Potassium permanganate is well known as an excellent oxidizing
agent. This is due in large part to the fact that manganese has
several oxidation states, the largest of which (+7) is present in
potassium permanganate. Manganese is often reduced to manganese(IV)
(as in manganese (IV) oxide)
or a soluble manganese(II) salt. Here are some findings on the
benefits and drawbacks of using potassium permanganate as an oxidizing
agent for a variety of compounds.
|Namely, the consideration was given towards
copper(I) chloride. Although this compound is green when in
solution, it was decided to see what happens since it is possible to
oxidize copper(I) to copper(II). This procedure was halted fairly
quickly when it was observed that the acidification of the copper(I)
solution oxidizes it to copper(II) (as noted by the light blue color of
solution) before any potassium permanganate was added. In truth,
the color of both copper(I) and copper(II) would have probably been
difficult to overcome.
|Most redox titration labs out there utilize an
iron(II) salt, frequently
iron(II) sulfate heptahydrate. This compound is used for two
reasons. First, it has a large molar mass, making it a good (but
not the best) primary standard. Second, its solution is near
colorless, so the pink/purple endpoint is easy to detect. Color
proves to be an issue when dealing with many other compounds.
Iron(II) salts are an excellent choice for a titration utilizing
|While lead(II) can oxidize to lead(IV), this is
not a feasible reaction for a number of reasons. Firstly, most
lead(II) salts are insoluble in water; truly only
lead(II) nitrate can
be considered. Secondly, the acidification of the solution will
immediately produce a precipitate. If
sulfuric acid is used,
lead(II) sulfate will begin to precipitate from the solution.
Hydrochloric acid will not produce much insoluble
lead(II) chloride, in
particular when it is heated, but the acid will react with the potassium
permanganate yielding chlorine gas. Consequently, lead(II) is not
|In the oxalate ion, carbon has a +3 charge.
When reacted with potassium permanganate the carbon is oxidized to the +4
ion in the form of carbon dioxide gas. Titrations involving the
oxalates are a little more involved, but the results were very good.
The reaction below describes the overall redox reaction:
|In the trial run performed, 0.51 g of
sodium oxalate was titrated against a 0.1M potassium
permanganate solution. The potassium permanganate was a premade
solution obtained from a supplier. You must gently heat the sodium
oxalate solution (most sources say between 55°C
and 60°C) before beginning the titration.
14.80 mL (0.00148 mol) of potassium permanganate was needed to reach the
endpoint. Stoichiometrically, 0.0015(22) mol should be needed to
fully titrate the original amount of oxalate. From this, it was
verified that sodium oxalate is excellent for use in a redox titration
with potassium permanganate.
|Tin(II) chloride dihydrate
|At first, one might think
tin(II) chloride dihydrate
would work well as a reducing agent for potassium permanganate.
After all, tin(II) can be oxidized to tin(IV). However, there are
several problems that arise. First, tin(II) chloride dihydrate,
while soluble in water, establishes an equilibrium with Sn(OH)Cl:
|Addition of chloride ions should shift the
equilibrium to the left and maintain a colorless solution of tin(II)
ions. Several sources confirmed that this was possible, but in
practice it is not so simple. Firstly, hydrochloric acid cannot be
used to acidify the solution and provide chloride ions. This is
due to the acid's willingness to react with potassium permanganate to
produce chlorine gas. To avoid this, sodium and potassium chloride
were used. After the addition of copious amounts of each salt, the
solution remained cloudy white. After performing several trials
the correct stoichiometric ratio of tin(II) needed to react with
permanganate was never verified. In the first trial, 0.66 g of
tin(II) chloride dihydrate was massed, an equivalent to 0.0029(2) mol.
9.32 mL (0.000932 mol) of potassium permanganate was needed to reach the
endpoint. Stoichiometrically, 0.0011(6) mol should have been
needed to fully titrate the tin. In the second trial, 0.52 g of
tin(II) chloride dihydrate was massed, an equivalent to 0.0023(0) mol.
7.96 mL (0.000796 mol) of potassium permanganate was needed to reach the
endpoint. Stoichiometrically, 0.00092(15) mol should have been
needed to fully titrate the tin. From this data, it was concluded
that tin(II) chloride dihydrate was not a good reducing agent. It
is assumed that this is because of a certain quantity of SnOCl that
refused to reverse back to tin(II) chloride.