Ferric thiosulfate

Williamson and Rimstidt (1993) reported the rate of decomposition of ferric thiosulfate to form ferrous iron and tetrathionate ions. The reaction brings two ferric thiosulfate molecules together into an activated complex, which decomposes to form two ferrous ions and a tetrathionate ion. 

Example 5.2. Finding the activation energy for the ferric thiosulfate reaction using the Arrhenius equation... 

Williamson and Rimstidt (1993) determined the rate of the ferric thiosulfate reaction at temperatures ranging from 12° to 34°C. When those data are plotted on a graph of log k versus /T (Figure 5.4) the equation for the best-fit straight line is... 

The very large activation energy of the ferric thiosulfate reaction means that a 5°C temperature increase, from 25° to 30°C, will cause the rate to double (Figure 5.3). 

Figure 5.5. Schematic diagram showing the steps in the ferric thiosulfate reaction. (a)Two FeSjOj molecules approaching each other. (b)The approximate geometry of the activated complex. (c)The reaction products, which are two ferrous iron ions and a tetrathionate ion.

Example 5.3. Testing the postulated stoichiometry of the ferric thiosulfate transition state using rate versus ionic strength data... 

In the previous two examples, the activated complex for the ferric thiosulfate reaction was postulated to be [(FeS203)2 ]. This molecule is a combination of two FeS203+ ions, each with a +1 charge, so the product Z is +1. If this pos-... 

Figure 5.6. The graph of log /rversus V/for the ferric thiosulfate reaction.This graph has a slope of 0.928 (= 1), which is consistent with each of the interacting molecules having a charge of +1.

The ferric thiosulfate reaction can be inhibited by the addition of Ag, which forms a very strong complex with thiosnlfate and that complex is urrreactive becanse Ag carmot be reduced to Ag by reaction with thiosnlfate or oxidized by reaction with Fe . ... 

Williamson, M.A., Rimstidt, XD. (1993). The rate of decomposition of the ferric-thiosulfate complex in acidic aqueous solutions. Geochimica et Cosmochimica Acta, 57, 3555-3561. 

Test by catalytic acceleration of the ferric-thiosulfate reaction ... 

Copper, Test by catalytic effect on the ferric-thiosulfate reaction, see page 205. 

The hberated iodine, as the complex triiodide ion, may be titrated with standard thiosulfate solution. A general iodometric assay method for organic peroxides has been pubUshed (253). Some peroxyesters may be determined by ferric ion-catalyzed iodometric analysis or by cupric ion catalysis. The latter has become an ASTM Standard procedure (254). Other reducing agents are ferrous, titanous, chromous, staimous, and arsenite ions triphenylphosphine diphenyl sulfide and triphenjiarsine (255,256). 

Analytical and Test Methods. An aqueous solution of sodium thiosulfate forms a white precipitate with hydrochloric acid and evolves sulfur dioxide gas which is detected by its characteristic odor. The white precipitate turns yellow, iadicatiug the presence of sulfur. The addition of ferric chloride to sodium thiosulfate solutions produces a dark violet color which quickly disappears. 

A number of methods have been proposed for the detection of rancidity. The determination of active oxygen consists of dissolving the fat in a suitable medium such as chloroform and acetic acid, adding potassium iodide, and titrating the liberated iodine with a standard thiosulfate solution (16, 20). This is perhaps the most widely used method at the present time. Another procedure which has been proposed for the detection of peroxides employs ferrous ammonium sulfate and ammonium thiocyanate in acetone. The resulting red color of ferric thiocyanate is measured spectrophotometrically, and is said by the authors to yield more reproducible results than do the usual titration methods (21). 

Redox (reduction-oxidation) titrimetry is used primarily for nitrate detns. Five systems are in current use ferrous sulfate—dichromate, io dome trie, periodic acid oxidation (NaOH titrant), K permanganate, and titanous chloride-ferric ammonium sulfate. The ferrous sulfate— dichromate system is used for MNT DNT detns (Vol 2, C162-Lff Vol 6, F17-Rff Ref 17). In the iodometric procedure, the sample (ie, NG) is treated in a C02 atm with a satd soln of Mn chloride in coned HC1, the vol reaction products are bubbled thru a K iodide soln, and the liberated iodine is titrated with standard thiosulfate soln (Refs 1 17). The periodic... 

Electrocatalysis in oxidation has apparently first been shown for ascorbic acid oxidation by Prussian blue [60] and later by nickel hexacyanoferrate [61]. More valuable for analytical applications was the discovery in the early 1990s of the oxidation of sulfite [62] and thiosulfate [18, 63] at nickel [62, 63] and also ferric, indium, and cobalt [18] hexacyanoferrates. More recently electrocatalytic activity in thiosulfate oxidation was shown also for zinc [23] hexacyanoferrate. Prussian blue-modified electrodes allowed sulfite determination in wine products [64], which is important for the wine industry. 

Starting with examples of Mechanism 10 in inorganic chemistry, one may cite the oxidation of cuprous thiosulfate by ferric, vanadate, molybdate, and chromate ions (38). 

Compound Name Ammonium Hydroxide Hexamethylenetetramine Ammonium Acetate Ammonium Bifluoride Ammonium Sulfamate Ammonium Sulfamate Ammonium Benzoate Ammonium Bicarbonate Ammonium Dichromate Ammonium Bifluoride Ammonium Carbonate Ammonium Chloride Ammonium Citrate Ammonium Citrate Ammonium Pentaborate Ammonium Dichromate Nickel Ammonium Sulfate Ferric Ammonium Citrate Ferric Ammonium Oxalate Ferrous Ammonium Sulfate Ammonium Fluoride Ammonium Silicofluoride Ammonium Formate Ammonium Gluconate Ammonium Bicarbonate Ammonium Bifluoride Ammonium Sulfide Ammonium Hydroxide Ammonium Thiosulfate Ammonium Thiosulfate Ammonium Iodide Ferrous Ammonium Sulfate Ammonium Lactate Ammonium Lactate Ammonium Lauryl Sulfate Ammonium Molybdate Ammonium Chloride Nickel Ammonium Sulfate Ammonium Nitrate Ammonium Nitrate-Urea Solution Ammonium Oleate... 

Supersaturation.—One of the most frequent difficulties in crystallization is due to super saturation. This condition arises when the normal saturation concentration of a salt solution is exceeded without the appearance of any crystals, and as the solution cools further it becomes steadily more supersaturated. When crystallization finally starts, it proceeds with great rapidity, forming a mass of poorly defined crystals unsuited to drying. The tendency toward supersaturation is most marked in the case of very soluble substances which form viscous or syrupy solutions. Lead acetate, sodium thiosulfate, ferric nitrate, and sulfuric acid are good examples. 

The dissolution rate of goethite by sulfide was found to increase with surface area and proton concentration. Pyzik and Sommer (21) suggested that HS" is the reactive species that reduces surface ferric iron after exchanging versus OH . A subsequent protonation of surface ferrous hydroxide would lead to dissolution of a surface layer. Elemental sulfur was the prominent oxidation product polysulfides and thiosulfate were found to a lower extent. The dissolution rate R (in moles per square meter per second) of hematite by sulfide was demonstrated to be proportional to the surface concentration of the surface complexes >FeHS and >FeS (22). 

Redox titrations find numerous applications in environmental analysis. Iodo-metric titration involving the reaction of iodine with a reducing agent such as thiosulfate or phenylarsine oxide of known strength is a typical example of a redox titration. This method is discussed separately in the next section. Another example of redox titration is the determination of sulfite, (S032-) using ferric ammonium sulfate, [NH4Fe(S04)2]. 

Applicable to <20 mg 1, in natural and waste water. Ammonium phosphate must be added to mask interferences of residual ferric chloride. The effect of other interferents (reducing agents and thiosulfate) may be important. 

Determination of the Hydroperoxide. Glacial acetic acid (30 ml.) containing 0.0005% ferric chloride hexahydrate and 5 ml. saturated potassium iodide solution was added to an aliquot of the hydroperoxide solution which had been flushed with nitrogen. The mixture was stored in the dark for 10 minutes, diluted with water (50 ml.), and the liberated iodine was titrated with sodium thiosulfate solution using starch as indicator. 

Iron-free ferrioxamine B ( Desferal", Ciba Pharmaceutical Company, Sum-mitt, New Jersey) reacts rapidly with ferrous ion, especially at neutral pH, forming the ferric chelate. This transition is blocked by mercaptoacetic acid, hydrosulfite and thiosulfate but not by weaker reducing agents such as ascorbic acid, hydroxylam-ine and sulfite (Nature 205, 281, 1965). 

Ferric Iron Dissolve about 2 g of sample, accurately weighed, in a mixture of 100 mL of water and 10 mL of hydrochloric acid contained in a 250-mL glass-stoppered flask, add 3 g of potassium iodide, shake well, and allow the mixture to stand in the dark for 5 min. Titrate any liberated iodine with 0.1 N sodium thiosulfate, using starch TS as the indicator. Perform a blank determination (see General Provisions), and make any necessary correction. Each milliliter of 0.1 N sodium thiosulfate is equivalent to 5.585 mg of ferric iron. 

Ferric Iron Transfer 2 g of sample into a 250-mL glass-stoppered Erlenmeyer flask, add 25 mL of water and 4 mL of hydrochloric acid, and heat on a hot plate until solution is complete. Stopper the flask, and cool to room temperature. Add 3 g of potassium iodide, stopper, swirl to mix, and allow to stand in the dark for 5 min. Remove the stopper, add 75 mL of water, and titrate with 0.1 A sodium thiosulfate, adding starch TS near the endpoint. Not more than 7.16 mL of 0.1 A sodium thiosulfate is consumed. 

With hydrochloric acid, solutions of thiosulfates yield a white precipitate that soon turns yellow, liberating sulfur dioxide, recognizable by its odor. The addition of ferric chloride TS... 

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