cum
 

CSEC Chemistry: Extraction of Metals (Using Electrolysis)

Updated: Jun 1, 2021

Hey you! Look around you… what do you think you take the most for-granted? Your parents? The roof over your head? Air? Not being paralyzed? The skip intro button on Netflix? Well that’s all very cute, look at you being grateful, here’s a gold star.



But YOU’RE WRONG! The thing you take most for-granted… are metals. That doesn’t make any sense, I hear you say, how is my not acknowledging aluminium or some other… metal… more important than not acknowledging my ability to literally breathe? Well, metals are literally all around you.





Copper wires carry the electricity you can barely function without, cars (are big boxes of metal, usually aluminium or steel), metals of all types create the device you’re reading this on, and metals are used in the industrial processes that produce so many of the other products we use. Hell, even the fractional distillation of crude oil usually happens in a fractionating tower made of metal.


If you’ve been paying attention in Chemistry class (if you haven’t, I’d really suggest you re-evaluate your priorities, or read our post here) you already know about electrolysis, the decomposition of electrolytes when an electric current is run through them. However, just like every topic in Math, you probably thought electrolysis was some kind of random thing some random guy discovered could be done and had no purpose other than showing up on your syllabus and adding an extra 3 pages to your exam. Well, while that “what use does this have” ideology generally does apply for math, electrolysis actually plays a big role in getting those metals you take so for-granted into a form that can be used.


You see, metals don’t just exist in nature in their pure forms- the earth’s crust isn’t some kind of grocery store we can go to and pick up whatever pure metal we want. Instead, metals occur naturally as ores- natural metallic compounds from which we can extract the metals we want. Ores have to contain enough of the metal to make it economically viable to extract the metals within as well, so we don’t waste massive quantities of power for a disappointing output, like a report card. In this post, we’ll be talking about the extraction of metals from these ores, and how exactly electrolysis ties into actual industrial applications.


As we said before, metals occur naturally as ores, these metallic compounds which are usually contaminated by other materials. These ores are ionic compounds that can be oxides, carbonates, chlorides or even sulphides of the metals we want. And since these are ionic compounds (hinting back to some fourth form knowledge on types of compounds there), the metals exist as positive ions, or cations, bonded strongly via electrostatic attraction to one or more anions.

So, extracting metals means converting metal cations to metal atoms with the help of reduction, the gain of electrons. When metal cations gain electrons, they lose their positive charges and are reduced to pure metal atoms that we humans can use to make all sorts of wacky inventions.


Alright, well, how do we get from some impure deposit of ore in the earth, to the clean magnesium casing of the all-new Surface Book 3 from Microsoft in Platinum grey and 13.5 and 15-inch variations?



Well, we start by mining the ore from deposits in the earth’s crust. We no longer do this by hand- mostly, that is.


The next step is to purify the ore. This can be done in several ways, but one of the simplest is the froth flotation method. The ore, like my motivation in online school, is ground into a fine powder and then mixed with water and a frothing agent like pine oil. Then a stream of air is blown through the mixture, causing it to bubble and froth. While the frothing occurs, impurities such as sand and rock are wetted by the water and sink to the bottom of the container. The metal ore does not adsorb water but does adsorb the pine oil.

(The metal powder ‘adsorbing’ the pine oil just means that it holds the oil as a thin film on the outsides of the granules of metal ore.)

The oil-coated ore floats to the top of the mixture, where it can be skimmed off. The next step is the most interesting part- the reduction of the ore to the metal.



We mentioned before that electrolysis played a role in extracting metals, but that is only partially true. Not all metal ores are generally reduced using electrolysis, so other types of metals use chemical reduction or their own special methods of extraction. The method of extraction used relies on two things- the nature of the ore and the reactivity of the metal we hope to extract. The more reactive metals, that is, the metals aluminium and above on the reactivity series, can be extracted using electrolysis of the molten chloride or oxide of the particular metal. The metals zinc, iron and tin can be extracted using chemical reduction. These less reactive metals can be reduced using the reducing agents carbon, carbon monoxide or hydrogen. Metals below tin, like lead and copper, require special methods of extraction due to their lower reactivity.


We’ll be focusing on electrolytic reduction, which is the most powerful method of reduction. These very reactive metals require a lot more energy to reduce than other metals, which is why we have to pull out the 'big guns' with electrolysis to extract them properly.


What metal do you think is the most abundant? You might think iron, just because we produce more of it than any other metal. But no, no- the most abundant metal is aluminium. Aluminium makes up about 8 percent of the earth’s crust, and ranks fairly high on the reactivity series. Because it is so reactive, aluminium requires the firm but gentle touch of electrolysis to coax it from its ores.


Aluminium, like most metals, is never found in its free state. Only very unreactive metals like gold and platinum hate mingling with other elements enough to be found pure in nature. Metals like those are called native metals.



Assuming that you’ve been alive for at least more than a year, you probably already know that bauxite is the main ore of aluminium, and is the raw material of one of the chief exports of countries like Jamaica. Bauxite is mostly made up of hydrated aluminium oxide (Al2O3.xH2O), but we can’t just throw a couple tonnes of red dirt into a vat and run a current through it and expect aluminium. No, my dear Watson, we (and by we, I mean a bauxite processing plant) must first convert the crude bauxite that we ripped from acres of now barren countryside into either calcined bauxite by heating it at 3000 degrees, or to pure alumina, anhydrous aluminium oxide. While the calcined bauxite is mostly used in roadways for ‘making roads safer and more durable,' the really cool stuff happens with alumina.


The alumina melts at 2050 degrees Celsius, but most companies don’t want to use that much power to keep it melted during electrolysis. So, the molten alumina is dissolved in molten cryolite (also called sodium aluminium fluoride) to reduce the melting point to 950 degrees. Wow! It’s almost cool enough to take a shower in. Seems kind of counter-intuitive, right? Purifying the alumina only to impurify it by mixing it with another compound? Nonetheless, it benefits the process by making the alumina more conductive, reducing the amount of power needed to keep it molten during electrolysis, and making it so that the mixture doesn’t mix with the aluminium produced via electrolysis.





Now, for the fun part- shocking a vat of molten aluminium oxide and cryolite to get some molten aluminium. While the apparatus of the electrolytic cell you’re used to looks like this:



in the industrial electrolysis of alumina (sometimes called the Hall-Heroult Process), the apparatus looks like this:


Our electrolyte is a nice and steamy 950-degree mixture of molten alumina and cryolite. That electrolyte is contained in a durable steel tank, with an opening at the bottom for us to extract the sweet nectar of molten aluminium afterwards. The cathode, the negatively charged electrode, is made of graphite and lines the tank filled with the electrolyte. The anode (positively charged) is also made of graphite or some titanium alloy. Unlike a little experiment with electroplating, this electrolysis requires massive amounts of electricity, around 5 Volts and 100000 Amperes. When we switch on this terrifying circuit, two main reactions occur, one at the cathode, and another at the anode.




At the cathode, aluminium ions from the aluminium oxide in alumina gain three electrons to be reduced to aluminium atoms. This can be shown in the half equation Al3+(l) + 3e-→Al(l). The aluminium ion is quite greedy, as one mole of ions needs 3 moles of electrons to be discharged. At the anode, on the other hand, oxidation occurs, where each oxide ion is discharged to an oxygen atom by losing two electrons to the positively charged anode. Two oxygen atoms will pair up to form molecular oxygen gas, or O2. This can be shown in the half equation: 2O2-(l) →O2(g) + 4e-.


The molten aluminium can be siphoned off through that little opening in the tank, giving us the product of all of our electricity and labour. Of course, Jamaica exports alumina rather than aluminium because our electricity is not even near cheap enough to make it economically worthwhile.


(Note: the graphite anode used in the industrial electrolysis of aluminium oxide needs to be periodically replaced, because over time, the carbon from the graphite can react with the oxygen liberated to form carbon dioxide and slowly degrade)


Other metals follow essentially the same process with extraction via electrolysis. The metal cation from the molten ore will be discharged at the cathode, gaining electrons and giving the molten pure metal. The anion from the ore, however, will be discharged at the anode, losing electrons.


In the end, electrolysis has many industrial applications, but the most useful application by far is the electrolytic reduction of metal ores. Now you know the ‘charge’ of all the metals you enjoy.