100 Comments

  1. There is no gas as a product of the reaction. The bubbles you see there are just boiling water (solution) due to the highly exotermic reaction. This is a simple redox reacton:
    Cu2+ (imagine as superscript) + Al – –> Cu + Al3+

  2. im afraid to say that is wrong :/ Cu+ is colourless due to its complete d-subshell. it is greeny blue due to the copper(II) chloride impurites. and in fact they are both not even green or blue unless hydrated (by either moisture in the air or in solution)

  3. Sounds like they finally got around to it that way, doesn't it?

    Martyn seems like the one of the funnest dads in the wholest widest world. As a husband…well, I bet flooring budget had more to it than a carpet. I bet there was more than a flooring budget.

  4. It does lead one to wonder why Bradey/TestTube/Nottingham doesn't sell Martyn dolls. I might play bingo just to pull out a small collection of Martyn dolls to confuse the Troll doll users ;o)

  5. I demo'd this reaction for my department today and it is now gonna be part of our year 8 scheme of work. Great little bit of chemistry this.

  6. I would just like to say, thank you so much for making all of these videos! They are so awesome! I learned so much

  7. i should be watching this 2 years ago when i'm having my final science exams. this is so much better compare to my teachers. 😀

  8. Actually, the solution of CuSO4 (Copper-II) is blue due to Copper-Sulfate-Pentahydrate-Complexes.
    CuO (also Copper-II) is black.

  9. 17 likes for something totally wrong lol. Copper(I) itself is d-10 and colorless, it is usually insoluble and unstable in aqueous solutions. Copper(II) sulphate is blue because [Cu(H2O)6]2+ is blue; For copper chloride copper makes a yellow complex ion [CuCl4]2- therefore copper(II) chloride appears more green than blue.

  10. As far as the gas goes it's probably a mix of steam (from the heat) and HCl; aluminium chloride, once formed in aqueous solutions, will hydrolyze immediately to form Al(OH)3 and HCl. I suspect that the reaction won't be as spectacular if one uses copper sulphate instead of copper chloride because aluminium sulphate doesn't hydrolyze as much as AlCl3 (but I could be wrong)

  11. Concentrated solution of copper(II) chloride is green but when the concentration is lower, the solution is blue.

  12. I did this experiment as part of a larger one in my Chem class. We started off with some copper wire and went through a few different compounds of copper and ended up with this. Really fun except we just put some aluminium foil into our beaker and saw it go from there and decanted the solution from there.

  13. No, no no. The truth is much more complex than that. Notice how the crystals are all green, and a concentrated solution is green? Contrary to what simply high school chemistry teaches you, copper almost never exists in solution as the unbound ion, it bonds with suitable ligands to form a complex. The green crystals probably results from the excess HCl contamination in the crystals, forming a negatively-charged complex CuCl4(2-) via the attraction of the positively charged-cont

  14. Cu2+ ions to the Cl- ions and acts as the ligand instead of water, because not much water is present. When the copper (II) chloride dries, water still remains as the hydrate, CuCl2.2H2O, in which the crystal is still green because the Cu2+ is still bound to the Cl-, making this more like CuCl4(2-).2H3O+ (approximation). When you dilute this with more water, the copper instead becomes attracted to the water molecule, because there are more water molecules than Cl-, so-cont

  15. The copper forms the typical complex erroneously called the Cu2+ ion, the [Cu(H2O)6]2+ complex/ion, which has a blue colour, instead of the green of CuCl4(2-). Please forgive me if I have made a mistake in my explanation, this colour difference is very complex indeed.

  16. Yet, the label clearly states that it is a copper (II) chloride. Actually, the ligands in solution may alter its color quite substantially. Here, chloride ions make complexes with Cu2+ ion that appear green, but on addition of concentrated HCl it can even change color to yellow due to formation of [CuCl4]2- complex anion. On the contrary, copper (I) compounds are mostly quite insoluble in plain water and appear green because of copper (II) impurities.

  17. wow lucky me, the other day i was doing electroplating with this and the only metal i didnt try plating with copper as aluminium XD

  18. I did a similar reaction using mercury(i) cloride instead of copper cloride, but the result was different (I obtained elemental HG and a foam of aluminium oxide) it was a very beautyfull reaction 😀

  19. Top draw on Prof's past home chemistry antics … nothing like a safe home demo that had a funny unplanned consequence 🙂 Three cheers for Poliakoff … hip-hip!

  20. this is a great video showing the reaction of metals in the reactivity series, aluminum is more reactive, so it takes the place of the copper and leave behind metallic copper. same thing happens when you introduce copper into a silver salt solution or even chloroauric acid as a result of dissolving gold into aqua regia. the less reactive metals like silver, gold and platinum will precipitate out of solution if you introduce copper or any or the more reactive metals.

  21. @periodicvideos Is the first reaction 12CuCl2 + 2Al2O3 -> 12Cu + 4Al2Cl3 + 3O2 as the aluminium oxide layer is consumed, and then as the copper(II) chloride reacts with the aluminium metal, 3CuCl2 + Al -> 3Cu + AlCl3, forming the copper metal deposit ? Or is there some copper oxide formed?

  22. @periodicvideos If you want more views on your videos, just put a picture of the professor as a thumbnail on each one. lol

  23. Everytime I see a thumbnail of the professor, I have to fight not to click on it… *must not watch it, won't be able to stop later, must finish homework first*

  24. Do chemists perform any reactions that cannot take place in a glass beaker because it would have a bad reaction with the glass?

  25. Actually,hot concentrated sodium hydroxide and hydrofluoric acid both react with glass pretty well, have to use polyethylene or teflon for those, dilute NaOH isn't too bad though.

  26. So.. If I got this right.. You took some Copper (II) Chloride from your lab to your home to make the expirement to your kids?! hehe.. <3

  27. Did the same experiment in my chemistry class except we put a piece of aluminum wire into the copper chloride solution which turned into a chunk of copper

  28. Aluminum is a correct spelling, & the American (& Canadian) way of pronouncing it is also correct. British chemist & inventor Humphry Davy settled on aluminum by the time he published his 1812 book Chemical Philosophy. I think calling it aluminum honours Humphry Davy, who identified the existence of the metal base of alum.

  29. I need to make a marble fall through an aluminum bridge with some sort of reaction. it needs to fall in 2-4 minutes. the Copper(ii)Chloride reacts too quickly. Should I decrease the molarity and use more?

  30. Hey. Add English captions as well because you have some subscribers that can't hear. With them the Portuguese captions is not good enough

  31. My guess is that Pr Poliakov was fed up with that old and ugly carpet and he pulled that trick of educating the children to alleviate mommy's anger.

  32. What's really interesting is what DOESN'T happen when this is performed with copper sulfate or nitrate.  The chloride is catalytic, and I'm not sure anyone yet has an explanation as to why.

  33. I would like to modify aluminium cylinder head intake tracks with this process-could you direct me to a copper chloride supplier? It sure is nice to hear aluminium called by it's real name, you hear that much in the racing industries…

  34. How will u find the exact percentage of copper in cuprous chloride that has been prepared… By instrumental analysis… To determine its purity

  35. The moment I saw and heard this Professor I subscribed because this may be the closest I will ever get to a Potion's Class at Hogwarts. I solemnly swear that…. I don't want to learn from anybody else.

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