PlatinumDrakes
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what happens during the electrolysis of potassium aluminate and koh

Postby PlatinumDrakes » Tue Jun 09, 2020 5:37 pm

So using nickel electrodes what happens in the solution of koh and potassium aluminate? The voltage would be around 6 volts and a 30% koh 10% potassium aluminate and 60% water (weight).

LouisCouryBiogen
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Re: what happens during the electrolysis of potassium aluminate and koh

Postby LouisCouryBiogen » Sun Oct 11, 2020 4:30 pm

This is a very good question. You have chosen a fairly complex system to study, but let's walk through the possibilities of what could be happening.

In general, an electrolysis cell (2 electrodes, electrolyte solution, power source) causes an oxidation reaction to occur at one electrode (the anode) and a reduction reaction at the opposite electrode (the cathode).

The most common system that is described in text books is the electrolysis of water. There are many references on the web that explain the details of water electrolysis. Here is one example:

https://chem.libretexts.org/Bookshelves/Introductory_Chemistry/Book%3A_Introductory_Chemistry_(CK-12)/23%3A_Electrochemistry/23.09%3A_Electrolysis_of_Water

As this link explains, electrolysis of pure water (no added salts) does not occur very rapidly. So, a small amount of a salt that will dissolve in water, called an electrolyte, should be added so that the solution will conduct electricity. The movement of ions allows current to pass through the solution, completing the electrical circuit, and allowing the oxidation and reduction reactions to occur. In the example described in the link above, they added a small amount of sulfuric acid to the water in the electrolysis cell. Sulfuric acid is a liquid that dissociates into hydrogen ions (H+) and sulfate anions (SO4)2-, and these ions conduct electricity by moving around in the water.

In the electrolysis reaction in the link, the reaction that occurs at the anode is oxidation of water to oxygen gas (O2) and hydrogen ions. The reaction that occurs at the cathode is the reduction of water to form hydrogen gas (H2) and hydroxyl ions (OH-). The sulfate ions that are in solution do not participate in the electrolysis reactions at either electrode, because sulfate is very difficult to either oxidize or reduce.

So, let's look at your set-up. Since you are using water as the solvent (main component), the same reactions could occur in your electrolysis cell. In your case, you have potassium aluminate and potassium hydroxide dissolved in the water in your cell, so you have the required electrolyte.

But there is one important difference between your electrolysis cell and the one described in the link. They use platinum electrodes (Pt), which are considered inert in water. That is, you cannot apply a sufficiently large voltage across two Pt electrodes in water to oxidize the platinum metal into platinum ions, without first electrolyzing all of the water present. In your case, you have used nickel electrodes (Ni). Nickel is not as inert as platinum, and Nickel metal can be oxidized to Nickel ions (most easily to the +2 oxidation state).

So how do you know what is happening in your cell? I will give you hint that Aluminum is extremely difficult to oxidize or reduce, so the aluminate is unlikely to be involved in any reactions, given the relatively low voltage you are applying. That is, aluminate, [Al(OH)4]- will not be oxidized or reduced. The same is true for potassium ions (K+). No change should occur to those ions either. The reason I know this is by looking at the standard oxidation/reduction potentials for those ions in a reference book. (Working with standard oxidation/reduction potentials may be little bit more advanced than you are ready for at this stage - no way for me to know. So, I am helping you out by telling you this is true.)

So, that eliminates everything present in your case except water and the Nickel anode electrode. If water is easier to oxidize than Nickel, then the same reactions should occur as at the Pt electrodes in the example above. If Nickel is easier to oxidize than water, then the Nickel electrode itself will undergo the reaction at the anode, and will dissolve away as Ni(2+) ions are formed from Nickel metal.

What do you observe when you operate your electrolysis cell for awhile? Do you see oxygen bubbles form on the anode, or do you see the electrode dissolve away (undergo pitting, cracking and other physical changes that you can see)? Note that if Nickel and water are about equally easy to oxidize, you could actually see both things occur (pitting of the anode as well as formation of oxygen bubbles).

By the way, the only reaction that could be occurring at the Nickel cathode is reduction of hydrogen ions (or water) to form hydrogen gas (H2).

I recognize that this has been a long answer, but since your electrolysis system is fairly complicated, it takes awhile to explain everything. Feel free to post follow-up questions, and I will try to answer them for you. Good luck!


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