Iodide precipitation titration

To about 0-5 g add 25 ml of O IN silver nitrate. After the addition of 10 ml of dilute nitric acid and heating on a water-bath for thirty minutes the salt is completely decomposed and the silver iodide precipitated. Titrate the excess of silver nitrate with 0 1 N potassium thiocyanate solution, using ferric alum as indicator. 1 ml O IN = 0 01269 g I. 

Active Oxygen Determinations. A sample (ca. 2.5 grams) was removed from the solution, weighed to 0.05 gram, and diluted with about 20 ml. of pentane. Nitrogen was blown over the solution for about 1 minute. The solution was diluted with 100 ml. of isopropyl alcohol 2 ml. of glacial acetic acid and 1 ml. of saturated potassium iodide solution were then added in that order. Enough water (5-10 ml.) was added to dissolve the potassium iodide precipitate. This mixture was titrated with standard sodium thiosulfate (0.01N) to a colorless endpoint. If the mixture were not titrated immediately, a piece of dry ice was added, and the solution was stored in the dark. 

Precipitation titrations can be extended to mixtures that form precipitates of different solubilities. The titration curve in Fig. 2 for a chloride/iodide mixture is a composite of the individual curves for the two anionic species. Because silver iodide has a much lower solubility than does silver chloride, the initial additions of the reagent result exclusively in formation of iodide. Thus, two equivalence points are evident. 

Curve A in Figure 13-6, which is the titration curve for the chloride/iodide mixture just considered, is a composite of the individual curves for the two anionic species. Two equivalence points are evident. Curve B is the titration curve for a mixture of bromide and chloride ions. Clearly, the change associated with the first equivalence point becomes less distinct as the solubilities of the two precipitates approach one another. In the bromide/chloride titration, the initial pAg values are lower than they are in the iodide/chloride titration because the solubility of silver bromide exceeds that of silver iodide. Beyond the first equivalence point, however, where chloride ion is being titrated, the two titration curves are identical. 

When the titration starts, Agl is precipitated, because it is much less soluble than AgBr. The initial part of the titration is therefore described by (5.6-4). As more silver nitrate is added, more silver iodide precipitates, until almost all iodide has been precipitated, at which point the silver concentration increases rapidly. At a given moment (which in this case will occur before the titration curve reaches its first equivalence point) silver bromide starts to... 

The potentiometric detection of the endpoint of precipitation titrations is very often used because not many visual indicators are available, in particular when mixtures of analytes are titrated. Halides, cyanide, sulfide, chromate, mercaptans, and thiols can be titrated with silver nitrate, using the silver sulfide-based ISE. Also complex mixtures, such as sulfide, thiocyanide, and chloride ions, or chloride, bromide, and iodide ions, can be titrated potentio-metrically with silver(I) ions. When the solubility of a compound formed during titration is too high, nonaqueous or mixed solvents are used, for example,... 

According to Hllllbrand and Lundell (H3) the best of the volumetric methods for lead Is precipitation of lead chromate from acetic acid solution, dissolution of the washed precipitate In hydrochloric acid, addition of excess potassium Iodide and titration of the liberated Iodine with a standard sodium thiosulfate solution. This procedure has recently been reviewed by Ryazanov (R2). 

Iodide ions can be determined quite well with argentimetric titrations. The Volhard method, electrometric endpoint indications, and adsorption indicators work well. The Mohr endpoint indication, however, does not give good results because of the adsorption of the chromate on the silver iodide precipitate. The presence of chloride and bromide ions disturbs the argentimeric iodide determination. 

The formation of stable soluble complexes has an analytical advantage over precipitation titration in that it avoids the problem of coprecipitation. In view of the rather unusual behavior of iodide ions in acetone, a short study of other inorganic anions which form insoluble silver salts in aqueous solutions was made for the acetone system. The results are shown in Table I. 

As Glass has observed, care must be taken that the end-point of the titration is decided upon with all the accuracy possible, since 0-1 ml of 0 1N solution affects the result by 0 66 per cent when using 0-2 g of iodoform. A gravimetric determination may be made by collecting the silver iodide precipitate and drying it in the usual manner, Agl X 0-5590 = CHI3. 

Dilute 50 ml of syrup to 400 ml, precipitate with sodium hydroxide and a little bromine water, to oxidise any iron in the ferrous state. Bring the mixture to the boil and allow the precipitate to settle. After filtration on paper in a Gooch crucible, dissolve the precipitate in hydrochloric acid, reprecipitate with sodium hydroxide and redissolve in hydrochloric acid. Add potassium iodide and titrate with 0 1 N thiosulphate. 1 ml = 0 005585 g Fe. 

For the determination of acetanilide, to an aliquot part of the filtrate from the iodide precipitation add sodium sulphite and sodium bicarbonate in slight excess add 2 drops of acetic anhydride and extract with three 60-ml portions of chloroform. Evaporate the chloroform to low bulk, add 10 ml of dilute sulphuric acid and evaporate the rest of the chloroform. Add 20 ml of water and digest on a water-bath for one hour. Add 10 ml of concentrated hydrochloric acid and titrate with 0 1 N potassium bromate-bromide, adding the volumetric solution very slowly to the well-shaken mixture until a faint yellow colour remains. Standardise the bromate-bromide solution against pure acetanilide. 

Precipitation titration Let us suppose our task is to determine to 0.1% the iodide content of an approximately 0.01 M I" solution. We choose for this a precipitation titration with silver ions. According to the tables, the Agl produced in the reaction has a solubility product of about 10 . We then calculate the 3p value ... 

Titration methods using adsorption indicators, based on the precipitation of insoluble iodides, have also been proposed (81—84). The sensitivity of these methods is less than that for the thiosulfate titration. Electrometric titration of the reaction between iodine and thiosulfate (85) was not found to be practicable for routine deterrninations of minute quantities of iodine. 

Analysis. The abiUty of silver ion to form sparingly soluble precipitates with many anions has been appHed to their quantitative deterrnination. Bromide, chloride, iodide, thiocyanate, and borate are determined by the titration of solutions containing these anions using standardized silver nitrate solutions in the presence of a suitable indicator. These titrations use fluorescein, tartrazine, rhodamine 6-G, and phenosafranine as indicators (50). 

Chlorinity When a sample of sea water is titrated with silver nitrate, bromides and iodides, as well as chlorides are precipitated. In calculating the chlorinity (Cl), the entire halogen content is taken as chloride, and chlorinity is defined as the weight in grams of silver required for precipitation of total halogen content per kilogram of sea water, multiplied by 0-328 533. (Chlorinity is always expressed as parts per thousand, using the symbol %o.)... 

The only difficulty in obtaining a sharp end point lies in the fact that silver cyanide, precipitated by local excess concentration of silver ion somewhat prior to the equivalence point, is very slow to re-dissolve and the titration is time-consuming. In the Deniges modification, iodide ion (usually as KI, ca 0.01 M) is used as the indicator and aqueous ammonia (ca 0.2M) is introduced to dissolve the silver cyanide. 

The method may be applied to those anions (e.g. chloride, bromide, and iodide) which are completely precipitated by silver and are sparingly soluble in dilute nitric acid. Excess of standard silver nitrate solution is added to the solution containing free nitric acid, and the residual silver nitrate solution is titrated with standard thiocyanate solution. This is sometimes termed the residual process. Anions whose silver salts are slightly soluble in water, but which are soluble in nitric acid, such as phosphate, arsenate, chromate, sulphide, and oxalate, may be precipitated in neutral solution with an excess of standard silver nitrate solution. The precipitate is filtered off, thoroughly washed, dissolved in dilute nitric acid, and the silver titrated with thiocyanate solution. Alternatively, the residual silver nitrate in the filtrate from the precipitation may be determined with thiocyanate solution after acidification with dilute nitric acid. 

Iodides can also be determined by this method, and in this case too there is no need to filter off the silver halide, since silver iodide is very much less soluble than silver thiocyanate. In this determination the iodide solution must be very dilute in order to reduce adsorption effects. The dilute iodide solution (ca 300 mL), acidified with dilute nitric acid, is treated very slowly and with vigorous stirring or shaking with standard 0.1 M silver nitrate until the yellow precipitate coagulates and the supernatant liquid appears colourless. Silver nitrate is then present in excess. One millilitre of iron(III) indicator solution is added, and the residual silver nitrate is titrated with standard 0.1M ammonium or potassium thiocyanate. 

Discussion. The theory of the titration of cyanides with silver nitrate solution has been given in Section 10.44. All silver salts except the sulphide are readily soluble in excess of a solution of an alkali cyanide, hence chloride, bromide, and iodide do not interfere. The only difficulty in obtaining a sharp end point lies in the fact that silver cyanide is often precipitated in a curdy form which does not readily re-dissolve, and, moreover, the end point is not easy to detect with accuracy. 

There are two methods for overcoming these disadvantages. In the first the precipitation of silver cyanoargentate at the end point can be avoided by the addition of ammonia solution, in which it is readily soluble, and if a little potassium iodide solution is added before the titration is commenced, sparingly soluble silver iodide, which is insoluble in ammonia solution, will be precipitated at the end point. The precipitation is best seen by viewing against a black background. 

Notes. (1) If in a similar determination, free mineral acid is present, a few drops of dilute sodium carbonate solution must be added until a faint permanent precipitate remains, and this is removed by means of a drop or two of acetic acid. The potassium iodide is then added and the titration continued. For accurate results, the solution should have a pH of 4-5.5. 

After the addition of the potassium iodide solution, run in standard 0.1M sodium thiosulphate until the brown colour of the iodine fades, then add 2 mL of starch solution, and continue the addition of the thiosulphate solution until the blue colour commences to fade. Then add about 1 g of potassium thiocyanate or ammonium thiocyanate, preferably as a 10 per cent aqueous solution the blue colour will instantly become more intense. Complete the titration as quickly as possible. The precipitate possesses a pale pink colour, and a distinct permanent end point is readily obtained. 

Discussion. Iodine (or tri-iodide ion Ij" = I2 +1-) is readily generated with 100 per cent efficiency by the oxidation of iodide ion at a platinum anode, and can be used for the coulometric titration of antimony (III). The optimum pH is between 7.5 and 8.5, and a complexing agent (e.g. tartrate ion) must be present to prevent hydrolysis and precipitation of the antimony. In solutions more alkaline than pH of about 8.5, disproportionation of iodine to iodide and iodate(I) (hypoiodite) occurs. The reversible character of the iodine-iodide complex renders equivalence point detection easy by both potentiometric and amperometric techniques for macro titrations, the usual visual detection of the end point with starch is possible. 

The indicator electrode must be reversible to one or the other of the ions which is being precipitated. Thus in the titration of a potassium iodide solution with standard silver nitrate solution, the electrode must be either a silver electrode or a platinum electrode in the presence of a little iodine (best introduced by adding a little of a freshly prepared alcoholic solution of iodine), i.e. an iodine electrode (reversible to I-). The exercise recommended is the standardisation of silver nitrate solution with pure sodium chloride. 

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