Cyanate ion: uncommon but interesting ion

This web page demonstrates some properties of the cyanate ion, OCN–. This ion is not really common. Many people may have heard of it, but only few people really have done experiments with it. It appears that the ion has interesting coordinating properties, but its behavior on addition of acids also is quite interesting.

This web page gives a series of experiments, which all can be performed easily and safely. Not all experiments need to be done. If one of the required compounds is not available, then other experiments on this page still can be done if those reagents are available.

 

Required chemicals:

  • sodium cyanate or potassium cyanate
  • moderately concentrated nitric acid (35% or so)
  • moderately concentrated sulfuric acid (30% or so)
  • cobalt sulfate
  • silver nitrate
  • vanadyl sulfate
  • chrome alum or normal (purple) chromium(III) sulfate
  • ferric ammonium sulfate
  • copper sulfate
  • nickel sulfate

Required equipment:

  • test tubes
  • small beaker
  • many plastic spatulas

Safety:

  • Nitric acid of moderate concentration is quite corrosive. Avoid contact with skin and eyes.
  • Sulfuric acid of moderate concentration is quite corrosive. Avoid contact with skin and eyes.
  • Vanadyl sulfate is fairly toxic. Avoid exposure.
  • Cobalt sulfate is toxic. Avoid exposure.
  • Nickel sulfate is toxic and possibly carcinogenic. Avoid exposure.
  • Silver nitrate is caustic and causes black stains on the skin. Avoid contact, particularly contact with the eyes must be avoided!

Disposal:

  • Waste of vanadium, cobalt, nickel, silver and copper should not be flushed down the drain. Bring waste of these metals as heavy metal waste to a municipal waste processing facility.

 
Sodium cyanate is non-toxic, despite it being derived from cyanide ion. In acidic conditions (such as exist in the stomach) it is converted to ammonium ions and carbon dioxide in a matter of minutes.

Preparation of a solution of sodium cyanate

If you perform a row of different experiments with a solution of sodium cyanate, then it is easiest to prepare a little stock (appr. 30 ml of solution) in advance. This is enough for all experiments on this web page. Below follows a picture of some solid NaOCN in a test tube. The solid is a white crystalline powder, not looking particularly spectacular.
     

     


I had a one-time opportunity to buy a nice amount of this chemical for just a few euros. Here follows a picture of the old bottle, which still is in very good condition.

     


Put the approximately 2 grams of solid sodium cyanate in a small beaker.

Add approximately 30 ml of cold distilled water to the solid.

Using a glass rod, stir the liquid, until all of the solid has dissolved. This takes some time. Sodium cyanate does dissolve fairly well in water, but it does so slowly.

The resulting liquid is clear and colorless, having a concentration of nearly 7% by weight of sodium cyanate. For each of the experiments, 4 to 5 ml of solution is used in test tubes.


 

Formation of isocyanic acid and decomposition

Cyanate ion reacts with strong acids, giving mainly isocyanic acid, O=C=N-H. This tautomerizes to cyanic acid H-O-C≡N in an equilibrium reaction. Approximately 97% of the acid exists as isocyanic acid and approximately 3% exists as cyanic acid, when in aqueous solution. The cyanic acid is unstable and quickly hydrolyzes in the acidic solution to give ammonium ions and carbonic acid. The latter decomposes to water and carbon dioxide.

In this experiment, the reaction between sodium cyanate and an excess amount of strong acid is demonstrated. Most interesting is the observation that the formation of carbon dioxide is not instantaneous. This is in strong contrast to adding an excess amount of strong acid to a solution of sodium carbonate or sodium bicarbonate, where the liquid immediately foams up strongly and all carbonate is decomposed at once after adding the acid.
With the cyanate, the liquid at first does not seem to react at all. After a few seconds, the liquid starts bubbling and slowly the intensity of the bubbling increases. The liquid also heats up a little. The 7% solution of sodium cyanate becomes quite warm (but not hot), a few tens of seconds after adding the acid.

        

In this video one can nicely see the delayed formation of carbon dioxide on addition an excess amount of 30% sulfuric acid.

 

 

White precipitate with silver(I) ions

In this experiment, it is shown that silver(I) cyanate is insoluble in water. This is a simple experiment.

Take 2 ml or so of the stock solution of sodium cyanate, and add 2 ml of water. This makes a somewhat more dilute solution.

In a separate test tube put some silver nitrate and add a few ml of water.

Pour the solution of silver nitrate in the solution of sodium cyanate.

The result is formation of a very pale beige (nearly white) precipitate of silver(I) cyanate. Click here for a video of the reaction. The pale beige solid, formed in the reaction, is not a true salt, but a polymeric species, with a structure as shown in the picture below. See also the Wikipedia page about this compound.


               

On addition of excess nitric acid, the solid dissolves, but only quite slowly. A lot of stirring is needed to get the solid dissolved and if the reaction must proceed at reasonable speed, then quite some excess acid is necessary. But if sufficient acid is present, then finally, after several minutes, the solid dissolves again, giving a colorless solution, and all cyanate is destroyed with slow production of carbon dioxide.


 

A bright blue cobalt(II) complex

In this experiment, an interesting deep blue cobalt complex is formed. This complex is quite labile. It is demonstrated that it is in equlibrium with hydrated cobalt(II) ions and when the cyanate ion is taken away, then the complex quickly falls apart again.

Put a few ml of the 7% solution of sodium cyanate in a test tube.

In another test tube, prepare a solution of cobalt sulfate. A small spatula of the solid in a few ml is sufficient for this experiment. The solution, thus obtained, is pink/rose.

Below, a picture is shown of the two solutions:

    


When these two solutions are mixed, then a bright blue solution is obtained. A video is available of this reaction. The picture below shows the final result of the reaction.

              

This blue complex is quite specific for the combination of cobalt(II) and cyanate. It exists at low concentration and is different from many other blue cobalt complexes, which require a high concentration of the ligand (e.g. for formation of the blue tetrachloro complex of cobalt(II) one needs a very concentrated solution of NaCl or quite concentrated HCl).

The complex also can exist at low concentrations, which means that the ligands are fairly strongly attracted by the cobalt(II) core of the complex ion. On the other hand, the complex is sufficiently labile to allow for sufficient free cyanate ions, which can be taken away from the solution very quickly. This is demonstrated in two ways:

  • Addition of excess amount of silver(I) ions, which leads to precipitation of AgOCN.
  • Addition of an excess amount of a strong acid, which leads to immediate formation of O=C=N-H. The protonated ion cannot coordinate to the cobalt metal ion and hence the blue color disappears again and the pink color of aqueous cobalt(II) appears again.

Click the above links for videos of these reactions.

The picture below shows the result of adding silver nitrate to the blue complex, after allowing the precipitate to settle somewhat at the bottom.

         

This picture nicely shows the pale beige color of the precipitate of AgOCN, the pink color of the free aqueous cobalt(II) ions and some blue liquid, which did not mix with the main part of the solution.

 

Several metal complexes

In the above experiment, the cobalt(II) complex of cyanate ion is prepared and some properties are demonstrated. Other metal ions, however, also form complexes with cyanate ion and some of them are quite interesting. All of the prepared complexes are quite labile, just like the cobalt(II) complex.

First, metal salts are dissolved in distilled water, in different test tubes, which were put in a glass container in order to be able to make pictures of the solutions.

 

The solutions are fairly dilute. From left to right these are solutions of the following chemicals:
  • vanadyl sulfate
  • chromium(III) sulfate
  • ferric ammonium sulfate
  • copper sulfate
  • nickel sulfate

Next, to each of the test tubes, some of the 7% solution of sodium cyanate is added. A picture is taken, immediately after adding the solution.

 

The picture shows that in general, all solutions obtain a more intense color, while the metal salt is diluted by the addition of the solution of sodium cyanate. The color of the chromium complex also shifts from purple to grayish green.


While looking at the complexes, one can see a slow change of the colors after the initial immediate color change. The color of the chromium complex shifts towards brighter green and the color of the iron(III) complex shifts towards red and also intensifies. A new picture was made three minutes later.

 

The picture above shows how the colors are observed under high quality LED light or under daylight by the eye. On a digital camera, the color of the chromium(III) complex looks different from what you see with the eye:

 

The last picture (immediately above) is the picture, as it comes directly from the digital camera, but this is definitely not how it looks to the eye. The previous picture is how it is perceived by the eye, but in order to get the image like that, it was necessary to change the hue settings for the chromium(III) complex (the hues for the other complexes were not adjusted).

The strange behavior in color rendering of the digital camera for the chromium complex is most likely due to the presence of narrow absorption bands in the visible spectrum. Digital photo cameras have difficulty rendering colors correctly, when the spectrum of the light is not spread smoothly, but has narrow bands and peaks.


Finally, some 30% sulfuric acid is added to each of the test tubes in order to see whether the complex remains present or is destroyed immediately. For each of the metals, the complex is destroyed immediately and the solutions become pale again. Several seconds later, the liquids start bubbling. So, in a very fast reaction, the metal complexes are broken down and the cyanate-ligands are protonated and then in a subsequent slower reaction, the protonated cyanate decomposes, giving carbon dioxide.

 

The iron(III) complex almost completely disappears and becomes colorless. This is due to the excess amount of added acid, which causes the brown partially hydrolyzed iron(III) to be converted to nearly colorless non hydrolyzed aqueous iron(III).

 


 

 

 

Discussion of results

Solid sodium cyanate exists in anhydrous form, as NaOCN. When this is dissolved, the solution obtains sodium ions and cyanate ions:

    NaOCN(s)  →  Na+(aq) + OCN–(aq)


 On addition of a strong acid, nearly all of the the cyanate is protonated:

   H+(aq)  +  OCN–(aq)  →  O=C=N-H(aq)        (this is isocyanic acid)

Isocyanic acid tautomerizes to cyanic acid in an equilibrium reaction. The equilibrium is at 97% at the isocyanic acid side and 3% at the cyanic acid side.

    O=C=N-H  ⥃  H-O-C≡N 

The latter is not stable and reacts with water and excess acid in a somewhat slower reaction with formation of ammonium ions and carbonic acid, which in turn decomposes to water and carbon dioxide:

    H-O-C≡N  +  H+  +  2 H2O  →  NH4+  +  H2CO3  →  NH4+  +  H2O  + CO2

This explains the bubbling after addition of a strong acid to a solution of a cyanate. In practice, cyanate ion cannot exist at low pH for more than a few tens of seconds.


 With silver ions, a polymeric species is formed, which is insoluble in water. This species is very pale beige, nearly white.

   2n Ag+(aq)  +  2n OCN–(aq)  →  [⋯Ag⋯N(=C=O)⋯Ag⋯N(=C=O)⋯]n (s)

The structure of the solid is given below, for one unit of the polymeric species:

                            

The silver ions are bound to two N-atoms, but not with full bonds between the silver atoms and nitrogen atoms.

This compound has very low solubility in water and it also is not easily protonated. On addition of a strong acid, it does dissolve again, but only very slowly. The acid protonates the cyanate-ions, and this leads to breaking of the half-bonds with the silver ions. This causes the chain to break apart and in this way, the solid slowly dissolves. At first, the liquid becomes somewhat opalescent, but after a longer time, when the chains are broken up to very short pieces, the liquid becomes clear.


 With many other metal ions, cyanate forms complexes. The usual coordinating site is at the N-atom, but some complexes also are at the O-atom. There even are bridging complexes, with one metal ion coordinated to the N-atom and another metal ion (which need not be of the same metal) at the O-atom. The complexes, formed in the experiments, most likely are coordinated at the N-atoms, but this is not certain. Further research is needed to be conclusive about the precise structure of the shown metal complexes.

 

   

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