage battery.

If you put a 9V transistor battery across a 1 ohm resistor, the current would be 9 Amperes for only a short time; it would fall off very rapidy.

Tiny batteries simply cannot sustain large amounts of power.

A laboratory power supply usually accepts "mains" voltage (110/120) as the input, but the voltage is stepped down (typically to 12 V) and current limited (usually no more than 5 A) for safety.