ELC_Electrolysis Lesson

Electrolysis

We have learned how galvanic cells use spontaneous reactions to produce electricity. Now, let's flip this around! If we pass electricity through a non-spontaneous reaction, we can force that reaction to occur.  This process is called electrolysis.    

Click below to watch this short overview of electrolysis.

Here is an example of the electrolysis of NaCl. This will force the NaCl to decompose into Na and Cl₂..

The substance undergoing electrolysis, NaCl in this example, is placed in a container. An external heat source is used to melt the NaCl. So, the container is full of molten Na+ and Cl- ions.

Two electrodes, one anode and one cathode, are placed in the container. Sound familiar? We use these terms in galvanic cells as well.

The electrodes are connected by a wire. In a galvanic cell, something can be placed between the wire that connects the anode and cathode that uses the electricity (because a galvanic cell produces electricity). But in electrolysis, we need a source of electricity, like a battery.

This overall set-up is called an electrolysis or electrolytic cell.

The battery "pumps" the electrons by pulling them away from one electrode and pushing them toward the other electrode. This gives one electrode a positive charge and the other a negative charge. The Na+ ions are attracted towards the negative electrode where they undergo this reaction, Na⁺(l) + e^- → Na(l). This reaction shows a gain of electrons, and is the reduction reaction.

The Cl- ions are attracted towards the positive electrode, where they undergo this reaction, 2Cl^-(l) → Cl₂(g) + 2e^-. You can see the chlorine bubbles above. This reaction shows a loss of electrons, and is the oxidation reaction.

So far, the terminology and conventions have been the same as they were for galvanic cells. But, let's look more carefully. At the negative electrode above, the sodium ions were reduced. In a galvanic cell, the negative electrode (the anode) is where oxidation occurred. Do you see the contradiction? In order to maintain the convention that reduction always occurs at the cathode (remember red cat) and oxidation always occurs at the anode, the charges for anode and cathode must be switched for electrolytic cells. You just need to remember that in electrolysis, a reaction that would not happen on its own is forced to happen. We are also "forced" to switch the sign of anode to + and cathode to -. So, the electrodes in electrolysis follow a different sign convention that we chemists use. (Labeling an electrode as positive or negative is beyond the scope of this course and the AP Exam, but I wanted you to be aware of this in case you continue studying chemistry or physics in college!)

Stoichiometry of Electrolysis

Michael Faraday, a British chemist and physicist, discovered that the amount of chemical change that occurs during electrolysis is directly proportional to the amount of electric charge passed through it. This means that we can use stoichiometry to calculate the amount of product formed! It will also be helpful to know some other basic units used in electricity.

Remember to work on the module practice problems as you complete each section of content.

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