One of the our chemistry practicals involve finding the molarity of an oxalic acid solution by titrating it with standard $\ce{KMnO4}$ solution.
The oxalic acid solution is heated and little sulfuric acid is added to titration flask before starting the titration.
It was my personal observation that adding the $\ce{KMnO4}$ solution rapidly causes a temporary appearance of a light yellow colour. Slow and gradual addition of $\ce{KMnO4}$ does not lead to this observation.
If sulfuric acid is not added to the titration flask, then also a yellowish tinge appears as we get closer to the end point.
What is the for this yellowish tinge? Is there any intermediate species formed (maybe a complex one) that leads to this observation?
Answer
Without knowing the concentrations of the involved substances and the values of some other parameters, we can only guess what has happened during your experiment.
The intended analytical reaction is the reduction of purple permanganate to colourless $\ce{Mn^2+}$:
$$\ce{MnO4- + 8H+ + 5e- <=> Mn^2+ + 4H2O}$$
If you add the permanganate solution too quickly, the reduction may be incomplete because not enough reducing agent (here: oxalic acid) or acid $(\ce{H+})$ is available in the reaction zone:
$$\ce{MnO4- + 4H+ + 3e- <=> MnO2 + 2H2O}$$
The observed colour may be caused by a cloud of finely dispersed particles of dark brown manganese(IV) oxide $(\ce{MnO2})$, which dissolves when the solution is well mixed again so that the intended reaction continues.
The same reaction may occur when you don’t add enough sulfuric acid, since $\ce{H+}$ is consumed during the reaction (8 mol $\ce{H+}$ per 1 mol $\ce{MnO4-}$).
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