I read that after leaching of oxidized zinc concentrates with sulfuric acid, the resulting solution is purified from harmful impurities like iron, copper, cobalt, nickel. And with a weakly acidic medium, pH = 5.2–5.3 iron ions are practically precipitated by the formation of hydroxides.

Greenish Fe(OH)2 begins to precipitate at pH=7.5 (from 0.01M solutions), and rust-colored Fe(OH)3 begins to precipitate at pH=2 (from 0.01M solutions).

Under mildly acidic conditions pH .1.5 , Fe (iii) hydrolyses easily to ferric hydroxides where as Fe(ii) is very stable. For detail please ref : N.Mbedzi 217.

Ferrous hydroxide precipitates at pH above 8.5. in the presence of oxygen ferrous iron oxidizes to ferric iron and ferris hydroxide starts to precipitate at pH above 3.5

It will depend on solution strength, as solubility decreases with rising pH. At very low concentrations, the onset of precipitation will be delayed. As the prior notes indicate, Fe3+ precipitates first. This is described by Monhemius (1977). Direct message me and I can send the chart, but on a log conc - linear pH chart, there is a straight line relationship between about 10^-4 molar at pH 2.3 and 10^-2 molar at pH 1.7.

all of the above answers are partially good else Lyle Trytten . He concluded what I want to write in emphasis form. but above pH 2.22 vanadium may be precipitated also. if you want to precipitate iron any way

acidify with sulfuric acid and add 2-3 c.c. conc. nitric. cover the beaker and boil gently (10 min). and go on precipitation path.

Meysam Shahrashoub You need a (Pourbaix) Eh- pH diagram. See, for example, attached

Dear Alen,

They have no equal pH. One can precipitate at pH 3.5(Fe3+) and and another Fe+2 at pH 50 to 8. Differences for precipitations are there in operating condition and temperature which influences the precipitate formation.

The pH required to precipitate most metals from water ranges from pH 6 to 9 (except ferric iron which precipitates at about pH 3.5). In the case of acid water, the treatment could be supplemented by a correction of the pH. Thus, the ferrous iron is oxidized in ferric iron, which precipitates in ironhydroxide, Fe(OH)3. The precipitate is then separated from water by filtration on sand or decantation.

Iron (III) hydroxide precipitate formed by adding sodium hydroxide (NaOH) to a solution containing iron (III) ions. Iron (III) hydroxide (Fe(OH)3) is precipitated out of solution as a rust-brown gelatinous solid.

The most dominant form of dissolved iron is the soluble Fe+2 under the pH range of 5 to 8. Iron bacteria growth is very dependent upon the pH level, occurring over a range of 5.5 to 8.2 with 6.5 being the optimum level so it expected in all samples.

Systematic variations of the experimental variables revealed greater than 99% of the ferric iron can be removed from solution at conditions similar to those used in standard partial neutralisation in zinc and nickel production, pH of 2.5 and temperature less than 100 °C (less than 0.5%) with minimal losses of both nickel and cobalt. Temperature variation from 55 to 90 °C had no significant effect on the magnitude of Fe (III) precipitation but led to a significant increase in aluminium removal from 67% to 95% and improved the filterability of the precipitates. No ferrous iron precipitation even at a pH of 3.75 in the absence of an oxidant with its removal (98%) achieved by oxidative precipitation with oxygen gas at pH 3.5. Unlike Fe (III) precipitation, the operating temperature significantly affects oxidative precipitation of Fe (II). Hence, in practical application, the hydrolytic precipitation and oxidation to remove iron must be operated at 85 °C to ensure both ferrous and ferric iron are precipitated.

Hope it is helpful for you.

Ashish

Dear Meysam,

They have no equal pH. One can precipitate at pH 3.5(Fe3+) and and another Fe+2 at pH 50 to 8. Differences for precipitations are there in operating condition and temperature which influences the precipitate formation.

The pH required to precipitate most metals from water ranges from pH 6 to 9 (except ferric iron which precipitates at about pH 3.5). In the case of acid water, the treatment could be supplemented by a correction of the pH. Thus, the ferrous iron is oxidized in ferric iron, which precipitates in ironhydroxide, Fe(OH)3. The precipitate is then separated from water by filtration on sand or decantation.

Iron (III) hydroxide precipitate formed by adding sodium hydroxide (NaOH) to a solution containing iron (III) ions. Iron (III) hydroxide (Fe(OH)3) is precipitated out of solution as a rust-brown gelatinous solid.

The most dominant form of dissolved iron is the soluble Fe+2 under the pH range of 5 to 8. Iron bacteria growth is very dependent upon the pH level, occurring over a range of 5.5 to 8.2 with 6.5 being the optimum level so it expected in all samples.

Systematic variations of the experimental variables revealed greater than 99% of the ferric iron can be removed from solution at conditions similar to those used in standard partial neutralisation in zinc and nickel production, pH of 2.5 and temperature less than 100 °C (less than 0.5%) with minimal losses of both nickel and cobalt. Temperature variation from 55 to 90 °C had no significant effect on the magnitude of Fe (III) precipitation but led to a significant increase in aluminium removal from 67% to 95% and improved the filterability of the precipitates. No ferrous iron precipitation even at a pH of 3.75 in the absence of an oxidant with its removal (98%) achieved by oxidative precipitation with oxygen gas at pH 3.5. Unlike Fe (III) precipitation, the operating temperature significantly affects oxidative precipitation of Fe (II). Hence, in practical application, the hydrolytic precipitation and oxidation to remove iron must be operated at 85 °C to ensure both ferrous and ferric iron are precipitated.

Hope it is helpful for you.

Ashish