Iron , titrant

Another reducing titrant is ferrous ammonium sulfate, Fe(NH4)2(S04)2 6H2O, in which iron is present in the +2 oxidation state. Solutions of Fe + are normally very susceptible to air oxidation, but when prepared in 0.5 M 1T2S04 the solution may remain stable for as long as a month. Periodic restandardization with K2Cr20y is advisable. The titrant can be used in either a direct titration in which the Fe + is oxidized to Fe +, or an excess of the solution can be added and the quantity of Fe + produced determined by a back titration using a standard solution of Ce + or... 

The principle of coulometric titration. This involves the generation of a titrant by electrolysis and may be illustrated by reference to the titration of iron(II) with electro-generated cerium(IV), A large excess of Ce(III) is added to the solution containing the Fe(II) ion in the presence of, say IM sulphuric acid. Consider what happens at a platinum anode when a solution containing Fe(II) ions alone is electrolysed at constant current. Initially the reaction... 

STRATEGY (a) To obtain the amount of iron(II) in the analyte, we use the volume and concentration of the titrant. We follow the first two steps of the procedure in Toolbox L.2. Then we convert moles of Fe2+ ions into mass by using the molar mass of Fe2+ because the mass of electrons is so small, we use the molar mass of elemental iron for the molar mass of iron(II) ions, (b) We divide the mass of iron by the mass of the ore sample and multiply by 100%. 

Redox titrants (mainly in acetic acid) are bromine, iodine monochloride, chlorine dioxide, iodine (for Karl Fischer reagent based on a methanolic solution of iodine and S02 with pyridine, and the alternatives, methyl-Cellosolve instead of methanol, or sodium acetate instead of pyridine (see pp. 204-205), and other oxidants, mostly compounds of metals of high valency such as potassium permanganate, chromic acid, lead(IV) or mercury(II) acetate or cerium(IV) salts reductants include sodium dithionate, pyrocatechol and oxalic acid, and compounds of metals at low valency such as iron(II) perchlorate, tin(II) chloride, vanadyl acetate, arsenic(IV) or titanium(III) chloride and chromium(II) chloride. 

QA First, determine the mass of iron that has reacted as Fe2+ with the titrant. The balanced chemical equation provides the essential conversion factor. 

Potassium permanganate can be standardized by titration of sodium oxalate (Na2C204) by Reaction 7-1 or pure electrolytic iron wire. Dissolve dry (105°C, 2 h) sodium oxalate (available in a 99.9-99.95% pure form) in 1 M H2S04 and treat it with 90-95% of the required KMn04 solution at room temperature. Then warm the solution to 55-60°C and complete the titration by slow addition of KMn04. Subtract a blank value to account for the quantity of titrant (usually one drop) needed to impart a pink color to the solution. 

The basic approach in coulometric titrations is to generate electrochemically (at constant current) a titrant in solution that subsequently reacts by a secondary chemical reaction with the species to be determined. For example, a large excess of cerium(III) is placed in the solution together with an iron(II) sample. When a constant current is applied, the cerium(III) is oxidized at the anode to produce cerium(IV), which subsequently reacts with the iion(II) ... 

Should any iron(II) reach the anode, it also would be oxidized and thus not require the chemical reaction of Eq. (4.13) to bring about oxidation, but this would not in any way cause an error in the titration. This method is equivalent to the constant-rate addition of titrants from a burette. However, in place of a burette the titrant is electrochemically generated in the solution at a constant rate that is directly proportional to the constant current. For accurate results to be obtained the electrode reaction must occur with 100% current efficiency (i.e., without any side reactions that involve solvent or other materials that would not be effective in the secondary reaction). In the method of coulometric titrations the material that chemically reacts with the sample system is referred to as an electrochemical intermediate [the cerium(III)/cerium(IV) couple is the electrochemical intermediate for the titration of iron(II)]. Because one faraday of electrolysis current is equivalent to one gram-equivalent (g-equiv) of titrant, the coulometric titration method is extremely sensitive relative to conventional titration procedures. This becomes obvious when it is recognized that there are 96,485 coulombs (C) per faraday. Thus, 1 mA of current flowing for 1 second represents approximately 10-8 g-equiv of titrant. 

The cell potential is followed as the Fe2+ is added in small increments. Once the first drop of titrant has been added, the potential of the left cell is controlled by the ratio of oxidized and reduced iron according to the Nernst equation... 

Many samples have redox potentials such fiiat fiiey can be oxidized by iodine. Therefore, file iodine in file titrant may be consumed by readily oxidizable samples fiiat will give a false high value for file water content. Some common substances fiiat can be oxidized by iodine are ascorbic acid, arsenite (As02 ), arsenate (As04 ), boric acid, tetraborate (3407 ), carbonate (COs ), disulfite (8205 ), iron(ll) salts, hydrazine derivatives, hydroxides (OH ), bicarbonates (HCOs"), copper(l) salts, mercaptans (RSH), nitrite (N02 ), some metal oxides, peroxides, selenite (SeOs "), silanols (RsSiOH), sulfite (SOs ), tellurite (TeOs ), fiiiosulfate (8203 ), and tin(ll) salts. For situations such as fiiese where file material under analysis reacts wifii iodine, an oven can be used to liberate fiie moisture from file sample, which is fiieii carried into file reaction vessel and titrated wifiiout interference. 

The data in the third column of Table 19-2 are plotted as curve B in Figure 19-3 to compare the two titrations. The two curves are identical for volumes greater than 25.10 mL because the concentrati ons of the two cerium species are identical in this region. It is also interesting that the curve for iron(Il) is symmetric around the equivalence point, but the curve for uranium(IV) is not. In general, redox titration curves are symmetric when the analyte and titrant react in a 1 1 molar ratio. 

A solution prepared by dissolving a 0.2256-g sample of electrolytic iron wire in acid was passed through a Jones reductor. The iron(Il) in the resulting solution required a 35.37-mL titration. Calculate the molar oxidant concentration if the titrant used was (a) Ce (product Ce +). 

All the methods of end point detection discussed in the previous paragraphs are based on the assumption that the titration curve is symmetrical about the equivalence point and that the inflection in the curve coiresponds to this point. This assumption is valid if the titrant and analyte react in a 1 1 ratio and if the electrode reaction is reversible. Many oxidation/reduction reactions, such as the reaction of iron(II) with permanganate, do not occur in equimolar fashion. Even so, such titration curves are often so steep at the end point that vei little error is introduced by assuming that the curves are symmetrical. 

For routine titrations, it is often convenient to calculate the titer of the titrant. The titer is the weight of Malyte that is chemically equivalent to 1 mL of the titrant, usually expressed in milligrams. For example, if a potassium dichromate solution has a titer of 1.267 mg Fe, each milliliter potassium dichromate will react with 1.267 mg iron, and the weight of iron titrated is obtained by simply multiplying the volume of titrant used by the titer. The titer can be expressed in terms of any form of the analyte desired, for example, milligrams FeO or Fe203. 

We detect the end point by adding iron(III) as a ferric alum (ferric ammonium sulfate), which forms a soluble red complex with the first excess of titrant ... 

If the precipitate, AgX, is less soluble than AgSCN, we do not have to remove the precipitate before titrating. Such is the case with I , Br", and SCN . In the case of I , we do not add the indicator until all the I is precipitated, since it would be oxidized by the iron(III). If the precipitate is more soluble than AgSCN, it will react with the titrant to give a high and diffuse end point. Such is the case with AgCl ... 

An iron alloy or ore is dissolved in HCl and the iron is then reduced from Fe(III) to Fe(n) with stannous chloride (SnCl2). The excess SnCla is oxidized by addition of HgCl2. The calomel formed (insoluble Hg2Cl2) does not react at an appreciable rate with the titrant. The Fe(II) is then titrated with a standard K2Cr207 solution to a diphenylamine sulfonate end point. 

Nicholson proposed a differential potentiometric tltrator involving two indicator electrodes for the automatic control of processes in industrial plants [35]. As can be seen from Fig. 7.20, the sample and reagent streams are split and led to two half-cells via capillary tubes adjusted to provide slightly different titrated fractions. The potential difference (AE) between the two indicator electrodes Is transmitted to a control and detection system (D) which regulates the flow of titrant in an automatic fashion by means of valve V, thereby maintaining the preselected AE between the two ends of the cell. The speed of titrant addition, reflected by the flow meter (M), is a measure of the sample composition. An evaluation of the instrument carried out by the titration of dichromate with iron(II) revealed that the conditions to be used must be carefully selected. Thus, stable electrode responses are only obtained in the zone where Fe(II) prevails, and not in that where dichromate prevails over the former as the process determining the potential obtained in such a zone is irreversible. This method therefore has limited application in the control of slow reactions. 

In a back-titration of silver by chloride ions, potassium chromate can be used to indicate the endpoint (Mohr s method) but with potassium thiocyanate as titrant, ammonium iron(III) sulfate ( ferric alum ) is preferred. In the direct titration (Gay-Lussac s method) the location of the turbidimetric endpoint has been improved in detail. ... 

Titrations based on oxidation-reduction reactions enjoy wide use. Permanganate, dichromate, and iodine and iron(II), tin(II), thiosulfate, and oxalate are commonly used oxidizing and reducing titrants, have been employed to determine components in both inorganic and organic analysis. As we saw in Chapter 7, solvent water does not play as central a role as in acid-base titrations. Oxidants or reductants strong enough to decompose water are not practical as titrants. 

Some of the most successful and widely used chelating reagents include dimethylglyoxime for the gravimetric determination of nickel 1,10-phe-nanthroline and its derivatives for the colorimetric determination of iron and copper dithizone for the separation and colorimetric determination of a number of metals but particularly lead, silver, zinc, cadmium, and mercury the dithiocarbamates such as diethylammonium diethyldithiocarbamate and ammonium pyrrolidinedithiocarbamate, used for colorimetry but more widely applied now as selective extractants and the most successful titrant, EDTA. 

The logs of the activities have been used in Figure 12-2 to produce the iron diagram. It shows clearly that attempts to obtain cell readings for Fe°-Fe(III) will fail because Fe(II) must first be produced in far higher activity than the Fe(III) to reach equilibrium. It shows that Fe(II) will be lowered below 1/1000 of the Fe(III) if an oxidizing titrant raises the E value of the solution above 1.0 V. It predicts suitable conditions for study of the equilibrium... 

In coulometric titrations, a constant current generates the titrant electrolyticaliy. In some analyses, the active electrode process involves only generation of the reagent." An example is the titration of halides by silver ions produced at a silver anode. In other titrations, the analyte may also be directly involved at the generator electrode. An example of this type of titration is the coulometric oxidation of iron(II)—in part by elec-trolytically generated ceriiun(IV) and in part by direct electrode reaction (Section 24B-2). Under any circumstance, the net process must approach 100% current efficiency with respect to a single chemical change in the analyte. 

General methods for the estimation of bismuth present in quantity include direct weighing as sulphide precipitation as sulphide, conversion to carbonate and ignition to oxide and precipitation as phosphate. Recently the use of EDTA as a titrant for bismuth has considerably simplified the determination of this metal in pharmaceutical mixtures. Bismuth forms a complex at pH 1 to 2 and at this degree of acidity few other metals likely to be encountered (with the exception of iron) interfere bismuth may therefore be selectively titrated in the presence of, say, aluminium, magnesium, or calcium. Various methods available have been discussed by Brookes and Johnson and the following general procedure is recommended ... 

The classical gravimetric determination of iron now finds little application in pharmaceutical work, partly because of the readiness with which other ions are adsorbed on to the precipitate but principally because of the variety of titrimetric methods which are available. Oxidising titrants are the most widely used and these may be applied directly to ferrous iron or, after suitable reduction, to ferric iron. Reducing titrants also find some application for the direct titration of ferric iron. Chelating titrants such as EDTA may be used but, because of the formation of a highly coloured complex and because other rapid titrimetric methods are already available, these are unlikely to find routine application. 

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