Uploaded September 2016 | Updated September 2026, 1 week ago
So Vanadium in one of these rainbow elements again, where many Oxidation states are accessible with lots of different colors. Usually this experiment is done with a Vanadate solution which is reduced by Zinc:Mercury (Amalgamate). So we wanted to do this reaction with something less toxic and looked for different reducing agents like Sulufours Acid and Nitrous Acid, both of them had no effect. In the end we came up with the plan to try nascent Hydrogen as we used this to reduce NIckel(II) to Nickel(I) (upcoming video) once and it kinda worked out quite well.
Now Vanadium is at this position where on one hand lots of Oxidation states can easily be achieved in aqueous (pH=0) solution which makes it suitable to be a good catalyst (since oxidation states vary this quick Vanadium is a great catalyst for Oxygen-reactions). On the other hand however we are at this point (mentioned in the Polychromate video) where aqueous chemistry stops as soon as come from Vanadium to Niobium or Tantalum, even aqueous Titanium chemistry is somewhat complicated. And these so called Isopolyacids, so long polymeric chains of condensed molecules which at one point precipitate to something we call an Oxide at some point makes Vanadium sometimes hard to handle as it is quite sensitive to pH environment.
Next problem was that we didn't have Sodium Vanadate but only the insoluble Ammonium Metavanadate. So we tried to convert it into the Oxide then dissolve it in NaOH to the Vanadate and add Aluminum so the NaOH and Al would form nascent Hydrogen. This however failed. As we added some dilute HCl to the metavanandate however it became dark red and the solution turned yellow. That was when we added Aluminum again and got all the oxidation states ! Now it's time to try and explain what happened, and it took us quite some time with a lot of literature to fully explain it.
The following explanation is taken from the Hollemann Wiberg chapter 26:
Okay assume you started with Vanadium(V)Oxide, which is the most stable compound. In bases this is converted to a Vanadate [VO4](2-). If this is acidified, the vanadates can much like the Chromate and Dichromate to Polychromate condense by giving off water and form corner-sharing tetrahedra in long chains. At pH 8 to 13 those shorter Vanadates and the in this experiment used Metavanadates (VO3(-))n form. At pH 2 to 6 the red/orange Decavanadates form and below pH 2 which I am not sure whether we really reached it as we only added some drops to it Vanadyl-cations form [VO2](+) or probably better if you saw our Ammine-copper/nickel video a [VO2(H2O)4](+). This can now be reduced with nascent Hydrogen or Hg:Zn and probably other reducing agents as well to Vanadium(IV) which is blue and looks like this [VO(H2O)5](2+). What is happening is the VO2(+) changes to a VO(2+) and as you see we loose an oxygen in this reaction.
The next step forms the green Vanadium(III) [V(H2O)6](3+) and the purple Vanadium(II) [V(H2O)6](2+). The color green appears in between but seems to be a mixture of the yellow Vanadium(V) and the blue Vanadium(IV).
There is an additional note on the last page in the book which says:
For V(III) its not [V(H2O)6]Cl3 but [VCl2(H2O)4]Cl as trans-Isomer.
For V(II) its not [V(H2O)6]Cl2 but again trans-[VCl2(H2O)4].
V(II) will react with water to form Hydrogen and V(III) again all of them will react with Oxygen to form V(V).
So Vanadium in one of these rainbow elements again, where many Oxidation states are accessible with lots of different colors. Usually this experiment is done with a Vanadate solution which is reduced by Zinc:Mercury (Amalgamate). So we wanted to do this reaction with something less toxic and looked for different reducing agents like Sulufours Acid and Nitrous Acid, both of them had no effect. In the end we came up with the plan to try nascent Hydrogen as we used this to reduce NIckel(II) to Nickel(I) (upcoming video) once and it kinda worked out quite well.
Now Vanadium is at this position where on one hand lots of Oxidation states can easily be achieved in aqueous (pH=0) solution which makes it suitable to be a good catalyst (since oxidation states vary this quick Vanadium is a great catalyst for Oxygen-reactions). On the other hand however we are at this point (mentioned in the Polychromate video) where aqueous chemistry stops as soon as come from Vanadium to Niobium or Tantalum, even aqueous Titanium chemistry is somewhat complicated. And these so called Isopolyacids, so long polymeric chains of condensed molecules which at one point precipitate to something we call an Oxide at some point makes Vanadium sometimes hard to handle as it is quite sensitive to pH environment.
Next problem was that we didn't have Sodium Vanadate but only the insoluble Ammonium Metavanadate. So we tried to convert it into the Oxide then dissolve it in NaOH to the Vanadate and add Aluminum so the NaOH and Al would form nascent Hydrogen. This however failed. As we added some dilute HCl to the metavanandate however it became dark red and the solution turned yellow. That was when we added Aluminum again and got all the oxidation states ! Now it's time to try and explain what happened, and it took us quite some time with a lot of literature to fully explain it.
The following explanation is taken from the Hollemann Wiberg chapter 26:
Okay assume you started with Vanadium(V)Oxide, which is the most stable compound. In bases this is converted to a Vanadate [VO4](2-). If this is acidified, the vanadates can much like the Chromate and Dichromate to Polychromate condense by giving off water and form corner-sharing tetrahedra in long chains. At pH 8 to 13 those shorter Vanadates and the in this experiment used Metavanadates (VO3(-))n form. At pH 2 to 6 the red/orange Decavanadates form and below pH 2 which I am not sure whether we really reached it as we only added some drops to it Vanadyl-cations form [VO2](+) or probably better if you saw our Ammine-copper/nickel video a [VO2(H2O)4](+). This can now be reduced with nascent Hydrogen or Hg:Zn and probably other reducing agents as well to Vanadium(IV) which is blue and looks like this [VO(H2O)5](2+). What is happening is the VO2(+) changes to a VO(2+) and as you see we loose an oxygen in this reaction.
The next step forms the green Vanadium(III) [V(H2O)6](3+) and the purple Vanadium(II) [V(H2O)6](2+). The color green appears in between but seems to be a mixture of the yellow Vanadium(V) and the blue Vanadium(IV).
There is an additional note on the last page in the book which says:
For V(III) its not [V(H2O)6]Cl3 but [VCl2(H2O)4]Cl as trans-Isomer.
For V(II) its not [V(H2O)6]Cl2 but again trans-[VCl2(H2O)4].
V(II) will react with water to form Hydrogen and V(III) again all of them will react with Oxygen to form V(V).
![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.")
