Reinhardt-Zimmermann solution

The Reinhardt-Zimmermann solution is a mixture of sulfuric acid , manganese (II) sulfate and phosphoric acid in water used in analytical chemistry . It is used for the redox titration of samples containing iron (II) ions with potassium permanganate . It is named after its inventors Julius Clemens Zimmermann and C. Reinhardt . The titration of iron-containing samples with permanganate under hydrochloric acid would not be possible without the Reinhardt-Zimmermann solution.

Problem

If a solution containing iron (II) ions is titrated with potassium permanganate and sulfuric acid, iron (II) is, as expected, oxidized to iron (III):

{\ displaystyle \ mathrm {MnO_ {4} ^ {\, -} + 5 \ Fe ^ {2 +} + 8 \ H_ {3} O ^ {+} \ \ longrightarrow \ Mn ^ {2 +} + 5 \ Fe ^ {3 +} + 12 \ H_ {2} O}}

However, if you want to use hydrochloric acid instead of sulfuric acid, some problems arise:

The titration takes longer because the end point is difficult to determine
The smell of chlorine gas can be heard during the titration

When comparing the amount of substance, it is noticeable that when hydrochloric acid is used instead of sulfuric acid, there is a significant increase in the consumption of permanganate solution in order to titrate the same amount of iron (II) ions. Since mass and charge cannot be lost according to the principles of chemistry, it is obvious that the resulting chlorine gas was created by the oxidation of the chloride ions from the permanganate in the solution. For this purpose, the standard electrode potentials must be considered.

reaction	Standard electrode potential
${\ displaystyle \ mathrm {{Cl_ {2}} _ {(g)} + 2 \ e ^ {-} \ longrightarrow 2 \ Cl _ {(aq)} ^ {\, -}}}$	${\ displaystyle 1 {,} 36 \, \ mathrm {V}}$
${\ displaystyle \ mathrm {{MnO_ {4} ^ {\, -}} _ {(aq)} + 8 \ H_ {3} O _ {(aq)} ^ {+} + 5 \ e ^ {-} \ longrightarrow Mn _ {(aq)} ^ {2 +} + 12 \ H_ {2} O _ {(l)}}}$	${\ displaystyle 1 {,} 51 \, \ mathrm {V}}$

The voluntariness of electrochemical processes is described like many other processes with the free enthalpy .

{\ displaystyle \ Delta G_ {m} ^ {\ circ} = - z \ cdot F \ cdot E ^ {\ circ}}

The greater the standard reduction potential, the more voluntary the process:

{\ displaystyle E ^ {\ circ} (\ mathrm {MnO_ {4} ^ {\, -} / Mn ^ {2+}})> E ^ {\ circ} (\ mathrm {Cl_ {2} / Cl ^ {-}})}

Consequently, it can be assumed that the permanganate ions oxidize the chloride ions to chlorine gas.

In 1954, however, the German chemist Hans J. Schäfer discovered that dilute hydrochloric acid is not oxidized by permanganate ions, although this would be entirely possible from the redox potential. The reaction is apparently kinetically inhibited and has overvoltage . Since iron (III) ions have no influence on the reaction of permanganate with chloride, the cause can be traced back to the iron (II) ions.

Mode of action

Julius Clemens Zimmermann found that the addition of manganese (II) sulfate at the start of the titration minimized the disturbance of the reaction. Consequently, the presence of ferrous ions promotes oxidation while manganese ions prevent it. The kinetic background of the reaction seems more complicated than expected. At the beginning of the 20th century it was Zimmermann and Wilhelm Manchot (chemists) who investigated the background to this reaction. Today it is assumed that short-lived manganese (III) ions are formed during the reaction, which enable the induction of the oxidation of chloride ions to chlorine. The reduction potential of the system

{\ displaystyle \ mathrm {Mn ^ {2 +} \ rightleftharpoons Mn ^ {3 +} + e ^ {-}}}

which is 1.51 V is decisive here. Generally speaking, the high concentration of manganese (II) ions leads to the formation of manganese (III) ions, which lower the oxidation potential and accelerate the reaction. The oxidation of chloride is thus inhibited. In addition, the phosphoric acid complexes manganese (III) ions to form manganese (III) phosphate . Just like manganese, iron is complexed to form colorless iron (III) phosphate . The yellowish chloric acids of iron, which otherwise arise in acidic solution, such as, for example, would make the detection of the end point more difficult. By reducing the iron (III) ion concentration, the redox potential of iron is also reduced and the oxidation of the iron (II) ions is facilitated. ${\ displaystyle \ mathrm {H_ {2} FeCl_ {6}}}$

application

A hydrochloric acid solution diluted to 100 ml is mixed with 10 ml Reinhardt-Zimmermann solution and titrated. If the iron concentration is 0.1 mol / l and the acid concentration does not exceed 1 mol / l, the error can be kept at 0.2%. The titration should be carried out slowly.

Manufacturing

A mix of

1 liter of phosphoric acid (at a density of 1.3 g / cm³)
0.6 liters of water
0.4 liters of sulfuric acid (at a density of 1.84 g / cm³)

is added to a solution of 200 g of crystalline manganese (II) sulfate in 1 liter of water.

This creates 3l of the Reinhardt-Zimmermann solution.

Alternatives

Research has shown that other substances such as

can prevent the oxidation of the chloride ions to chlorine to some extent.

Individual evidence

↑ Gerhart Jander, Karl year: Measure analysis - theory and practice of titrations with chemical and physical indicators. 18th edition. De Gruyter, 2012, ISBN 978-3-11-024898-2 , pp. 165ff.
↑ All potentials are taken from: Michael Binnewies and others: Allgemeine und Anorganische Chemie. 1st edition. Spektrum Akademischer Verlag, Heidelberg, 2004, ISBN 3-8274-0208-5 .

[1] Gerhart Jander, Karl year: Measure analysis - theory and practice of titrations with chemical and physical indicators. 18th edition. De Gruyter, 2012, ISBN 978-3-11-024898-2 , pp. 165ff.

[2] All potentials are taken from: Michael Binnewies and others: Allgemeine und Anorganische Chemie. 1st edition. Spektrum Akademischer Verlag, Heidelberg, 2004, ISBN 3-8274-0208-5 .