Reactions can be painfully slow without surface area on your side. I will be converting copper sulfate into copper powder through the aid of aluminium, in the form of kitchen foil, along with a special ingredient that is also very widely available. The convenience of the surface area gained from using copper powder instead of solid copper metal is worth the trouble. The difficulty of mechanically powdering the copper gives this wet chemical process the edge. Any copper salt should be suitable for this redox reaction, but copper sulfate is the easiest to purchase for the home chemist and who does not want to work with the beautiful blue of this compound?
Sure, metal sheets and wires can be a great source for copper. The extensive applications of copper means it can be found almost everywhere in the household and easily purchased from various stores. On average, 175kg of copper is associated with each person in the UK as pipes and in electronics.
However, the solid metal can be very slow to react in experiments, and the chemist is always looking to speed up a reaction if they can. The high surface area of powdered copper successfully gives it that extra boost in reaction rate that’s worth the trouble of powdering the copper.
Mechanically with solid copper this can be difficult, especially for the amateur, but the ability to do this chemically with copper salts can be granted by just a few household materials: aluminium, in the form of kitchen foil, and sodium chloride, obviously as table salt. The idea is that we are performing a redox reaction between aluminium and the copper salt.
Copper(II) sulfate is the salt of choice due to being cheap and easy to get hold of, especially in bulk. If you want to turn solid copper metal into copper powder chemically then you will need to convert it into a copper salt – hydrochloric acid, sulfuric acid or other acids should get the job done but may require some hydrogen peroxide to help oxidise the copper. Throughout the rest of this post I will be referring to copper(II) sulfate only, although as mentioned other copper salts should function similarly but remember to check the safety of the copper salt as well as the aluminium salt produced in this synthesis.
The redox (reduction-oxidation) reaction as shown by the ionic equation:
2Al (s) + 3Cu2+ (aq) → 2Al3+ (aq) + 3Cu (s)
The redox reaction as shown by the full equation if you would prefer:
2Al (s) + 3CuSO4 (aq) → 3Cu (s) + Al2(SO4)3 (aq)
Usually aluminium will not react with copper(II) sulfate; the passivation layer of unreactive aluminium oxide, developed when aluminium is exposed to air, prevents contact between the copper ions and the otherwise reactive aluminium. This is where the sodium chloride comes in. The aggressive chloride anions attack the oxide layer, causing it to become unstable and degrade. This allows the contact between aluminium metal and the copper ions. It also allows the contact between aluminium and water, producing a small amount of aluminium hydroxide and hydrogen gas.
Any other metals in the kitchen foil, if present, should also undergo a redox reaction with the copper, due to copper’s low position on the reactivity series. Theoretically, iron from nails or zinc from galvanised nails could be used in place of the aluminium foil, removing the requirement for sodium chloride, but the higher reactivity and surface area of the aluminium foil makes the reaction much faster and convenient.
Once the reaction is complete, we should be left with an aqueous solution with mainly aluminium sulfate and excess copper sulfate as well as a precipitation of copper powder. From here it is just purification steps to wash the copper powder to remove any soluble salts and then leaving it to dry.
It is harmful if swallowed, especially saturated solutions. The solid may irritate the eyes and skin. Water added to the anhydrous solid produces heat.
|Sodium chloride||Low hazard|
|Aluminium sulfate||Low hazard
Solutions are acidic.
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.
- Copper(II) sulfate (including Benedict’s reagent) is known as quite harmful to swallow and can also be irritating to the eyes. The drill of washing your eyes for at least 10 minutes should be used for eye contact but if you do swallow some, unusual for harmful materials, vomiting is recommended, as copper(II) ions are quite toxic to ingest. From experience though if you get this on your skin it is hardly dangerous, washing it off with water immediately solves the problem unless it has contacted any hair or your nails, of which, particularly your nails, will be stained blue, the blue only being removed by cutting the offending nail or hair (it clings to protein). Keep it in mind that it can be irritating to the skin for some and any prolonged contact is certainly not good so, again, gloves are highly recommended.
- The hydrogen is your main worry due to being potentially explosive, but such a small volume is produced at the small-scale like this experiment that it is a minor worry. I still do recommend performing the reaction outside or in a well ventilated area though.
- A small hazard may be the inhalation hazard of the dry copper powder so a dust mask is recommended and an air tight storage container is too.
Originally I decided to complete this reaction because I saw an interesting demonstration concerning this reaction and I thought I would adapt it to prepare copper powder. I do not know the name of the demonstration, but I have dubbed it ‘Copper out of the Blue’ for hopefully obvious reasons – I may do a post on it at some point. Without recording any measurements, I added a few scrunched up balls of aluminium foil to a 250ml beaker containing a solution of copper(II) sulfate and sodium chloride. I left the beaker outside to react as I was not sure how much hydrogen would be produced.
The reaction was very exothermic and required a water bath to keep it under control. Once the reaction had reached completion after a few hours, the beaker was taken inside. The solution was still slightly blue from the excess copper sulfate. A brown-red precipitate that had sunken to the bottom showed the formation of copper powder.
The upper aqueous layer was almost clear with seemingly no copper in suspension so it was decanted and discarded. An arbitrary volume of distilled water was added, the copper precipitate was stirred with a glass rod thoroughly, the solution was left to allow the copper precipitate to settle out again and then the upper aqueous layer was decanted and discarded. These steps were repeated 2-3 times with fresh distilled water to act as a purifying step, washing away any soluble compounds that would otherwise contaminate the copper.
Once I was satisfied with the washing steps, I transferred the solution to a piece of coffee filter paper with printer paper underneath to help absorb water and allow the copper sludge to dry.
Once the copper was dry it had a nice powdered consistency. It was transferred to a new sheet of printer paper cut to size so that the copper powder could be weighed. Afterwards it was poured into a 41ml glass jar for storage.
The beautiful brown-red of copper is very intense as a powder. The experiment seemed like quite a success! I do not usually have that without some problems in between. Easily available chemicals to the home chemist can be used to produce powdered copper, a more practical form of copper for experiments instead of solid metal. Now it was time to repeat the experiment, except with actual measurements so we can calculate a percentage yield.
I decided I would aim to theoretically form 5g of copper. Doing the calculations, this required 1.418g of aluminium foil and 21.622g of copper(II) sulfate in 5% excess. The copper(II) sulfate is in excess as it is easier to remove left over copper(II) sulfate than aluminium foil when purifying the copper formed. I used 2.5g of sodium chloride because I felt that was a good amount and there are no equations which I can use to work out how much sodium chloride I need.
I dissolved the solid copper(II) sulfate(VI) and sodium chloride in 150ml of distilled water. A slightly dirty solution reminded me I had not got around to purifying my copper(II) sulfate(VI) yet. I carried out a gravity filtration to remove any solid particles, successfully clearing up the solution.
I moved the 250ml beaker into a water bath to cool the reaction as the previous run revealed to me how exothermic the reaction is. The aluminium foil was then all thrown into the beaker and left to react by an open window.
The solution still got quite warm but not enough for much water to evaporate. Some cling-film was still put over the top of the beaker to catch any water that evaporated with a gap to allow any pressure build up to be released. The solution slowly lost its brilliant blue and became a murky blue-green as the aluminium reacted and dissolved, leaving copper to build up in the solution and all over the surface of the aluminium.
Once the reaction had died down after being left for a few hours, the solution left had the copper precipitate build up at the bottom like before but the solution was not blue. This suggests that all of the copper(II) sulfate(VI) has reacted with the aluminium which is not a good sign – this could mean there is some left over aluminium.
I decided I would carry on anyways with the washing and purifying steps as shown with the first run.
The copper sludge that resulted was left as before on some filter paper with printer paper below to help absorb any water and dry it. After a few days of drying, the copper powder was transferred to a sheet of printer paper to be weighed.
The first run of the experiment resulted in about 9g of red-brown copper powder although I do not know what the theoretical value is so I can not calculate a percentage yield. The second run though resulted in 5.645g of light brown copper powder. Wait, we know the theoretical yield for the second run – 5.000g. I could have got a very high recovery of copper and the rest of the mass is just water but the powder was very dry.
The best conclusion is that the copper from the second run is very impure and contaminated. The extra mass could be due to contaminants and is supported by how all the copper(II) sulfate had reacted during the experiment shown by the lack of blue in the solution, possibly due to losses at the added filtration step. This would leave some aluminium left, adding to the mass of the copper powder. The contamination of the copper is further supported by the colour of the powder. The light brown should be a dark brown, almost red, colour.
There is no point calculating a percentage yield as I do not know how much copper I even have. The pathway does work though and I do recommend it. I would recommend not doing what I did and recrystallise and purify your copper(II) sulfate beforehand so you actually do get an excess. Or you could just bump up your excess from 5% to 10% or 15%. I will try to revisit this experiment in the future.