Uploaded April 2017 | Updated September 2026, 1 week ago
This is the first video of hopefully a huge collection of Ruthenium related experiments. We've been trying to somehow get RuCl3 to react in the past weeks but most of our experiments didn't work out. So we tried to improve them a bit by either changing pH environments or temperature setups which resulted in better solutions. We believe that, as we are using a nearly anhydrous RuCl3, every time we prepare a fresh solution the Ru-Cl bonds are still too strong to react with weak ligands. Ru in a chloro and aqua environment goes through several ligand exchanges till an aqua complex is formed. So we believe that for example heating it helps to speed up these processes a little. At least our experiments finally worked when we used a little bit of heat
In this experiment we added Ammonium Thiocyanate to a Ruthenium(III) Chloride solution. At first nothing happens although it is supposed to turn red...at little bit of heat and we can see a lovely red colour which suddenly turns blue.
As far as we could find any literature on it this colour change is believed to be caused by a linkage isomerism of the Thiocyanate.
There are famous Nitro and Nitrito complexes whose colours change if the ligands coordinate the same way but with other elements, like with N or O in the nitro and nitrito isomers. This seems to be a similar case.
Tests, where this Thiocyanatoruthenate(III) had been isolated with a tetrbutylammonium-cation showed that the complex anion looked something like this
[Ru (NCS)x (SCN)6-x](3-).
Depending on how x varies the colour changes, while a big concentration of Isothiocyanate (4 x N-coordination) is red, 5x an S-coordination seems to be blue.
This is the first video of hopefully a huge collection of Ruthenium related experiments. We've been trying to somehow get RuCl3 to react in the past weeks but most of our experiments didn't work out. So we tried to improve them a bit by either changing pH environments or temperature setups which resulted in better solutions. We believe that, as we are using a nearly anhydrous RuCl3, every time we prepare a fresh solution the Ru-Cl bonds are still too strong to react with weak ligands. Ru in a chloro and aqua environment goes through several ligand exchanges till an aqua complex is formed. So we believe that for example heating it helps to speed up these processes a little. At least our experiments finally worked when we used a little bit of heat
In this experiment we added Ammonium Thiocyanate to a Ruthenium(III) Chloride solution. At first nothing happens although it is supposed to turn red...at little bit of heat and we can see a lovely red colour which suddenly turns blue.
As far as we could find any literature on it this colour change is believed to be caused by a linkage isomerism of the Thiocyanate.
There are famous Nitro and Nitrito complexes whose colours change if the ligands coordinate the same way but with other elements, like with N or O in the nitro and nitrito isomers. This seems to be a similar case.
Tests, where this Thiocyanatoruthenate(III) had been isolated with a tetrbutylammonium-cation showed that the complex anion looked something like this
[Ru (NCS)x (SCN)6-x](3-).
Depending on how x varies the colour changes, while a big concentration of Isothiocyanate (4 x N-coordination) is red, 5x an S-coordination seems to be blue.
![Chromium Chemistry: [2xfaster] Chromium in Gas-Phase (Chromylchloride)](https://i.ytimg.com/vi/qeQLe06sV0E/mqdefault.jpg "Chromium Chemistry: [2xfaster] Chromium in Gas-Phase (Chromylchloride)
Chromyl Chloride is a well known compound in organic chemistry. We tried using its property to become airborne quite easily and tested if the gaseous Chromium would react with some Hydrogen Peroxide in another vial to form the blue Peroxo-complex.
So we prepared two vials, one containing some 3% Hydrogen Peroxide and diluted H2SO4 and the other containing a mixture of KCl and K2Cr2O7. To the second one a few drops of conc. H2SO4 were added. It starts to foam and an orange (hard to see on camera) gas forms. On the video it looks quite pale but it was quite dense in reality. On the glass walls you can see some red drops of liquid Chromyl Chloride. We covered both vials with a beaker and waited. After some time a blue complex forms proofing that the yellow gas is not only Chlorine but also a gaseous Chromium compound.
To understand the structure of Chromylchloride a bit more it is easier to remind yourself that this compound can be prepared if Chromic Acid is reacted with Hydrochloric Acid, too.
Now we had this in the Polychromate video already. Chromic acid is a tetrahedron having two Cr-O bonds and two Cr-OH bonds. Back then we added acid to form a Cr-O-Cr bond. But what if we dont have the H(+) acting here but the Cl(-) ? The Cl(-) could substitute the OH(-)
Cr-OH + HCl to Cr-Cl + H2O. this happens to both of the Cr-OH groups forming the CrO2Cl2. Now the water reacts with this compound so the would would have to be captured here. One way to do this is using conc. H2SO4 which is strongly hygroscopic. And the HCl can be substituted by NaCl for example to reduce the amount of water even more. And we showed that Chromic acid forms if Chromates meet Acids, so the setup can be reduced to having a Dichromate or Chromate, a Chloride salt and some conc. H2SO4.")
](https://i.ytimg.com/vi/qzjkOAttUHw/mqdefault.jpg "Bismuth Chemistry: [Bi(SCN)6](3-)
If Thiocyanate is added to a Bismuth(III) solution the color will change to yellow upon the formation of [Bi(SCN)6](3-).")
![Chromium Chemistry: Burning (NH4)2[Cr2O7] (no volcano)](https://i.ytimg.com/vi/uCwS1bTaX-8/mqdefault.jpg "Chromium Chemistry: Burning (NH4)2[Cr2O7] (no volcano)
You are probably familiar with the famous Ammonium Dichromate volcano. If you light some (NH4)2[Cr2O7] it will form Cr2O3 while creating burning sparks that form a sort of volcano. We tried to do that but our Dichromate is quite wet, unfortunately. So we tried a different idea and heated it directly with the burner. The temperature is high enough to vaporize the water and burn the Dichromate, creating much more sparks than usually.
What is happening is that the system contains two parts, in Ammonia the Nitrogen is reduced. If it is oxidized a stable N-N triple bond forms. On the other hand, the Dichromate contains a highly oxidized Chromium(VI). The stable form is Cr(III) so in exchange Cr(VI) is reduced and N(-III) oxidized.
This is also the reason why we didnt dry our Dichromate for the volcano. Many of these combinations of Ammonium and strong oxidizers (e.g. Ammonium Nitrate) tend to combust when heated. The dry compound starts to decompose around the boiling point of water. Although Ammonium Dichromate is usually not that reactive.")
 an NO (+*-) as Ligand](https://i.ytimg.com/vi/uLh69ukYmrU/mqdefault.jpg "Iron Chemistry: [Fe(H2O)5NO](2+) an NO (+*-) as Ligand
This is a very famous test for Nitrate-Ions:
In this case we added some Iron nails for one day to 38% H2SO4 and then added some of the clear solution to a KNO3 solution. If conc. H2SO4 is very carefully added a brown-purple ring forms.
The ring has the composition [Fe(H2O5)NO](2+).
For many years people believed that Iron has the Oxidation state +I in this compound since NO is actually a Nitrosyl NO(+). Other suggestions say its Fe(II) or Fe(III).
..................................................................................
Just thinking about this is sounds strange for this to be Fe(I). Fe(II) is really hard to store as solution since it easily oxidizes to Fe(III). Thus it should become an Fe(III) really easy and not an Fe(I) (where no real compounds exist, some exceptions of course). On the other hand Nitrate is reduced to NO so something clearly has to be oxidized for this to happen which is Fe(II) becoming Fe(III).
We can also look this up in the redox series:
NO3(-) + 4H3O(+) + 3e(-) to NO + 6H2O (U = 0,96 V)
Fe(II) to Fe(III) + e(-) (U = 0,77 V).
..............................................................................
But can NO act in three different ways as a ligand ?
As you can see in this picture, neutral NO is already a radical
http://daten.didaktikchemie.uni-bayreuth.de/umat/stickstoffoxide/no_moschema.gif
So we could write it as NO* for a second. If another electron is added to form NO(-) we reach something that reminds us of Oxygen and if we take an electron away to make NO(+) we can get something which is isoelectronic to Carbonyl (CO). And Oxygen and Carbonyl are famous Ligands. This is why all three seem possible depending on the electronic situation making the Iron a (I), (II) or (III). It gets even harder as there are multiple ways how NO can coordinate to a metal center, but we wont further discuss this here.
Quite new calculations and research show that it is really an Fe(III) and an NO(-) present.
But there is more! If you followed our video series you might have seen videos where we used Nitroprusside. In Nitroprusside an NO(+) is present. If this compound is reduced even further you might expect NO* to form. Analyzing the products showed that for some percent, some say even 25%, Fe(I) forms !
Still many of these compounds remain a mystery and are worth to be further researched.")
 (Pd in HCl/H2O2)](https://i.ytimg.com/vi/wQA5K90Ijqg/mqdefault.jpg "Palladium Chemistry: [Pd(Cl4](2-) (Pd in HCl/H2O2)
Palladium is an interesting element. Being a Platinum-group metal and standing between Nickel and Platinum one might expect it to be unreactive. Funnily this is not the case and some people even joke it could be more reactive than Nickel. Obviously, this might be a bit too much but the reactivity can be compared to the one of silver with one important difference: Palladium usually does not build up many passive layers. Meaning that it can dissolve in HCl for example.
In this video, however, we cheated a little and this is first to make it shorter and second to make it more visible. Usually, oxygen in HCl is enough to slowly dissolve Palladium but we used H2O2 instead of the oxygen here. Now you might say there is no big difference but there is. The redox reaction now changed to the one of an acidified Peroxide which has a huge redox potential being capable of even dissolving gold and platinum to some extent!
And PdCl2 forms Pd(IV) in H2O2 as well!
Note however that we also tried it with HCl and Oxygen and a lot of boiling and it turned yellow as well!
What forms here is the complex [PdCl4](2-). While Nickel(II) already shows some square planar geometry, especially when working with cyano-Ligands for Palladium almost all complexes are square planar.
Also the coordination number of 4 is really common in Palladium chemistry. So if you have to guess, a square planar geometry is always a good guess when dealing with Pd(II). Even the aqua-complex only has 4 Ligands.")
