Equivalence point precipitation titration

At the equivalence point, the titration curve is steepest for the least soluble precipitate... 

The solubility of metal-hydroxide precipitates in water varies depending on ionic strength and number of pairs and/or complexes (Chapter 2). A practical approach to determining the pH of minimum metal-hydroxide solubility, in simple or complex solutions, is potentiometric titration, as demonstrated in Figure 12.3. The data show that potentiometric titration of a solution with a given heavy metal is represented by a sigmoidal plot. The long pH plateau represents pH values at which metals precipitate the equivalence point, or titration end point, indicates the pH at the lowest metal-... 

Advantages and Limitations of Radiometric Titrations. Radiometric detection of the equivalence point is a general method that does not depend on the chemical reaction employed. This contrasts with other methods of detection, which depend on specific chemical or physical transitions at the equivalence point. Amperometric titrations are applicable only to electrochemically active systems conductometric titrations apply only to ionic solutions, and so on. In principle, any titration system in which a phase separation can be effected is amenable to radiometric detection, provided there exist suitable radioactive labels. The major limitation of the method is the requirement for phase separation. In precipitation titrations, the phase separation is automatic and the method is well suited to this class of titrations. For other classes of titrations, special phase-separation methods, such as solvent extraction, need to be applied. At the present time, the method suffers from a lack of phase-separation techniques suitable for continuous monitoring of the titration curves. 

Fig. 21. Hypothetical titration curve for solutions of free bile salts or for glycine conjugates. IV — first equivalence point where titration of bile salt with hydrochloric acid commences, Y = last point where bile salt solution is in thermodynamic equilibrium as a single aqueous phase. T = Tyndall effect noted in this region of titration curve. X = point where precipitation of bile acid crystals commences. X — equilibrium pH at point of bile acid precipitation. Z = second equivalence point where titration of bile salt with hydrochloric acid is complete. TOT = total amount of acid required to complete the titration. HA = the amount of acid added from the first equivalence point (IV), to point V, which represents the maximum solubility of the bile acid (HA) in the bile salt solution (A"). For further explanation of the symbols, see text.

For a titration to be accurate we must add a stoichiometrically equivalent amount of titrant to a solution containing the analyte. We call this stoichiometric mixture the equivalence point. Unlike precipitation gravimetry, where the precipitant is added in excess, determining the exact volume of titrant needed to reach the equivalence point is essential. The product of the equivalence point volume, Veq> and the titrant s concentration, Cq, gives the moles of titrant reacting with the analyte. 

Potcntiomctric Titrations In Chapter 9 we noted that one method for determining the equivalence point of an acid-base titration is to follow the change in pH with a pH electrode. The potentiometric determination of equivalence points is feasible for acid-base, complexation, redox, and precipitation titrations, as well as for titrations in aqueous and nonaqueous solvents. Acid-base, complexation, and precipitation potentiometric titrations are usually monitored with an ion-selective electrode that is selective for the analyte, although an electrode that is selective for the titrant or a reaction product also can be used. A redox electrode, such as a Pt wire, and a reference electrode are used for potentiometric redox titrations. More details about potentiometric titrations are found in Chapter 9. 

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. 

A. Direct titration. The solution containing the metal ion to be determined is buffered to the desired pH (e.g. to PH = 10 with NH4-aq. NH3) and titrated directly with the standard EDTA solution. It may be necessary to prevent precipitation of the hydroxide of the metal (or a basic salt) by the addition of some auxiliary complexing agent, such as tartrate or citrate or triethanolamine. At the equivalence point the magnitude of the concentration of the metal ion being determined decreases abruptly. This is generally determined by the change in colour of a metal indicator or by amperometric, spectrophotometric, or potentiometric methods. 

Similar remarks apply to the determination of bromides the Mohr titration can be used, and the most suitable adsorption indicator is eosin which can be used in dilute solutions and even in the presence of 0.1 M nitric acid, but in general, acetic (ethanoic) acid solutions are preferred. Fluorescein may be used but is subject to the same limitations as experienced with chlorides [Section 10.77(b)], With eosin indicator, the silver bromide flocculates approximately 1 per cent before the equivalence point and the local development of a red colour becomes more and more pronounced with the addition of silver nitrate solution at the end point the precipitate assumes a magenta colour. 

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. 

Titrations can be carried out in cases in which the solubility relations are such that potentiometric or visual indicator methods are unsatisfactory for example, when the reaction product is markedly soluble (precipitation titration) or appreciably hydrolysed (acid-base titration). This is because the readings near the equivalence point have no special significance in amperometric titrations. Readings are recorded in regions where there is excess of titrant, or of reagent, at which points the solubility or hydrolysis is suppressed by the Mass Action effect the point of intersection of these lines gives the equivalence point. 

In principle, any type of titration can be carried out conductometrically provided that during the titration a substantial change in conductance takes place before and/or after the equivalence point. This condition can be easily fulfilled in acid-base, precipitation and complex-formation titrations and also the corresponding displacement titrations, e.g., a salt of a weak acid reacting with a strong acid or a metal in a fairly stable complex reacting with an anion to yield a very stable complex. However, for redox titrations such a condition is rarely met. 

In fact, any type of titration can be carried out potentiometrically provided that an indicator electrode is applied whose potential changes markedly at the equivalence point. As the potential is a selective property of both reactants (titrand and titrant), notwithstanding an appreciable influence by the titration medium [aqueous or non-aqueous, with or without an ISA (ionic strength adjuster) or pH buffer, etc.] on that property, potentiometric titration is far more important than conductometric titration. Moreover, the potentiometric method has greater applicability because it is used not only for acid-base, precipitation, complex-formation and displacement titrations, but also for redox titrations. 

EXPERIMENTAL PROCEDURES AUTOMATED TURBIDIMETRIC TITRATION. A method for the automated aqueous turbidimetric titration of surfactants has been published (10) in which anionic surfactants are titrated against N-cetylpyridinium chloride to form a colloidal precipitate near the equivalence point. N-cetylpyridinium halides have a disadvantage in that they have the tendency to crystallise out of solution (15), consequently the strength of the solution may alter slightly without the knowledge of the operator, also the crystals suspended in solution may cause damage to the autotitrator. In view of these drawbacks hyamine was preferred as the titrant. 

The end points of precipitation titrations can be variously detected. An indicator exhibiting a pronounced colour change with the first excess of the titrant may be used. The Mohr method, involving the formation of red silver chromate with the appearance of an excess of silver ions, is an important example of this procedure, whilst the Volhard method, which uses the ferric thiocyanate colour as an indication of the presence of excess thiocyanate ions, is another. A series of indicators known as adsorption indicators have also been utilized. These consist of organic dyes such as fluorescein which are used in silver nitrate titrations. When the equivalence point is passed the excess silver ions are adsorbed on the precipitate to give a positively charged surface which attracts and adsorbs fluoresceinate ions. This adsorption is accompanied by the appearance of a red colour on the precipitate surface. Finally, the electroanalytical methods described in Chapter 6 may be used to scan the solution for metal ions. Table 5.12 includes some examples of substances determined by silver titrations and Table 5.13 some miscellaneous precipitation methods. Other examples have already been mentioned under complexometric titrations. 

The electrical conductance of a solution is a measure of its current-carrying capacity and is therefore determined by the total ionic strength. It is a nonspecific property and for this reason direct conductance measurements are of little use unless the solution contains only the electrolyte to be determined or the concentrations of other ionic species in the solution are known. Conductometric titrations, in which the species of interest are converted to non-ionic forms by neutralization, precipitation, etc. are of more value. The equivalence point may be located graphically by plotting the change in conductance as a function of the volume of titrant added. 

The potentiometric titrations of [Cu1(MeCN)4](CIO4), AgIC104, and AuIC104 with (Bu4N)0H(in MeOH) are illustrated in Figure 4, and demonstrate that each process has one-to-one stoichiometry. The three systems form precipitates such that all of the metal is removed from solution at the equivalence point. Addition of excess "OH causes some dissolution of the CuOH and AgOH precipitates, and appears as a second step for the titration curve of Ag(I) (Figure 4b). 

Example The titration of Pb2+ with S042 or C2042 ions. An appreciably high potential is usually applied to yield a diffusion current for lead. From Fig. 1(A), one may evidently observe a linear decrease in current because Pb2+ ions are removed from the solution by precipitation. The small curvature just prior to the end-point (or equivalence point) shows the incompleteness of the analytical reaction in this particular region. However, the end-point may be achieved by extrapolation of the linear portions, as shown in the said figure. 

In the next chapter, you will extend your knowledge of equilibria involving aqueous ions. You will learn how to calculate the pH at an equivalence point, so you can select an appropriate indicator for any acid-hase titration. You will also learn why equilihrium is important to the solubility of compounds that are slightly soluble, and how to predict whether a precipitate will form as the result of a reaction between ions in solution. 

Consider the addition of AgN03 to a solution containing KI and KC1. Because Kip(AgI) A spCAgCl), Agf precipitates first. When precipitation of I is almost complete, the concentration of Ag+ abruptly increases and AgCl begins to precipitate. When Cl" is consumed, another abrupt increase in [Ag+] occurs. We expect to see two breaks in the titration curve. The first corresponds to the Agl equivalence point, and the second to the AgCl equivalence point. 

Common adsorption indicators are anionic dyes, which are attracted to the positively charged particles produced immediately after the equivalence point. Adsorption of the negatively charged dye onto the positively charged surface changes the color of the dye. The color change is the end point in the titration. Because the indicator reacts with the precipitate... 

In a spectrophotometric titration, absorbance of light is monitored as titrant is added. For many reactions, there is an abrupt change in absorbance when the equivalence point is reached. The Fajans titration is based on the adsorption of a charged indicator onto the charged surface of the precipitate after the equivalence point. The Volhard titration, used to measure Ag+, is based on the reaction of Fe3+ with SCN- after the precipitation of AgSCN is complete. 

Concentrations of reactants and products during a precipitation titration are calculated in three regions. Before the equivalence point, there is excess analyte. Its concentration is the product (fraction remaining) X (original concentration) X (dilution factor). The concentration of titrant can be found from the solubility product of the precipitate and the known concentration of excess analyte. At the equivalence point, concentrations of both reactants are governed by the solubility product. After the equivalence point, the concentration of analyte can be determined from the solubility product of precipitate and the known concentration of excess titrant. 

The text claims that precipitation of I is not complete before Cl- begins to precipitate in the titration in Figure 7-8. Calculate the concentration of Ag+ at the equivalence point in the titration of 1 alone. Show that this concentration of Ag+ will precipitate Cl... 

Examine the procedure in Table 7-1 for the Fajans titration of Zn2+. Do you expect the charge on the precipitate to be positive or negative after the equivalence point ... 

The figure shows the results of a bipotentiometric titration of ascorbic acid with If. Ascorbic acid (146 mg) was dissolved in 200 mL of water in a 400-mL beaker. Two Pt electrodes were attached to the K-F outlets of the pH meter and spaced about 4 cm apart in the magnetically stirred solution. The solution was titrated with 0.04 M If (prepared by dissolving 2.4 g of K1 plus 1.2 g of I2 in 100 mL of water), and the voltage was recorded after each addition. Prior to the equivalence point, all the If is reduced to I- by the excess ascorbic acid. Reaction B can occur, but Reaction A cannot. A voltage of about 300 mV is required to support a constant current of 10 pA. (The ascorbate dehydroascorbate couple does not react at a Pt electrode and cannot carry current.) After the equivalence point, excess If is present, so Reactions A and B both occur, and the voltage drops precipitously. 

Add 0.1 M tartaric acid in the following portions 4x5 ml, 10 x 1 ml, 4x5 ml. Allow the potential of the GC/PANI electrode to stabilize between each addition. Recording of the titration curve in this manner takes ca. 2 h. An example of the titration curve is shown in Fig. 5.2. The occurrence of a potential maximum at 5-10 ml and the minimum at 15 ml tartaric acid indicates where precipitation of trimeprazine tartrate starts. The increasing potential around 25 ml tartaric acid indicates the equivalence point. Titration of 1.4 M trimeprazine base with 0.7M tartaric acid results in a larger potential change at the equivalence point, as can be seen in Fig. 5.3. 

Another class of indicators, known as adsorption indicators, adsorb to (or desorb from) the precipitate or colloidal particles of the silver salt of the analyte at the equivalence point. The indicator anions are attracted into the counterion layer surrounding each colloidal particle of silver salt. Thus, there is a transfer of color from the solution to the solid or from the solid to the solution at the end point. The concentration of the indicator, which is an organic compound, is not large enough to cause its precipitation as a silver salt. Thus, the color change is an adsorption and not a precipitation process. Fluorescein is a typical example of an adsorption indicator in argentometric titration. 

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