Uploaded January 2018 | Updated September 2026, 1 week ago
In this video, we take the impure HgO made from the Nitrate in the previous video and gently heat it below the point of decomposition. It seems to get darker red and reversibly changes to orange again.
However, we were unable to find any explanation. It seems strange, but we saw this for so many simple compounds like the various lead compounds, too. How did nobody ever notice that color change in the history of analyzing various phases?
Perhaps you are luckier and find some good information we weren't able to find. What we know is that all yellow forms turn red because the small crystals stick together and thus become red but the reversible part is something we were unable to track down.
Youtube user 2 AM Productions came up with an interesting explanation:
youtu.be/lkiebvwfris
He had the idea that perhaps small droplets of Hg already form and coat the surface of the remaining oxide.
The question is how well the structure actually allows only partial cleavage. TiO2, for example, is known for its oxygen deficit structures and we showed you ZnO as well. It is not uncommon for oxides to lose some oxygen which later reacts with the cooled product again. That usually causes a color shift.
Also, the way how some crystals form could be an answer. The still hot Hg might be more reactive if adsorbed to the remaining HgO, making it easier for oxygen to attack.
If you have any suggestion we would love to hear about it.
In this video, we take the impure HgO made from the Nitrate in the previous video and gently heat it below the point of decomposition. It seems to get darker red and reversibly changes to orange again.
However, we were unable to find any explanation. It seems strange, but we saw this for so many simple compounds like the various lead compounds, too. How did nobody ever notice that color change in the history of analyzing various phases?
Perhaps you are luckier and find some good information we weren't able to find. What we know is that all yellow forms turn red because the small crystals stick together and thus become red but the reversible part is something we were unable to track down.
Youtube user 2 AM Productions came up with an interesting explanation:
youtu.be/lkiebvwfris
He had the idea that perhaps small droplets of Hg already form and coat the surface of the remaining oxide.
The question is how well the structure actually allows only partial cleavage. TiO2, for example, is known for its oxygen deficit structures and we showed you ZnO as well. It is not uncommon for oxides to lose some oxygen which later reacts with the cooled product again. That usually causes a color shift.
Also, the way how some crystals form could be an answer. The still hot Hg might be more reactive if adsorbed to the remaining HgO, making it easier for oxygen to attack.
If you have any suggestion we would love to hear about it.
![Copper Chemistry: Cu(II)/As(III) chameleon [no sound, 2x speed]](https://i.ytimg.com/vi/xHic2pew9bI/mqdefault.jpg "Copper Chemistry: Cu(II)/As(III) chameleon [no sound, 2x speed]
In this video, we show another way to make a beautiful spectrum of colors using Copper chemistry. The reaction is the following: Arsenic(III) is oxidized to Arsenic(V), while Copper(II) is reduced to Copper(I), all happening at a high pH value. But of course, we cannot just add cations and so different precipitates form and dissolve in this process creating many different colors.
We start off with a CuSO4 solution, which is light blue. To this, we add some NaAsO2. We already showed you Copperarsenite in another video. The blue-green precipitate also leaves a greenish aqueous layer on top. When we add KOH the remaining CuSO4 forms the dark blue Cu(OH)2. You can see the surface of the KOH turn dark blue. There is also an equilibrium between Arsenite and
Hydroxide and so the overall color turns more blueish in this process. Now the actual reaction also requires some heat. The KOH dissolving in so little water already causes it to become quite hot. This is why you suddenly see red-brown spots appearing everywhere. When we finally heat it all the Cu(II) is reduced to Cu(I) forming Cu2O.
Much like with our silver video, this is another case where some redox-reactions are just much easier accessible when high pH values are used.
For Arsenic(III to V) the potentials are 0,56 and -0,71 for pH=0 and 14 and for Copper(II to I) the potentials are 0,159 and -0,08.
To reduce Copper(II) to Copper(I) the redox potential of the reducing agent, here Arsenic(III) needs to be lower, which is only fulfilled at higher pH values.")
![Sulfur Chemistry: [Fe(CN)5NOS] from Sodium Polysulfide](https://i.ytimg.com/vi/xm7IYOrwQfA/mqdefault.jpg "Sulfur Chemistry: [Fe(CN)5NOS] from Sodium Polysulfide
There is an interesting reagent called Sodium Nitroprusside. It is sometimes used to identify several ions and is capable of distinguishing between Sulfite and Sulfide for example as well as forming many colored complexes. As we showed in our Cadmium Sulfide video, Sodium Polysulfide, made from elemental Sulfur and Sodium Hydroxide seems to be an alternative for the rather unstable Sodium Sulfide. So we tried to make the famous Sulfide test using Nitroprusside and a selfmade Polysulfide solution.
Here is the result. Note, as we didnt have any other source of H2S or Sulfide we could not compare this to the real test. So we assume the compound formed should have the composition as given in the title.")
