You have probably seen the term bauxite or aluminum ore before. Aluminum is widely used in industry, such as in the making of planes. To be able to extract aluminum for aluminum ore, industries need to perform the electrolysis of the aluminum ore. What does this mean? Let's find out!
- First, we will talk about electrochemical cells.
- Then, we will look at the definition of electrolysis.
- After, we will look at the electrolysis of water and how electrolysis is used to remove rust.
- Lastly, we will look at the electrolysis of lead bromide.
Electrochemical Cells
Before diving into the world of electrolysis, let's recall the basics of electrochemical cells.
Electrochemical cells are systems that integrate
redox reactions to produce, or use, electrical energy.
Electrochemical cells can be galvanic (or voltaic) cells or electrolytic cells.
Galvanic cells are spontaneous. Spontaneous processes are those that are thermodynamically favored to happen without the input of energy from the external surroundings. During a spontaneous process, the change in Gibbs Free Energy is negative (ΔG < 0) and the cell potential (Ecell) surpasses zero, Ecell > 0. Galvanic cells produce energy in the form of electricity.
Cell potential (Ecell) is a measurement of the potential difference between two electrodes. Cell potential can also be called the voltage of the cell, or electromotive force (emf).
Electrolytic cells, on the other hand, involve non-spontaneous processes. Here, the change in Gibbs free energy is positive (ΔG > 0) since the electrolytic cell requires the input of free energy from the surroundings. These cells also have a negative cell potential, Ecell < 0. Electrolytic cells consume energy, so they need a power source to force a non-spontaneous reaction to happen!
Let's make a table to summarize the relationship between ΔG, Ecell and Spontaneous/Non-spontaneous reactions.
| | Spontaneous or Non- Spontaneous? |
ΔG < 0 | Ecell > 0 | Spontaneous |
ΔG > 0 | Ecell < 0 | Non-Spontaneous |
If you're not sure about what this means, check out "Gibbs Free Energy"!
Anodes and Cathodes
Electrochemical cells have two electrodes: anode and cathode. The anode is the site of oxidation, whereas the cathode is where reduction takes place. Electrons always tend to flow through the wire from the anode to the cathode. In other words, electrons transfer from where they are lost to where they are gained.
Oxidation is the loss of electrons. Reduction is the gain of electrons.
In galvanic (voltaic) cells, the anode is negative, and the cathode is positive. Since the electrons are negatively charged, they want to move away from the negative electrode.
In electrolytic cells, it's the other way around: the anode is positive, and the cathode is negative! This makes sense because we are adding external energy to the compartment containing the cathode electrode and, in reality, the electrons would prefer to travel in the opposite direction (away from the added negative electrons). In the case of an electrolytic cell, we are forcing a non-spontaneous reaction to occur, by adding energy from the external surroundings in the form of electric current!
Another important thing to know about anodes and cathodes is that, for metal/metal salt voltaic cells, the cathode gains mass whereas the anode loses mass!
Electrolysis Definition
Now that we know more about voltaic and electrolytic cells, let's talk about electrolysis, which is a process that happens in electrolytic cells, typically to produce certain elements!
Electrolysis is referred to as the process in which electrical energy is used to cause a non-spontaneous chemical reaction.
There are two types of electrolysis that you should be familiar with: molten and aqueous electrolysis.
Molten Electrolysis usually only occurs at very high temperatures due to the high melting points of ionic compounds. Let's use sodium bromide, NaBr, as an example. In this case, Na has a +1 charge, while Br has a -1 charge. So, we can use electrolysis to turn Na+1 into elemental Na, and Br-1 into elemental Br.
When dealing with electrolysis, we can use half equations to describe what is happening that the electrodes!
Aqueous Electrolysis involves water, as expected. With aqueous electrolysis, it is a little harder to predict the products because water can also oxidize and reduce. So, we need to figure out which reaction will happen at the cathode and at the anode. (Tip: only one reaction can happen at each electrode!)
Let's look at a similar example, now involving the aqueous salt, NaBr (aq). In this example, we are given the reduction potentials of some half-reactions.
At the cathode, we could have:
At the anode, we could have:
The general rule is that the reaction with the least negative reduction potential will be the one occurring at the cathode. In this case, the reduction potential of the sodium reaction is -2.71 V, and the reduction potential of the reaction involving water is -0.83. So, at the cathode, the only half-reaction that will occur is the one involving the reduction of water!
At the anode, oxidation happens, so the values for reduction potential need to be inverted! In other words, If the reduction potential of the bromine reaction is + 1.07 V, then the oxidation potential will be -1.07 V. Similarly, the reaction with the least negative oxidation potential will be the one happening at the anode. So, at the anode, the only half-reaction that will occur is the one involving the oxidation of the bromine ion!
If you calculate the overall cell potential (E°cell) in this example, it would be, E°cell = -0.83 + -1.07 = -1.90 V. Since this value is negative, then we can tell precisely that it is a non-spontaneous reaction! So, this means that the power source in order for this electrolysis to occur needs to be able to supply at least 1.90 volts.
Electrolysis of Water
Hydrogen and oxygen combine spontaneously to make water. Electrolysis can be used to initially break H2O molecules into hydrogen and hydroxide ions, eventually producing hydrogen gas (H₂) and oxygen gas (O2).
In this case, water is the electrolyte. When the two electrodes are connected to a power source such as a battery by using waterproof wiring, the positive hydrogen ions will be in attraction with the negative cathode and hydrogen gas will form at the negative cathode. Negative ions will be attracted to the anode. So, oxygen gas will form at the positive anode and be collected in a test tube.
Electrolysis Rust Removal
We can also use electrolysis to remove the rust from metal. This works by adding water and sodium carbonate (or water and sodium hydroxide) into a container or beaker and connecting waterproof wires to the anode (usually a piece of iron or steel) and to the metal piece covered in rust.
When the power is turned on, hydrogen and gas bubbles will be formed, letting you know that electrolysis is working! Depending on the rust, you might have to leave it for hours, or even weeks! When you start seeing that the rust turned into a black oxide, remove the metal from the water and scrub until the black oxide comes off!
Hydrogen Electrolysis
Hydrogen is the most abundant element in the universe and the third most abundant element on earth. Hydrogen is widely used in chemical and industrial fields. So, hydrogen needs to be produced on a large scale!
Hydrogen can be synthesized in different ways, such as through the electrolysis of water (as shown above), the steam reformation of hydrocarbons, and the water-gas shift reaction.
Steam reformation of hydrocarbons involves heating up a hydrocarbon (most commonly methanol or methane) with water in the presence of an iron or copper catalyst. At high temperatures, the steam mixture undergoes a chemical reaction, producing syngas, which is a mixture containing H₂, CO, and CO2.
The water-gas shift reaction uses the carbon monoxide (CO) byproduct from the steam reformation process. This leads to an increase in the concentration of H₂ gas in the system.
Electrolysis in Chemistry
Did you know that we can also use electrolysis to extract metals? Let's look at the separation of lead bromide into lead and bromine. For this electrolysis to occur, we need to use a Bunsen burner to melt the lead bromide so that the ions in the compound can move freely.
In this process, the lead ions (Pb2+) will get attracted to the negative cathode, while the Br- ions will be attracted to the positive anode. So, lead metal will be produced at the cathode, and bromine gas will be produced at the anode.
Chemical changes cause a substance to turn into a chemically different substance. Electrolysis is an example of a chemical change because, in electrolysis, substances gain or lose electrons.
Now, I hope that you feel more confident in your understanding of electrolysis!
Electrolysis - Key takeaways
- Electrolysis is referred to as the process in which electrical energy is used to cause a non-spontaneous chemical reaction.
- Electrolysis happens in electrolytic cells, where the anode is positive, and the cathode is negative.
- During electrolysis, the negatively charged ions get attracted to the positive cathode, whereas the positively charges ions get attracted to the negative cathode.
- Galvanic cells are spontaneous. During a spontaneous process, the change in Gibbs Free Energy is negative (ΔG < 0) and the cell potential (Ecell) surpasses zero, Ecell > 0.
- Galvanic cells produce energy in the form of electricity.
References
- Theodore Lawrence Brown, Eugene, H., Bursten, B. E., Murphy, C. J., Woodward, P. M., Stoltzfus, M. W., & Lufaso, M. W. (2018). Chemistry : the central science (14th ed.). Pearson.
- N Saunders, Kat Day, Iain Brand, Claybourne, A., Scott, G., & Smithsonian Books (Publisher. (2020). Supersimple chemistry : the ultimate bite-size study guide. Dk Publishing.
- Chad’s Videos (n.d.). Chad’s Prep -- DAT, MCAT, OAT & Science Prep. Retrieved June 15, 2022, from https://courses.chadsprep.com/courses/take/general-chemistry-1-and-2/pdfs/217201-complete-general-chemistry-outlines
- Nelson, P. (n.d.). Removing rust with your battery charger (electrolysis method). Retrieved June 15, 2022, from http://fergusontractors.org/nfs/wp-content/uploads/technical-articles/Rust-Removal-with-a-Battery-Charger.pdf
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