Synthesis of alkali metal hydroxides via electrolysis

Abstract:
A long quest for me has finally been completed on whether cotton, often found cheap as cotton wool balls for removing make-up, actually can be used by the home chemist to act as an ion-permeable membrane in electrolysis. The aim is turning water and alkali metal chlorides into hydrogen, chlorine and, most importantly, the respective alkali metal hydroxide. It is often misleading how to actually synthesise alkali metal hydroxides by electrolysis, but with exploring further and problem solving, this process amazing works.

Introduction:
I have been meaning to get hold of some potassium hydroxide for a while now. The main use I had in store for it was in the construction of an electrolysis cell that would split water to produce hydrogen and oxygen. Potassium hydroxide is an excellent electrolyte, allowing electricity to easily flow from electrode to electrode. I usually prefer to make a chemical myself instead of buying it so I can enjoy exploring the science around the synthesis. Also it is sometimes cheaper.

Many sources, including BBC Bitesize and others, seem to say the electrolysis of a concentrated sodium chloride solution would produce sodium hydroxide, hydrogen and chlorine. If potassium chloride was used, this should then produce potassium hydroxide. On the other hand though, some sources said performing this electrolysis would produce almost no sodium hydroxide, and instead mainly produce sodium hypochlorite and chlorate. Quoting myself from one of my previous blog posts, chlorates are illegal and potentially dangerous:

“Also you shouldn’t be interested in the potassium chlorate as it is dangerously oxidising, landing it on the UK regulated substances list as an explosive precursor, requiring an EPP license to have the chemical in your possession.”

(18thtimelucky.wordpress.com/extracting-red-phosphorus-from-match-boxes)

So instead of finding out by experimenting who was right or wrong, and risk making chlorates, I searched a bit longer, stumbling across the idea of using an ion-permeable membrane to connect otherwise separated electrolysis cells – membrane electrolysis. I couldn’t afford any expensive industrial membranes so I took to YouTube in case some of the amateur chemists on the site had a few ideas of making some cheap membranes. I was in luck. A few scattered videos showed a few different salt bridge designs, even using gelatine, but one video that stuck with me was using cotton as a membrane.

I was immediately sceptical as it sounded too good to be true, but it seemed to work so I thought I would give it at least a good attempt. Here are the half-cell reactions and the overall reaction.

At the anode the following reaction occurs:
2Cl(aq) → Cl2 (g) + 2e

At the cathode the following reaction occurs:
2H2O (l) + 2e → H2 (g) + 2OH(aq)

This gives an overall reaction of:
2Cl(aq) + 2H2O (l) → Cl2 (g) + H2 (g) + 2OH (aq)

I will use potassium as an example of an alkali metal. The idea is the ion-permeable membrane connecting the two cells allows the potassium ions from dissolved potassium chloride in the anode cell to cross it into the cathode cell although the hydroxide ions formed at the cathode cannot cross it into the anode cell. The potassium ions then react with the hydroxide ions to form dissolved potassium hydroxide in the cathode cell.

This is useful as it helps in keeping the potassium chloride and potassium hydroxide separate for easier purification. Also, more importantly, it prevents the reaction between chlorine molecules formed at the anode and hydroxide ions formed at the cathode from forming a variety of side products such as chloride ions, hypochlorite ions, chlorate ions, and water.

An example of a reaction avoided by using membrane electrolysis rather than single cell electrolysis:
Cl2 (g) + 2OH(aq) → Cl (aq) + ClO (aq) + H2O (l)

I will be very happy if this works, although I will not be bothered if it does not work as that is what I am expecting; hopefully I am surprised!

Safety:

Substance Hazard information
Potassium chloride Low hazard
Sodium chloride Low hazard
Potassium hydroxide Corrosive
It causes severe burns; it is particularly dangerous to the eyes.
Sodium hydroxide Corrosive
It causes severe burns; it is particularly dangerous to the eyes.
Chlorine Toxic
This is toxic if breathed in, causing severe lung damage. It irritates the eyes, skin and respiratory system. It may trigger an asthma attack and the effects of exposure may be delayed for some hours. For a 15-minute exposure, the concentration of the gas in the atmosphere should not exceed 1.5 mg m-3. It is very toxic to the aquatic environment and so is used to kill microbes in public water supplies.
Hydrogen Extremely flammable
It forms explosive mixtures with air and oxygen. Mixtures with air between 4% and 74% hydrogen by volume are explosive. Explosive mixtures will ignite below 500°C and well below this temperature in the presence of catalysts such as transition metals and their oxides.
The explosion with oxygen produces a very loud noise which can damage hearing.
  • Gloves and goggles should always be worn when handling chemicals to protect your sensitive eyes and skin.
  • Good ventilation is extremely important due to the hydrogen and chlorine produced. You must be outside or using a fume hood.
  • Avoid sparks and naked flames which could potentially cause a hydrogen explosion. Use of insulated wires could help prevent sparks from the wires.
  • Hydrogen and chlorine are both dangerous, chlorine much more so, but the production of these gases is dependent on the speed of the electrolysis. Regularly check on your apparatus while it is running to ensure the production of the gases is always under control. If they are being produced too fast, a lower voltage and current power supply should be used. If you purposely want the reaction to go fast, you should consider building a chlorine scrubber to remove the chlorine as it is produced.
  • Alkali metal hydroxides are extremely caustic and corrosive in high concentrations, even eating slowly through glass. Lab coats or disposable clothes are advised to quickly allow the removal of contaminated clothing if there is a spill.
  • If you are making your own equipment, thoroughly make sure it is water tight before beginning any electrolysis. Also certain plastics, such as PET, can be damaged by sodium hydroxide so know your materials and their exposure limits to sodium hydroxide. I used PET plastic for my first run and the apparatus held up fine over two days as I thought it would, but I would give it maximum a week before the sodium hydroxide eats a hole through the PET, especially as the concentration of sodium hydroxide increases over the experiment as we are making it.

Experimental:

IMG_7110[1]
The apparatus I had designed for the experiment. Going extremely low-tech, my first apparatus consisted of two PET plastic bottles connected by the bottle screw threads that were cut off each of them. They are joined and sealed with hot glue. Two cotton balls were stuffed in the connecting tube to act as the membrane between the cells as mentioned.

The experiment was set-up outside where it is well ventilated. I prepared a saturated solution of potassium chloride using 93.6g of potassium chloride and 300ml of distilled water (potassium chloride solubility: 31.2g per 100ml of water at 10oC). 300ml of distilled water was added slowly to the cathode cell and then immediately the 300ml of saturated potassium chloride solution was then added slowly to the anode cell. The solutions were added slowly as the first attempt of adding the solutions dislodged the loose cotton membrane due to pouring too quickly and forcefully. This mixed the solutions and so I had to start over.

I then had a problem as I just noticed that the power source I was going to use (7V, 200mA) was AC/AC and so was not converting into DC. AC would obviously cause the anode and cathode to constantly switch places therefore the power source was unfit for our needs. This was remedied by acquiring another old electronics charger (12V, 1.5A) that was AC/DC although which wire was positive and which negative was unknown and an hour of searching online yielded no conclusive answer. I know, I have got off to a great start already!

In the end, a quick water electrolysis cell was constructed to tell me which wire was positive and which negative. A 2L plastic bottle was used as the cell with two stainless steel butter knives (we have so much cutlery, my father said I was helping by reducing their numbers) were hot glued into holes at the base of the bottle. This allowed crocodile clips to be attached underneath to provide electricity to these stainless steel electrodes. I added an arbitrary volume of tap water to the cell and an arbitrary amount of impure solid sodium hydroxide (>=30%) drain cleaner, allowing it to mostly dissolve with gentle stirring. Two test tubes were filled with the dilute sodium hydroxide solution and turned upside down over the top of the electrodes without any air being trapped inside so as to collect hydrogen gas in one and oxygen gas in the other.

I stripped the two wires of the AC/DC power source with wire cutters to reveal some bare copper wire and this was connected with crocodile clips to the electrodes of the setup. Gas bubbles were formed on the electrodes immediately which were then collected by the test tubes. Over about 5 minutes of running, enough gas had been collected to show twice the volume of gas had been collected in one of the test tubes – hydrogen. The other test tube therefore had oxygen and this was confirmed by a pop sound as the hydrogen was lit with a flame from a match. Hydrogen is formed at the cathode therefore the wire of my power source connected to the cathode is the negative wire and so the other wire must then be the positive wire. A piece of grey electrical tape was used to designate the negative wire.

IMG_7115[1]
The quick water electrolysis cell I constructed to find out which wire was positive and which was negative. A mystery solved chemist-style.

I had constructed a water electrolysis cell on a smaller scale before this but it was unsuccessful. The thin copper wire electrodes used did not have a large enough surface area to produce hydrogen and oxygen fast enough, so I scaled it up.

The now known negative wire was connected with a crocodile clip to the cathode and the positive wire to the anode of the membrane electrolysis cell. The electrodes consisted of carbon graphite rods scavenged from Zinc Chloride batteries (see my post ‘Extracting zinc, manganese(IV) oxide and carbon electrodes from Zinc Chloride batteries’ – 18thtimelucky.wordpress.com/extracting-chemicals-from-zinc-chloride-batteries) with copper wire wrapped around it multiple times and secured with electrical tape. The copper wire is there to allow the crocodile clips to easily be attached as the carbon electrodes are very wide and the crocodile clips find it difficult to attach to them. Also it allows the electrodes to be hung on the side of the container and make the cell design simpler and easier.

To begin the electrolysis process the power source was turned on, although, after waiting a few hours, nothing seemed to be happening apart from some minute bubbling only on the anode. Both solutions were tested with universal indicator paper to reveal they were neutral. This suggested no reaction was occurring. It was possible electricity was finding it difficult to pass through the cathode solution due to just being distilled water so roughly a gram of potassium chloride was added and allowed to dissolve to act as an electrolyte. After a few more hours the results were the same so I rejected my hypothesis.

I transferred the solution of potassium chloride into a container and the experiment was repeated with sodium chloride. Potassium chloride is slightly harder to get hold of so wasting sodium chloride while testing was more desirable. I had only planned making potassium hydroxide originally, but now I decided I would make sodium hydroxide too, hence the title of ‘alkali metal hydroxides’.

IMG_7114[1]
Some of the bits and bobs I ended up using, although I used quite a few others things too.

After a few hours the results were the same again, as expected as the chemistry should remain the same. Although I spotted the cotton membrane did not stay in place properly in the tube connecting the cells due to being too wide so the solutions were mixing. This was most likely not the source of the problem but it was not helping either.

I sat down and pondered the possible reasons for the failure, included too low a voltage used, inadequate time allowed for the reaction and/or the electrodes were non-functional. Increasing the voltage would have been difficult, so to remedy the low voltage it was thought to reduce the resistance; the tube connecting the cells could be moved higher, less cotton could be used or the volume of solutions used could be lowered. If the electrodes were non-functional I could test this by using different materials such as stainless steel.

IMG_7130[1]
A new membrane electrolysis cell was made with a smaller tube that was comparatively nearer the top of the connected cells. Less cotton needs to be used. A much better design than the previous.

The experiment was set-up indoors, providing more comfort until the experiment was detected to be working from the smell of chlorine, then it would be turned off and moved outside. The only differences were thought to be a thinner tube, less cotton in the tube and the tube was nearer the top of the connected cells. When the power source was turned on, nothing happened. The power source was turned off.

I was running out of ideas; I decided to now check if the electrodes were functional. I thought I would try connecting the crocodile clips directly to the carbon electrodes, despite their difficulty of holding on properly. When the power source was turned on, bubbling immediately started. A chlorine smell was witnessed along with a very slight yellow colour developing in the sodium chloride solution and so the power source was turned off. This is very strange as this suggests the copper wire was not conducting the electricity. Whatever the reason, it was now working.

IMG_7113[1]
Damn you, cheap copper wires. I will have my revenge!

The set-up was moved outside. 72g of sodium chloride were dissolved in 200ml of distilled water and then filtered to produce a saturated solution and 200ml of tap water was prepared. The 200ml of tap water was added to the cathode cell and quickly after the saturated solution of sodium chloride was added. I used tap water due to the dissolved salts to act as electrolytes, although this is probably not needed and distilled water should work fine as the sodium hydroxide produced will act as an electrolyte. The power was turned on and the progress was recorded every hour apart from the first 30 minutes where it was monitored.

Within a few hours, the anode cell solution had become very noticeably yellow and the cathode cell solution turned universal indicator paper a blue-purple colour, highlighting the presence of a strong alkali (sodium hydroxide). It was surprising to witness such a dramatic change so quickly.

IMG_7132[1]
It is hard to tell the colour, but the right solution (anode cell solution) is slightly yellow-green. If the cotton was not functioning properly as a membrane, the yellow of the chlorine would have also have spread to the left solution (cathode cell solution). This is some more evidence the cotton can function as a membrane.

I decided to also test the pH of the anode cell solution, although I did not expect the universal indicator paper to turn slightly orange, showing the presence of a very weak acid.

The only possible hypothesis I could think of to explain why the anode cell solution was turning acidic was that the chlorine molecules being produced at the anode were undergoing homolytic fission due to UV light in the direct sunlight to form chlorine radicals. These radicals could react with hydrogen molecules produced at the cathode to form hydrogen chloride. When hydrogen chloride dissolves in water it forms hydrochloric acid, so small amounts formed could cause the very slightly acidic pH.

I had put a clear plastic box over the membrane electrolysis cell to protect from possible rain, aiding the above process as the gases are trapped together. To possibly prevent the formation of hydrochloric acid the gases produced could be kept separate with two boxes – one over each electrode cell. An opaque box could be used instead of a transparent box to possibly prevent some of the UV light from the sun from getting through. Either way, the hydrochloric acid would not be causing much trouble and could be left.

IMG_7133[1]
An overview of the whole set-up at this point.

I did not want to leave the cell running overnight but when I was packing it up I noticed little white specks all over the walls of the anode cell above the surface of the solution. The crocodile clip holding the anode was also covered in white specks with the exposed steel corroding with large rusting patches all over it.

It is possible the bubbling at the electrode as chlorine gas is produced could throw droplets of sodium chloride solution everywhere. The water then would evaporate off leaving the white specks of sodium chloride crystals. The sodium chloride solution could then cause the steel of the crocodile clip to rust, and if hydrogen chloride was being produced this would not help the rust problem either. I gave it a quick clean by scrubbing with some water and drying to remove some of the rust on the teeth of the clip. To try prevent further corrosion, I covered the crocodile clip with electrical tape to protect it from any spitting sodium chloride solution.

The electrolysis went without problems and at the end of the day I decided to stop the process and test the presence and concentration of the sodium hydroxide in my product. I also completed another run except now finally with the potassium chloride to produce potassium hydroxide. I used a slightly different cell this time, again with the same design, but at a larger scale, using thicker walled plastic containers with a slightly wider PVC tube connecting the two. This ran similarly and I did not encounter any problems, showing how the process should function the same for all alkali metals.

IMG_7184[1]
My newest cell design, again notice the vivid colour difference of the solutions. There is many things to notice about this picture. For instance, the surface area of the electrodes has been increased, using two electrodes on each side and larger stainless steel butter knives; there is no stainless steel as the anode as for some reason this causes the reaction to stop. I tried to speed up the electrolysis by connecting 9V batteries a friend donated to try amplify the voltage but this was very ineffective. I suggest only to use electronic chargers or battery packs for electrolysis from my experience.

Results:
The sodium hydroxide solution was weighed and the mass of 100ml of the solution came out at 104g, giving a density of 1.04kgL-1. Using a density-concentration table, this corresponded to a concentration of roughly 4% sodium hydroxide. In total, the yield was 205cm3 of ~1moldm-3 sodium hydroxide solution.

 

The potassium hydroxide solution was put through the same treatment, corresponding to a concentration of roughly 1.4%. In total, the yield was 600cm3 of ~0.25moldm-3 potassium hydroxide solution.

With a little further analysis, after 15 hours of electrolysis 0.205mol of sodium hydroxide was produced in total at 0.0137molh-1. The yield is a little miserable but 1moldm-3 is a useful concentration of sodium hydroxide to have and I am quite satisfied with the results.

The purity should be relatively high, although I do expect at least some contamination of the respective alkali metal chloride, hypochlorite and chlorate. I do not currently own a burette or any pH indicator solutions so I can not perform a titration to calculate a more accurate concentration of the hydroxide.

In the future I may try to use a higher voltage and current and keep the electrodes closer to each other to speed up the rate of reaction. I am still quite amazed how cotton wool functioned as a membrane though! The synthesis was a lot of work and did have many problems to solve but in the end it is a very nice route to alkali metal hydroxides through electrolysis and commonly available materials. The disbelieve and intrigue of the possibility of cotton as a membrane saw me through. I will definitely revisit this synthesis in the future.

Thank you for sticking with me through this nearly 4000 word post! I’ll leave you with some quantitative tests I did to see if I had actually made alkali metal hydroxides.

IMG_7398[1]
First off, a pH test. The solution turns universal indicator paper a deep purple showing the presence of an alkali. It is a very basic test so I did it again even though the next tests also show the presence of a hydroxide. Very basic test, eh? No? Bad joke? Sorry. No, I did not just do this test for the pun. Ignore the graduations on the glass container I am using to store the sodium hydroxide solution, for some reason they were printed higher up than they should be so give inaccurate measurements.
IMG_7411[1]
Ion tests are never a bother, they often gift you with such beautiful colours. Here, four test tubes are in a rack with dilute copper(II) sulfate solution in the first two, and dilute iron(II) sulfate in the last two. The iron(II) sulfate should be a faint green but is white due to impurities as I am using iron sulfate in the form of 98% fertiliser. Nevertheless it should still give a noticeable effect.
IMG_7401[1]
Hey Presto! As a few drops of the ‘sodium hydroxide solution’ is added to the first tube, a wispy blue precipitate crashes out. This precipitate is insoluble copper(II) hydroxide, showing the presence of a hydroxide which undergoes a double displacement reaction with the copper(II) sulfate. As some of the ‘potassium hydroxide solution’ is added to the second test tube, again copper(II) hydroxide crashes out showing the presence of a hydroxide. This is repeated in the third and fourth test tube, the formation of insoluble iron(II) hydroxide again suggests the presence of hydroxides.
IMG_7405[1]
I am not very keen on working with ammonia – toxic and stinky do not make a nice pair – but I decided to use the test for ammonium ions. In the presence of sodium hydroxide, ammonium compounds will react to form ammonia. A spirit burner here provides the gentle heating of a test tube containing ammonium hydroxide solution and some of the ‘sodium hydroxide solution’. The test tube is held in place with a clamp (which I forgot to put at an angle oops) so I can free up my hands to hold a piece of damp universal indicator paper at the mouth of the test tube to detect any ammonia produced.
 

IMG_7407[1]
The ammonia produced dissolves into the water on the damp universal indicator paper to form ammonium hydroxide again. Ammonium hydroxide is an alkali, conveniently turning the universal indicator paper blue-purple. As you can see from the purple colour of the paper, ammonia was produced showing the presence of sodium hydroxide.

Further reading or watching:

A great density to concentration conversion table website for many chemicals – extremely useful and interesting to have a look at:
http://www.handymath.com/cgi-bin/naohtble3.cgi?submit=Entry

Unfortunately I could not find the video I watched where he showed how you could use cotton as a membrane. If I find the video I will add the link here. There are a few others who have done similar videos on YouTube on this idea though.

One thought on “Synthesis of alkali metal hydroxides via electrolysis

Add yours

Leave a Reply

Fill in your details below or click an icon to log in:

WordPress.com Logo

You are commenting using your WordPress.com account. Log Out /  Change )

Google photo

You are commenting using your Google account. Log Out /  Change )

Twitter picture

You are commenting using your Twitter account. Log Out /  Change )

Facebook photo

You are commenting using your Facebook account. Log Out /  Change )

Connecting to %s

Website Powered by WordPress.com.

Up ↑

%d bloggers like this: