Kinetics of the Oxidation of Iron (II) Sulfate by Nitric Acid at Constant and Varying Ionic Strength

Oxidation of iron (II) ions to iron (III) by nitric acid produces the complex ion [Fe(H2O)5NO]+2 which gives the characteristic brown color to the ring test for nitrates and nitrites. With sufficient excess of nitric acid, the complex is ultimately destroyed, and a clear, light yellow solution of the hydrated iron (III) ions in acid medium remains. The color change which accompanies the final destruction of the ferrous nitrosyl complex is abrupt, and hence, this is classified as a clock reaction. The time which elapses from the addition of iron (II) sulfate solution to the nitric acid and until the complex is destroyed may be taken as a measure of the rate of reaction. As with all nitric acid oxidations, the reaction is complex, and the mechanism is not fully understood. The original work on the kinetic study of this reaction as a clock reaction was done by Dr. John G. Surak and Mr. Duane Kobs of Marquette University at the suggestion of Dr. Joe Savolainen, Oak Ridge National Laboratory, Oak Ridge, Tennessee. The study was further advanced by Miss Betty Behl, Marquette University, who studied the reaction at various temperatures. This research is an investigation of the effects of varying the concentration of the nitrate ion and of the iron (II) ion on the rate of reaction. The purpose of this present paper has been to obtain more precise data under conditions already studied, and further, to study the effect of constant initial ionic strength on the rate of reaction. This problem also included the effect of various acids on the reaction rate. The ultimate purpose of the study of the ferrous nitrosyl reaction has been to gain some insight into the mechanism of the oxidation of iron (II) by nitric acid