# Analysis of group 3 cations

Group 3 cations is also called the hydroxides group, because it is made up of cations which precipitate as hydroxides in ammonia alkaline solution. More specifically they precipitate around pH 9, the pH made with ammonia and ammonium chloride (NH3 / NH4Cl), a very common buffer solution.

NH3 + H2O $\dpi{120}&space;\rightleftharpoons$ NH4+ + OH-

NH4Cl $\dpi{120}&space;\rightleftharpoons$ NH4+ + Cl-

As you can see two equilibria are involved. NH4Cl has a threefold function:

1) By acting as a common ion in the equilibrium of NH3, it reduces the concentration of OH- ions, thus preventing the precipitation of undesirable hydroxides, such as Mg(OH).

2) Still by acting as a common ion, it ensures a great availability of NH3, which forms soluble amino complexes with species such as Zn2+, Co2+ and Ni2+ in which this cations coordinate 4 or 6 molecules of ammonia.

3) As electrolyte reduces the likelihood of colloidal precipitates formation.

Cations that may precipitate as 3 group hydroxydes are:

Manganese is actually analyzed as a group 4 cation and precipitates only partially as an hydroxide.

Some theory: hydroxides precipitation mechanism

We can easily predict when our hydroxides will start to precipitate and even when they will be quantitatively precipitated. We take 10-2 M as start of precipitation and 10-5 M as quantitative precipitation (we're obviously referring to the free cations in solution).

Take for example ferric hydroxide, Fe(OH)3.

$\dpi{120}&space;Fe^{3+}&space;+&space;3OH^-&space;\rightleftharpoons&space;Fe(OH)_3$

$\dpi{120}&space;Kps&space;=&space;[Fe^{3+}][OH^-]^{3}&space;=&space;10^{-35}$

$\dpi{120}&space;[OH^-]^3&space;=&space;\frac{kps}{[Fe^{3+}]}$ given, as we said, $\dpi{120}&space;\dpi{120}&space;[Fe^{3+}]&space;=&space;10^{-2}M$

we get that the concentration of $\dpi{120}&space;[OH^-]$ necessary to precipitate the ferric hydroxide is equal to:

$\dpi{120}&space;\dpi{120}&space;[OH^-]=&space;\sqrt[3]{\frac{10^{-35}}{[10^{-2}]}}&space;=&space;10^{-11}$ of course, we've already got a relationship with the pH:

$\dpi{120}&space;pH&space;=&space;14&space;-&space;pOH&space;=&space;14&space;-&space;(-log&space;[OH^-])&space;=&space;14&space;-&space;(log[10^{-11}])&space;=&space;14&space;-&space;11&space;=&space;3$

So, the ferric hydroxide starts its precipitation around pH 3. That's why we have not heard about it during group 1 and group 2 analysis (pH 0 and 2, respectively).

We can also calculate the pH value  which makes the precipitation quantitative ( we consider the precipitation quantitative when $\dpi{120}&space;[Fe^{3+}]$ in solution it is equal or beneath 10-5  ,our limit of detection).

$\dpi{120}&space;\dpi{120}&space;[OH^-]=&space;\sqrt[3]{\frac{10^{-35}}{[10^{-5}]}}&space;=&space;10^{-10}$

$\dpi{120}&space;\dpi{120}&space;\dpi{120}&space;\dpi{120}&space;\dpi{120}&space;pH&space;=&space;14&space;-&space;pOH&space;=&space;14&space;-&space;(log[10^{-10}])&space;=&space;4$

The precipitation is therefore complete at pH 4. The other hydroxides (aluminum and chromium based), are little more soluble and precipitate at slightly higher pH. Anyway, these three hydroxydes start to precipitate between pH 3 and 6, and once the solution gets (see below) to pH 9 they all precipitate quantitatively.