Chemical precipitation

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Chemical precipitation is used for removal of both impurities and wanted metals from leaching solutions. When the desired metal is precipitated it is usually treated further before a desired product, metal or metal salt, is obtained. Precipitation is achieved by the addition of proper reagents (alkali, salts) and can under correct conditions be done selectively. Examples of chemical precipitations commonly used are hydroxide- and sulphide precipitation.

Formation of a solid substance that can be separated from a solution. This is achieved by converting the soluble metal ions into an insoluble form by changing the composition of the aqueous solution so that the solubility of the metal ions is reduced.

Chemical precipitation is used to remove dissolved metals from solutions, such as process wastewaters containing toxic metals. Soluble metal ions are converted to insoluble form by chemical reaction with a precipitating reagent. The solid particles formed by this reaction are removed from solution by settling and/or filtration.

The effectiveness of a chemical precipitation process is dependent on several factors which include:

  • type and concentration of metal ions present in solution
  • precipitant used
  • reaction conditions (pH, temperature, redox potential, etc.)
  • presence of other constituents like complexing agents that may inhibit the precipitation reaction.

Metal ions such as zinc can be precipitated from a solution by addition of sodium hydroxide:

Zn SO4 (aq) + 2NaOH (aq) → Zn(OH)2 (s) + Na2SO4 (aq)

Hydroxide precipitation

The precipitation of metal hydroxides such as for example Al(OH)3, Fe(OH)3 and Zn(OH)2 is the most widely used chemical precipitation process in hydrometallurgy. The precipitation is achieved by increasing pH in the solution by the addition of alkaline reagents such as limestone (CaCO3), slaked lime (Ca(OH)2), sodium hydroxide (NaOH) or ammonia (NH3).

Solubility diagram for some metal hydroxides
Solubility diagram for some metal hydroxides

To give an indication on the conditions needed to separate different elements by hydroxide precipitation solubility diagrams of metal hydroxides are helpful. In figure ?? a solubility diagram with some of the common metal hydroxides is shown.

In the figure the line for each metal ion represents the equilibrium between a 1 M solution of respective metal ion and its hydroxide. On the y-axis the metal ion concentration is given and pH is on the x-axis. At low pH-values the hydroxides are soluble and at higher pH-values the metal ion concentration decreases as the hydroxides are formed. It can be seen that trivalent metal ions can be precipitated at lower pH values compared with the divalent metal ions. Metal ions located close to each other like Fe2+ and Zn2+ are difficult to separate and give mixed hydroxide precipitates if both are present during neutralisation. The most difficult metals to precipitate in the diagram are the alkaline earth metals calcium and magnesium.

Another even more powerful tool for prediction of metal ion behaviour in aqueous solutions is the Eh-pH (Porbaix) diagrams. Image:Eh-pH diagram of Fe.png

In the diagram the dotted line represents the stability area of water, i.e. the area of relevance for hydrometallurgy. The upper line is the equilibrium between water and oxygen gas. If the solution potential is increased above the dotted line water is oxidised forming oxygen and hydrogen ions. The lower line shows where hydrogen ions in solution are reduced to hydrogen gas.

From the figure it can be seen that metallic iron (Fe(s)) is not stable in water since iron is located below the stability region of water. This is an example of the well known fact that iron corrodes in aqueous solution. If iron metal is placed in water it is oxidised to Fe2+ and if pH is higher than 5-6 the ferrous iron is precipitated as Fe(OH)2(s). It can also be seen that the only soluble iron species are ferrous iron (Fe2+) and ferric iron (Fe3+), ferric iron is only stable at a limited area at low pH-values and high redox potential while ferrous iron is stable over a wider region at lower redox potential and extending to a higher pH-values. In conditions where the iron containing solution is in contact with the atmosphere the iron is further oxidised to ferric iron by the oxygen in the air forming a brownish red precipitate at pH values above 1-2.

The normal procedure to remove iron from solution is to precipitate it as a ferric hydroxide (Fe(OH)3). If iron exists as Fe3+ in an acidic leaching solution precipitation is affected by additions of alkali like lime (CaO) or limestone (CaCO3) to increase the pH. In the case when iron is in the Fe2+ state precipitation is usually done by first oxidising it to the ferric form followed by a pH increase. In practice iron removal is usually achieved by maintaining pH at a level of 2-3 and at the same time air is blown into the tank to oxidise ferrous iron to ferric.

In the following figures examples of Eh-pH diagrams for copper, cobalt and nickel are given.

Examples of Eh-pH diagrams:

Chemical precipitation can be used for either solution purification or metal recovery. Eh-pH diagrams is a good tool to visualize the stability areas of metal species in solution depending on the solutions redox potential (Eh) and pH. These diagrams can be used to indicate how different metal ions can be separated from each other by making changes in these variables.

Sulphide precipitation

Precipitation of metal sulphides is also practised in hydrometallurgical operations especially for metals like copper, zinc, nickel and cobalt. Common reagents used for precipitation are H2S(g), NaHS and Na2S. In the figure below a solubility diagram of metal sulphides is shown. The scale of the x-axis is pS which is defined in the same way as pH where p means the negative logarithm, i.e. a pS of 10 is equal to a sulphide ion concentration of 10-10 M. From the figure it can be seen that metal ions of mercury, silver and copper are most easily precipitated and can be affected at extremely low concentrations of free sulphide ions. Image:Solubility of metal sulphides.png

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