Eh-pH diagram

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[[Image:Eh-pH diagram of Fe.png]]
[[Image:Eh-pH diagram of Fe.png]]
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== Why using Eh-pH diagrams? ==
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== How to use Eh-pH diagrams - an example ==
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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.
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.

Revision as of 12:18, 24 August 2007

Image:Eh-pH diagram of Fe.png

Why using Eh-pH diagrams?

xxxx xxx xxcx

How to use Eh-pH diagrams - an example

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.

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