gravimetric method chemical analysis, Lecture notes of Chemistry

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GRAVIMETRIC

METHOD

Precipitation

Reggie M. Gutierrez

PRECIPITATION GRAVIMETRY  (^) In precipitation gravimetry an insoluble compound forms when we add a precipitating reagent, or precipitant , to a solution containing our analyte. In most methods the precipitate is the product of a simple metathesis reaction between the analyte and the precipitant; however, any reaction generating a precipitate can potentially serve as a gravimetric method.  (^) Most precipitation gravimetric methods were developed in the nineteenth century, or earlier, often for the analysis of ores  (^) for example, a precipitation gravimetric method for the analysis of nickel in ores.

SOLUBILITY CONSIDERATIONS  (^) To provide accurate results, a precipitate’s solubility must be minimal. The accuracy of a total analysis technique typically is better than ±0.1%, which means that the precipitate must account for at least 99.9% of the analyte. Extending this requirement to 99.99% ensures that the precipitate’s solubility does not limit the accuracy of a gravimetric analysis.  (^) We can minimize solubility losses by carefully controlling the conditions under which the precipitate forms. This, in turn, requires that we account for every equilibrium reaction affecting the precipitate’s solubility. For example, we can determine Ag+^ gravimetrically by adding NaCl as a precipitant, forming a precipitate of AgCl. Ag+(aq)+Cl−(aq)⇌AgCl(s) (8.1) A total analysis technique is one in which the analytical signal—mass in this case—is proportional to the absolute amount of analyte in the sample. See classifying techniques

 (^) If this is the only reaction we consider, then we predict that the precipitate’s solubility, S AgCl, is given by the following equation. SAgCl=[Ag+]=KspCl−  (^) Equation 8.2 suggests that we can minimize solubility losses by adding a large excess of Cl

• . In fact, as shown in Figure 8.1, adding a large excess of Cl

increases the precipitate’s solubility.

 (^) To understand why the solubility of AgCl is more complicated than the relationship suggested by equation 8.2, we must recognize that Ag+also forms a series of soluble silver-chloro metal–ligand complexes. Ag+(aq)+Cl−(aq)⇌AgCl(aq)logK1=3.70 (8.3) AgCl(aq)+Cl−(aq)⇌AgCl−2(aq)logK2=1.92 (8.4) AgCl−2(aq)+Cl−(aq)⇌AgCl2−3(aq)logK3=0.78 (8.5)  (^) The actual solubility of AgCl is the sum of the equilibrium concentrations for all soluble forms of Ag+, as shown by the following equation. SAgCl=[Ag+]+[AgCl(aq)]+[AgCl−2]+[AgCl2−3] (8.6)  (^) By substituting into equation 8.6 the equilibrium constant expressions for reaction 8.1 and reactions 8.3–8.5, we can define the solubility of AgCl as

S AgCl=Ksp[Cl−]+K1Ksp+K1K2Ksp[Cl−]+K1K2K3Ksp[Cl−]2 (8.7)  (^) Equation 8.7 explains the solubility curve for AgCl shown in Figure 8.1. As we add NaCl to a solution of Ag+, the solubility of AgCl initially decreases because of reaction 8.1. Under these conditions, the final three terms in equation 8.7 are small and equation 8.1 is sufficient to describe AgCl’s solubility. At higher concentrations of Cl–, reaction 8.4 and reaction 8.5 increase the solubility of AgCl. Clearly the equilibrium concentration of chloride is important if we want to determine the concentration of silver by precipitating AgCl. In particular, we must avoid a large excess of chloride.  (^) The predominate silver-chloro complexes for different values of pCl are shown by the ladder diagram along the x -axis in Figure 8.1 Note that the increase in solubility begins when the higher-order soluble complexes, AgCl 2 –^ and AgCl 3 2–, become the predominate species.

 (^) Substituting the equilibrium constant expressions for reaction 8.8 and reaction 8.9 into equation 8.10 defines the solubility of CaF 2 in terms of the equilibrium concentration of H 3 O

. SCaF2=[Ca2+]=⎧⎩⎨Ksp4(1+[H3O+]Ka)2⎫⎭⎬1/3 (8.11)

 (^) Figure 8.2 shows how pH affects the solubility of CaF

Depending on the solution’s pH, the predominate form of fluoride is either HF or F

• . When the pH is greater than 4.17, the predominate species is F

and the solubility of CaF 2 is independent of pH because only reaction 8.8 occurs to an appreciable extent. At more acidic pH levels, the solubility of CaF 2 increases because of the contribution of reaction 8.9.

 (^) When solubility is a concern, it may be possible to decrease solubility by using a non-aqueous solvent. A precipitate’s solubility is generally greater in an aqueous solution because of water’s ability to stabilize ions through solvation. The poorer solvating ability of non- aqueous solvents, even those which are polar, leads to a smaller solubility product. For example, the K sp of PbSO 4 is 2 × 10

• in H 2 O and 2.6 × 10 - in a 50:50 mixture of H 2 O and ethanol.

PRACTICE EXERCISE  (^) You can use a ladder diagram to predict the conditions for minimizing a precipitate’s solubility. Draw a ladder diagram for oxalic acid, H 2 C 2 O 4 , and use it to establish a suitable range of pH values for minimizing the solubility of CaC 2 O 4.

AVOIDING IMPURITIES  (^) In addition to having a low solubility, the precipitate must be free from impurities. Because precipitation usually occurs in a solution that is rich in dissolved solids, the initial precipitate is often impure. We must remove these impurities before determining the precipitate’s mass.  (^) The greatest source of impurities is the result of 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 of cations and anions. Those cations and anions at the precipitate’s surface carry, respectively, a positive or a negative charge because they have incomplete coordination spheres. In a precipitate of AgCl, for example, each silver ion in the precipitate’s interior is bound to six chloride ions. A silver ion at the surface, however, is bound to no more than five chloride ions and carries a partial positive charge (Figure 8.3). The presence of these partial charges makes the precipitate’s surface an active site for the chemical and physical interactions that produce impurities.

 (^) One common impurity is an inclusion. A potential interfering ion whose size and charge is similar to a lattice ion, may substitute into the lattice structure, provided that the interferent precipitates with the same crystal structure (Figure 8.4a). The probability of forming an inclusion is greatest when the concentration of the interfering ion is substantially greater than the lattice ion’s concentration. An inclusion does not decrease the amount of analyte that precipitates, provided that the precipitant is present in sufficient excess. Thus, the precipitate’s mass is always larger than expected.

Figure 8.4 Three examples of impurities that may form during precipitation. The cubic frame represents the precipitate and the blue marks are impurities: (a) inclusions, (b) occlusions, and (c) surface adsorbates. Inclusions are randomly distributed throughout the precipitate. Occlusions are localized within the interior of the precipitate and surface adsorbates are localized on the precipitate’s exterior. For ease of viewing, in (c) adsorption is shown on only one surface.