# Analysis of group 4 cations

Zinc identification

Zinc is inside the alkaline solution (NaOH) as tetrahydroxy zincate, Zn(OH)42- . There are various ways to confirmate zinc. We will divide then the solution in as many quotes as the number of identification tests that you want to perform.

-precipitation as zinc sulfide, ZnS

The solution is added with diluted HCl. Zn(OH)42- is dissolved to restore Zn 2+ in solution. However, we can't use this solution to precipitate zinc sulphide. Indeed, at acidic pH the concentration of sulfide ion is too low to trigger the precipitation. It is then added concentrated NH4OH to obtain an alkaline environment. Zinc is now present as amino complex. Subsequently we add the source of sulfide ions (Na2S), which now may precipitate, if present, zinc sulfide, black.

Zn(OH)42- + 4H+ $\dpi{120}&space;\rightleftharpoons$ Zn2+ + 4H2O

Zn2+ +  xNH3 $\dpi{120}&space;\rightleftharpoons$ Zn(NH3)x2+

Zn(NH3)x2+ + S2- $\dpi{120}&space;\rightleftharpoons$ ZnS + xNH3

-precipitation as zinc ferrocyanide Zn2Fe(CN)6

The same solution is added with a few drops of potassium ferrocyanide, K4Fe(CN)6. If Zn was in the initial unknown substance, we will get white-light blue precipitate of zinc ferrocyanide, Zn2Fe(CN)6. If we didn't get the precipitate then try to acidify with CH3COOH, in order to ensure a higher concentration of Zn2+ ions. Furthermore, it avoids the decomposition of ferrocyanide, unstable compound in alkaline solutions.

It might not be easy to appreciate this color and is therefore recommended to accurately wash an the precipitate. An excess of potassium ferrocyanide may precipitate mixed ferrocyanide of zinc and potassium, white.

- Identification with dithizone (diphenylthiocarbazone)

The presence of zinc can be detected through a colorimetric test using dithizone. Dithizone is a dark green solid. The zinc-dithizone complex is colored in red; the coloring is instead yellow-orange in absence of copper.

Manganese identification

We're goint to confirmate manganese in the precipitate that is supposed to be manganous hydroxide, Mn(OH)2,. This will be partially burnished due to the action of atmospheric oxygen MnO(OH) or Mn(OH)3 . 1 mL of concentrated nitric acid (HNO3) is requested to solubilize this precipitate; then we add a spatula of PbO2, lead dioxide (or plumbic oxide).

This compound is such a strong oxidant that is able to oxidize Mn 2+ to permanganate, MnO4where manganese has o.n. +7.

2Mn(OH)2 + 5PbO2 + 8H3O+ $\dpi{120}&space;\rightleftharpoons$ 2MnO4- + 5Pb2+ + 14H2O

Our goal is the oxidation to permanganate ion, because it is extremely easy to identificate this compound; the intense violet color of its solutions should give you an idea!

The oxidation of manganese (II) to manganese (VII) is justified by the respective standard reduction potentials of the two species involved:

MnO4- / Mn2+ = 1,49V       ;         E°PbO2 / Pb2+ = 1.75 V

Cobalt identification

We have already seen that the identification of cobalt and nickel is performed directly on the ammonia solution coming from the 3rd group analysis.

- Vogel reaction

On a quote of the solution  from group 3 we add a few drops of 2N HCl to restore the cobalt ion (otherwise in the form of amino complex). We then add 2-3 mL of a mixture of amyl alcohol/ether and finally 1-2 spatula of  ammonium thiocyanate, NH4SCN. Shaking the tube weshould get a colored organic phase (ether) in blue, more or less intense.

The blue coloration is due to the formation of ammonium tetrathiocyanatocobaltate , a compound who has a good coefficient of distribution for the organic phase but is not particularly stable, hence we need to work using an excess of reagent.

The stoichiometry of the reaction explains why the complex is observable in the organic phase: H2O moves the equilibrium to the dissociation of the complex.

Co2+ + 4SCN- + 2NH4+ $\dpi{120}&space;\rightleftharpoons$ (NH4)2[Co(SCN)4]

or (better): Co(H2O)62+ + 4SCN- $\dpi{120}&space;\rightleftharpoons$ [Co(SCN)4]2- + 6H2O

Nickel forms  a very similar complex but less stable and with a less favorable ratio for the organic layer. The complex Ni(SCN)42- is colored in green and it is possible then to get as result of the reaction ia blue colored organic layer and a green colored aqueous layer.

Fe 3+ might interfere with this test. It could indeed form a red precipitate with thiocyanate (have a look at iron

Since the complex which iron forms with thiocyanate is more stable than the cobalt one, we will get this preferentially. That's why before doing this identification test is a good practice to add a spatula of NaF to the solution. Fluorine act as a maskerating agent for iron (III), which could result from mistakes made during group 3 analysis.

- Fisher reaction

To the same solution (ammoniacal from group 3) is added with CH3COOH, to increase the pH to 5. It is then added with an excess of solid potassium nitrite KNO2 .

The pH must be acidic, but not too much! This compound can undergo decomposition both in acidic and basic environment:

-in acidic medium:  [Co(NO2)6]3- + 10H+ $\dpi{120}&space;\rightleftharpoons$ 2Co2+ + 5NO + 7NO2 + 5H2O

-in basic medium:   [Co(NO2)6]3- + 3OH- $\dpi{120}&space;\rightleftharpoons$ Co(OH)3+ 6NO2-

In this reaction the nitrite performs two activities at the same time; oxidizes cobalt (II) to cobalt (III) and complexes cobalt (III) to give a hexacoordinated species.

The stoichiometry of the reaction explains the need for a large excess of reagent (7NO2-). The formation of the Fisher salt (cobaltinitrite) is quite slow, and the inner walls of the test tube should be rubbed with the glass rod (do you remember lead chloride?). Furthermore, it is recommended to heat in order to increase the speed.

A 'trick to increase the yield of the reaction is the addition of a soluble salt of potassium, such as KCl or KI, which provides quantitatively K+ ion. Considering the common ion effect, we will certainly encourage the cobaltinitrite formation.