D. Harvey - Modern Analytical Chemistry (794078), страница 71
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Inparticular, a large excess of chloride must be avoided.Another important parameter that may affect a precipitate’s solubility is the pHof the solution in which the precipitate forms. For example, hydroxide precipitates,such as Fe(OH)3, are more soluble at lower pH levels at which the concentration ofOH– is small. The effect of pH on solubility is not limited to hydroxide precipitates,but also affects precipitates containing basic or acidic ions.
The solubility ofCa3(PO4)2 is pH-dependent because phosphate is a weak base. The following fourreactions, therefore, govern the solubility of Ca3(PO4)2.Kspt 3Ca2+(aq) + 2PO43–(aq)KPO43–(aq) + H2O(l) t HPO42–(aq) + OH–(aq)KHPO42–(aq) + H2O(l) t H2PO4–(aq) + OH–(aq)KH2PO4–(aq) + H2O(l) t H3PO4(aq) + OH–(aq)Ca3(PO4)2(s)b1b2b38.88.98.108.11Depending on the solution’s pH, the predominate phosphate species is either PO43–,HPO42–, H2PO4–, or H3PO4. The ladder diagram for phosphate, shown in Figure8.2a, provides a convenient way to evaluate the pH-dependent solubility of phosphate precipitates. When the pH is greater than 12.4, the predominate phosphatespecies is PO43–, and the solubility of Ca3(PO4)2 will be at its minimum becauseonly reaction 8.8 occurs to an appreciable extent (see Figure 8.2b).
As the solutionbecomes more acidic, the solubility of Ca3(PO4)2 increases due to the contributionsof reactions 8.9–8.11.Solubility can often be decreased by using a nonaqueous solvent. A precipitate’s solubility is generally greater in aqueous solutions because of the abilityof water molecules to stabilize ions through solvation. The poorer solvatingability of nonaqueous solvents, even those that are polar, leads to a smaller solubility product.
For example, PbSO4 has a Ksp of 1.6 × 10–8 in H2O, whereas in a50:50 mixture of H2O/ethanol the Ksp at 2.6 × 10–12 is four orders of magnitudesmaller.14(a) Ladder diagram for phosphate;(b) Solubility diagram for Ca3(PO4)2showing the predominate form ofphosphate for each segment ofthe solubility curve.1400-CH08 9/9/99 2:17 PM Page 238238Modern Analytical ChemistryCl–Ag+Figure 8.3Schematic model of AgCl showingdifference between bulk and surface atomsof silver.
Silver and chloride ions are notshown to scale.inclusionA coprecipitated impurity in which theinterfering ion occupies a lattice site inthe precipitate.Bulk silver ionsurrounded by 6 chloride ionsSurface silver ionsurrounded by 4 chloride ionsAvoiding Impurities Precipitation gravimetry is based on a known stoichiometrybetween the analyte’s mass and the mass of a precipitate.
It follows, therefore, thatthe precipitate must be free from impurities. Since precipitation typically occurs ina solution rich in dissolved solids, the initial precipitate is often impure. Any impurities present in the precipitate’s matrix must be removed before obtaining itsweight.The greatest source of impurities results from chemical and physical interactions occurring at the precipitate’s surface. A precipitate is generally crystalline,even if only on a microscopic scale, with a well-defined lattice structure of cationsand anions. Those cations and anions at the surface of the precipitate carry, respectively, a positive or a negative charge as a result of their incomplete coordinationspheres. In a precipitate of AgCl, for example, each Ag+ ion in the bulk of the precipitate is bound to six Cl– ions.
Silver ions at the surface, however, are bound to nomore than five Cl– ions, and carry a partial positive charge (Figure 8.3).Precipitate particles grow in size because of the electrostatic attraction betweencharged ions on the surface of the precipitate and oppositely charged ions in solution. Ions common to the precipitate are chemically adsorbed, extending the crystallattice. Other ions may be physically adsorbed and, unless displaced, are incorporated into the crystal lattice as a coprecipitated impurity. Physically adsorbed ionsare less strongly attracted to the surface and can be displaced by chemically adsorbed ions.One common type of impurity is an inclusion.
Potential interfering ions whosesize and charge are similar to a lattice ion may substitute into the lattice structure bychemical adsorption, provided that the interferent precipitates with the same crystalstructure (Figure 8.4a). The probability of forming an inclusion is greatest when theinterfering ion is present at substantially higher concentrations than the dissolvedlattice ion. The presence of an inclusion does not decrease the amount of analytethat precipitates, provided that the precipitant is added in sufficient excess.
Thus,the precipitate’s mass is always larger than expected.Inclusions are difficult to remove since the included material is chemically partof the crystal lattice. The only way to remove included material is through reprecipitation.
After isolating the precipitate from the supernatant solution, it is dissolved1400-CH08 9/9/99 2:17 PM Page 239239Chapter 8 Gravimetric Methods of Analysisin a small portion of a suitable solvent at an elevated temperature. The solution isthen cooled to re-form the precipitate.
Since the concentration ratio of interferentto analyte is lower in the new solution than in the original supernatant solution, themass percent of included material in the precipitate decreases. This process of reprecipitation is repeated as needed to completely remove the inclusion. Potentialsolubility losses of the analyte, however, cannot be ignored. Thus, reprecipitationrequires a precipitate of low solubility, and a solvent for which there is a significantdifference in the precipitate’s solubility as a function of temperature.Occlusions, which are a second type of coprecipitated impurity, occur whenphysically adsorbed interfering ions become trapped within the growing precipitate.Occlusions form in two ways.
The most common mechanism occurs when physicallyadsorbed ions are surrounded by additional precipitate before they can be desorbedor displaced (see Figure 8.4a). In this case the precipitate’s mass is always greater thanexpected. Occlusions also form when rapid precipitation traps a pocket of solutionwithin the growing precipitate (Figure 8.4b). Since the trapped solution contains dissolved solids, the precipitate’s mass normally increases. The mass of the precipitatemay be less than expected, however, if the occluded material consists primarily of theanalyte in a lower-molecular-weight form from that of the precipitate.Occlusions are minimized by maintaining the precipitate in equilibrium withits supernatant solution for an extended time.
This process is called digestion andmay be carried out at room temperature or at an elevated temperature. During digestion, the dynamic nature of the solubility–precipitation equilibrium, in whichthe precipitate dissolves and re-forms, ensures that occluded material is eventuallyexposed to the supernatant solution. Since the rate of dissolution and reprecipitation are slow, the chance of forming new occlusions is minimal.After precipitation is complete the surface continues to attract ions from solution (Figure 8.4c). These surface adsorbates, which may be chemically or physicallyadsorbed, constitute a third type of coprecipitated impurity.
Surface adsorption isminimized by decreasing the precipitate’s available surface area. One benefit of digestion is that it also increases the average size of precipitate particles. This is notsurprising since the probability that a particle will dissolve is inversely proportionalto its size. During digestion larger particles of precipitate increase in size at the expense of smaller particles. One consequence of forming fewer particles of larger sizeis an overall decrease in the precipitate’s surface area. Surface adsorbates also maybe removed by washing the precipitate. Potential solubility losses, however, cannotbe ignored.Inclusions, occlusions, and surface adsorbates are called coprecipitates becausethey represent soluble species that are brought into solid form along with the desired precipitate.
Another source of impurities occurs when other species in solution precipitate under the conditions of the analysis. Solution conditions necessaryto minimize the solubility of a desired precipitate may lead to the formation of anadditional precipitate that interferes in the analysis. For example, the precipitationof nickel dimethylgloxime requires a pH that is slightly basic. Under these conditions, however, any Fe3+ that might be present precipitates as Fe(OH)3. Finally,since most precipitants are not selective toward a single analyte, there is always arisk that the precipitant will react, sequentially, with more than one species.The formation of these additional precipitates can usually be minimized bycarefully controlling solution conditions.
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