APCH 231 CHEMICAL ANALYSIS PRECIPITATION TITRATIONS Titrations based on reactions that produce sparingly soluble substances are referred to as precipitation titrations. They are limited in their scope because so many fail to meet either the stoichiometry or speed requirements for a successful titration. Coprecipitation often leads to nonstoichiometric reactions. The technique of digestion used to minimize coprecipitation in gravimetry cannot be applied to direct titrations because they require considerable amount of time to become effective. The rates of formation of some precipitates are quite slow, especially in the dilute solutions near the equivalence point. Only procedures using silver ion as the titrant or analyte have remained competitive with newer analytical methods. Precipitation Equilibria Inorganic solids which have limited water solubility show an equilibrium in solution represented by the so called solubility product. For example, AgCl dissolves slightly in water giving Cl - and Ag + where AgCl (s) Ag + + Cl - Keq = [Ag + ][Cl - ]/[AgCl (s) ] However, the concentration of a solid is constant and the equilibrium constant can include the concentration of the solid and thus is referred to, in this case, as the solubility product, K sp where K sp = [Ag + ][Cl - ] 43
It should be clear the product of the ions raised to appropriate power as the number of moles will fit in one of three cases: 1. When the product is less than K sp : No precipitate is formed and we have a clear solution. 2. When the product is equal to K sp : We have a saturated solution. 3. When the product exceeds the K sp : A precipitate will form. It should also be clear that at equilibrium of the solid with its ions, the concentration of each ion is constant and the precipitation of ions in solution does occur but at the same rate as the solubility of precipitate in solution. Therefore, the concentration of the ions remains constant at equilibrium. The same equilibrium concepts discussed earlier control equilibrium governing the behavior of solutions containing sparingly soluble substances. Two situations will be studied here: a. Solubility in pure water b. Solubility in presence of a common ion Calculating Solubility Solubility is defined as the concentration of a dissolved solute at equilibrium with its undissolved form. When a slightly soluble ionic salt dissolves, it produces cations and anions in the solution. If these ions do not undergo any reaction other than combining to re-form the salt, the solubility can be easily calculated from the solubility product expression. 44
a. Solubility in Pure Water The calculations involved in this type of equilibrium is straightforward. Question: 1. Calculate the molar solubility of silver bromide in water at 20 C. The solubility product at this temperature is 5.0 x10-13. 2. What must be the concentration of Ag + to just start precipitation of AgCl in a 1.0x10-3 M solution. 3. What ph is required to just start precipitation of Fe(III) hydroxide from a 0.1 M FeCl 3 solution. K sp = 2 x10-39. 4. Calculate the solubility, in g/100 ml, of calcium iodate in water at 20 C. b. Solubility in Presence of a Common Ion In the absence of any competing equilibria, a precipitate is less soluble in a solution containing an excess of one of the ions common to the precipitate than it in pure water. According to Le Chatelier s principle, we expect a decrease in solubility where a common ion will shift the equilibrium to the left (reactants side). Question 5. Calculate the molar solubility of lead iodide in (a) water and (b) in 0.200 M sodium iodide solution. The Ksp for PbI 2 is 7.9 x 10-9. 45
The Effect of ph If the cation of a substance is a weak acid or the anion is a weak base, the solubility of the substance will be affected by the ph of the solution. The calculation of the solubility of salts in which the anion and/or cation are involved in acid-base equilibria can be done in a routine manner if conditional solubility product constants are used. Effect of Complex-ion Formation The presence of complexing agents that can combine with either the cation or anion of a slightly soluble substance will lead to an increase in its solubility. Thus silver chloride is quite insoluble in water but dissolves in dilute, aqueous ammonia: AgCl(s) Ag + + Cl - Ag + + + NH 3 AgNH 3 AgNH + + 3 + NH 3 Ag(NH 3 ) 2 Solubilites are calculated using the conditional solubility product constants. Titrants and Standards Silver nitrate is the titrant used for determining halide and thiocyanate ions. It is available in primary-standard-grade purity, but it is expensive. Titrants prepared from less pure silver nitrate can be standardized with primary-standard potassium chloride. 46
Titration Curves Precipitation titrations can be divided into four basic regions based on composition: Initial conditions Before the equivalence point At the equivalence point After the equivalence point Example Consider the determination of Cl - by titration with AgNO 3. (The sample and titrant are 0.1000 M solutions and volume of sample = 50.00 ml) The titration reaction is: AgNO 3 + NaCl AgCl(s) + NaNO 3 Before the start of the titration At this point no titrant has been added. [Cl - ] = 0.1000 M pcl = -log [Cl - ] = 1.000 Before the equivalence point This would be any point after you have started to add some reagent but before reaching the equivalence point. As soon as we start adding Ag +, AgCl begins to form. In this region, virtually all Ag + will precipitate as AgCl. [Cl - ] can be determined by calculating the amount of AgCl that has precipitated. [Cl - ] = starting moles Cl - - moles AgCl precipitated total volume 47
After addition of 5 ml Ag + Solution volume = 55 ml mmol AgCl formed = 5 ml x 0.1000M Ag + = 0.5 mmoles AgCl [Cl - ] = 5 mmol - 0.5 mmol = 0.08181 M 55 ml Thereafter, determine pcl. pcl = -log (0.08181) = 1.0871 We can also calculate the [Ag + ] using our Ksp expression. [Ag + ] = Ksp / [Cl - ] = 1.8x10-10 / 0.08181 = 2.20 x 10-9 pag = -log (2.20 x 10-9 ) = 8.657 48
At the equivalence point The only source of chloride is the precipitate. Since silver ion and chloride ion are formed simultaneously when the precipitate dissolves and there is no other source of either ion, [Ag + ] = [Cl - ] Substituting in the Ksp expression (Ksp AgCl = [Ag + ] [Cl - ] ) gives Ksp = [Ag + ] 2 [Ag + ] = Ksp = 1.8 x 10-10 = 1.34 x 10-5 = [Cl - ] pcl = pag = 4.873 After the equivalence point In this region, the excess silver nitrate is the major supplier of silver ion. The amount of excess silver nitrate is found by taking the difference between the amount of silver nitrate added and the amount of silver nitrate consumed. [Ag + ] = moles Ag + added - moles Ag + used total volume and [Cl - ] = Ksp/ [Ag + ] After the addition of 55.00 ml of titrant : [Ag + ] = 0.0048 M pag = 2.32 pcl = 7.42 49
Chemical Indicators The indicators used in precipitation reactions are usually specific compound formers: that is they react selectively with the titrant to form a coloured substance. Since both the analyte A and the indicator In can react with the titrant T, they may be viewed as competitors: Titration reaction: Indicator reaction: A + T AT(s) In + T InT Applications Methods for Endpoint Detection Three approaches for endpoint detection of the Ag + and Cl - titration will be reviewed. Mohr method competitive anion Volhard method - complex formation Fajans Method adsorption indicator Mohr method for chloride This method was developed for the determination of chloride, bromide and cyanide ions. It uses silver nitrate as the titrant and sodium chromate as the indicator. Titration reaction: Ag + + Cl - AgCl(s) white Indicator reaction: 2 Ag + + CrO 4 2- Ag 2 CrO 4 (s) orange-red Precipitation is dependent on both the silver and the chromate concentration. 50
This approach relies on the Ksp differences for two insoluble silver salts. Ag + + Cl - AgCl(s) (titration reaction) K sp = 1.8 x 10-10 2Ag + + CrO 4 2- Ag 2 CrO 4(s) (at endpoint) K sp = 1.1 x 10-12 AgCl is less soluble than Ag 2 CrO 4 so it will precipitate first. We do not want any Ag 2 CrO 4 to precipitate until after we reached the AgCl equivalence point. At equivalence point : [Ag + ] = [Cl - ] And so [Ag + ] = Ksp AgCl = 1.34 x 10-5 M At higher concentrations Ag 2 CrO 4 can start to form. The maximum chromate concentration that we can have without forming a precipitate at [Ag + ] = 1.34 x 10-5 M is: [CrO4 2- ] = Ksp / [Ag + ] 2 = 6.1 x 10-3 M By trial and error 2.5 x 10-3 M has been found to work best. Titration cannot be conducted in acidic solutions. A ph of about 8 is best. Volhard Method This method involves the addition of an excess of silver ion to a halide solution, followed by back titration of the excess with thiocyanate ion. Ferric ion is used as the indicator, the endpoint being marked by the appearance of red FeSCN 2+. This is an indirect method for chloride determination based on competitive complex formation. The main advantage of this method is that it can be carried out in strongly acidic solutions. Analyte reaction: Ag + + Cl - AgCl(s) white Titration reaction: Ag + + SCN - AgSCN(s) white Indicator reaction: Fe 3+ + SCN - FeSCN 2+ 51
Question 6. A 500.0 mg sample of butter was warmed and shaken vigorously with water. The undissolved material was removed by filtering and the aqueous portion was made 1.0 M in HNO 3 and 0.025 M in Fe(NO 3 ) 3. This acidified solution was treated with 10.00 ml of 0.1755 M AgNO3 to precipitate the chloride ion and, after the addition of a small amount of nitrobenzene, 14.22 ml of 0.1006 M KSCN was required to back titrate the excess Ag +. Calculate the % NaCl in the butter. Most procedures call for the indicator to be added after the excess silver nitrate. In the determination of iodide, the particular order of reagent addition is crucial because iodide can reduce the ferric-ion indicator: 2Fe 3+ + 2I - 2Fe 2+ + I 2 The resulting ferrous ion does not form a colored complex with thiocyanate ion, and no endpoint can be detected. Fajans method This is an absorption indicator method where the endpoint reaction occurs on the surface of the AgCl precipitate. It relies on the change in the primary absorbed ion which occurs when we go past the equivalence point. This method uses an indicator that is absorbed by the electrostatically charged colloidal precipitate immediately after the equivalence point. Such absorption indicators have different colours in the free states and the adsorbed states. The indicators used most commonly in the titration of halides are dichlorofluorescein and tetrabromofluorescein, whose common name is eosin. 52
The dichlorofluorescein anion (DCF - ) has a greenish-yellow colour in solution that turns pink when it is adsorbed on silver chloride. The change in colour is thought to be due to a formation of the indicator anion resulting from the adsorption process. Ag + + Cl - AgCl Titration reaction Ag + + AgCl + DCF - AgCl : Ag + : DCF - Indicator reaction Greenish-yellow pink Adsorption indicators To illustrate how an adsoption indicator functions, we need to review the electrical behavior of colloidal silver halides. Colloidal silver chloride in the presence of excess chloride ion will adsorb some of these ions and acquire a negative charge. This negatively charged particle then attracts positive ions to a loosely bound counter-ion layer. In the presence of excess silver, the colloidal silver chloride will acquire a positive charge, which in turn will attract negative ions to the counter-ion layer.. During the titration of chloride ion with silver ion, an abrupt change in the charge on the silver chloride precipitate occurs in the vicinity of the equivalence point where excess chloride ion is replaced by excess silver ion. When the primary adsorption layer becomes positive, the indicator anion is attracted to the counter-ion layer and the colour change occurs. If the indicator is adsorbed more strongly than the analyte ion, it cannot be used in the titration because the colour change will occur at the beginning of the titration. Dichlorofluorescein is adsorbed less strongly than Cl -. Br -, I - or SCN - and can be used in the titration of any of these ions. Eosin is adsorbed more srongly than Cl - but less strongly than Br -, I - or SCN - and cannot be used in the titration of chloride ion. For the titration of Br -, I - or SCN -, where a choice of indicators is possible, eosin is preferred because it can be used at a lower ph., thereby avoiding a number of potential interferences. Dichlorofluorescein cannot be used below about ph 6.5 because it is a very weak acid and too little DCF - exists at a low ph. 53
For adsorption indicators to function properly, at least a portion of the silver halide precipitate must remain colloidal in the vicinity of the equivalence point. Unfortunately, a significant fraction of the precipitate tends to coagulate just before the equivalence point, especially when high concentrations are being titrated. If the titration is continued slowly with vigorous shaking, a good end point can be obtained. Some procedures determining chloride call for the addition of dextrin as a protective colloid to retard coagulation. Factors affecting the adsorption endpoint Intensity of colour is determined by the number of indicators molecules adsorbed. The indicator ion must not be able to displace the primary adsorbed ion. It must be adsorbed by the counter ion present at the endpoint. Dependent on indicator concentration and precipitate surface area. 54