Coprecipitation solids

The stoichiometry must be exact. Coprecipitation by solid-solution formation, foreign ion entrapment, and adsorption are possible sources of error. 

Occlusions, which are a second type of coprecipitated impurity, occur when physically adsorbed interfering ions become trapped within the growing precipitate. Occlusions form in two ways. The most common mechanism occurs when physically adsorbed ions are surrounded by additional precipitate before they can be desorbed or displaced (see Figure 8.4a). In this case the precipitate s mass is always greater than expected. Occlusions also form when rapid precipitation traps a pocket of solution within the growing precipitate (Figure 8.4b). Since the trapped solution contains dissolved solids, the precipitate s mass normally increases. The mass of the precipitate may be less than expected, however, if the occluded material consists primarily of the analyte in a lower-molecular-weight form from that of the precipitate. 

Inclusions, occlusions, and surface adsorbates are called coprecipitates because they 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 necessary to minimize the solubility of a desired precipitate may lead to the formation of an additional precipitate that interferes in the analysis. For example, the precipitation of nickel dimethylgloxime requires a plT that is slightly basic. Under these conditions, however, any Fe + that might be present precipitates as Fe(01T)3. Finally, since most precipitants are not selective toward a single analyte, there is always a risk that the precipitant will react, sequentially, with more than one species. 

An additional 3-5 g. (6-10%) of crude product can be obtained by concentrating the filtrate to about 800 ml., cooling the concentrate to 0°, and filtering the solid that separates. This solid must be washed well with cold water to remove coprecipitated potassium chloride. 

Zinc, Cu and Ni have similar ionic radii and electron configurations (Table 5.6). Due to the similarity of the ionic radii of these three metals with Fe and Mg, Zn, Cu and Ni are capable of isomorphous substitution of Fe2+ and Mg2+ in the layer silicates. Due to differences in the electronegativity, however, isomorphous substitution of Cu2+ in silicates may be limited by the greater Pauling electronegativity of Cu2+ (2.0), whereas Zn2+ (1.6) and Ni2+ (1.8) are relatively more readily substituted for Fe2+ (1.8) or Mg2+ (1.3) (McBride, 1981). The three metals also readily coprecipitate with and form solid solutions in iron oxides (Lindsay, 1979 Table 5.7). 

Whereas studies have been carried out on the factors (surface coverage, residence time, pH) which influence the desorption of arsenate previously sorbed onto oxides, phyllosilicates and soils (O Reilly et al. 2001 Liu et al. 2001 Arai and Sparks 2002 Violante and Pigna 2002 Pigna et al. 2006), scant information are available on the possible desorption of arsenate coprecipitated with iron or aluminum. In natural environments arsenic may form precipitates or coprecipitates with Al, Fe, Mn and Ca. Coprecipitation of arsenic with iron and aluminum are practical and effective treatment processes for removing arsenic from drinking waters and might be as important as sorption to preformed solids. 

A variety of preconcentration procedures has been used, including solvent extraction of metal chelates, coprecipitation, chelating ion exchange, adsorption onto other solids such as silica-bonded organic complexing agents, and liquid-liquid extraction. 

Another method of segregating and removing a portion of the dissolved organic matter includes incorporation into a solid phase, which can then be removed by filtration or centrifugation. This incorporation can result either from coprecipitation with a solid phase formed in a reaction in the solution or, as discussed in the next section, from adsorption of the organic material onto a pre-existing solid phase. 

YDC has been prepared by various methods, including solid-state reaction, coprecipitation, glycine-nitrate process and metal organic chemical vapor deposition (MOCVD). Table 1.4 shows that the properties depend on the preparation method [118,119,127,129,130], Zha et al. [129] have studied the influence of sintering... 

FIGURE 1.37 The Arrhenius plot of Ce08Sm0i2Oli9 from different methods (A) solid-state reaction [95] (B) sol-gel process [116] (C) oxalate coprecipitation [91] (D) carbonate coprecipitation, our work and (E) glycine-nitrate process [157]. 

Usually, however, the distribution coefficients determined experimentally are not equal to the ratios of the solubility product because the ratio of the activity coefficients of the constituents in the solid phase cannot be assumed to be equal to 1. Actually observed D values show that activity coefficients in the solid phase may differ markedly from 1. Let us consider, for example, the coprecipitation of MnC03 in calcite. Assuming that the ratio of the activity coefficients in the aqueous solution is close to unity, the equilibrium distribution may be formulated as (cf. Eq. A.6.11)... 

We already touched on some aspects of carbonate surface chemistry e.g., in Chapter 3.4. We have already illustrated some of the factors that affect surface charge and the point of zero charge, pHpzc, in Chapter 3.5, and have discussed certain elementary aspects of CaC03 nucleation in Chapter 6.5 and of coprecipitation (and solid solution formation) in Chapter 6.7. 

Carbonate minerals in natural systems precipitate in the presence of various other solutes This trace amounts of all components present in the solution may get incorporated into the solid carbonate minerals ("coprecipitation"). 

The phenomena of surface precipitation and isomorphic substitutions described above and in Chapters 3.5, 6.5 and 6.6 are hampered because equilibrium is seldom established. The initial surface reaction, e.g., the surface complex formation on the surface of an oxide or carbonate fulfills many criteria of a reversible equilibrium. If we form on the outer layer of the solid phase a coprecipitate (isomorphic substitutions) we may still ideally have a metastable equilibrium. The extent of incipient adsorption, e.g., of HPOjj on FeOOH(s) or of Cd2+ on caicite is certainly dependent on the surface charge of the sorbing solid, and thus on pH of the solution etc. even the kinetics of the reaction will be influenced by the surface charge but the final solid solution, if it were in equilibrium, would not depend on the surface charge and the solution variables which influence the adsorption process i.e., the extent of isomorphic substitution for the ideal solid solution is given by the equilibrium that describes the formation of the solid solution (and not by the rates by which these compositions are formed). Many surface phenomena that are encountered in laboratory studies and in field observations are characterized by partial, or metastable equilibrium or by non-equilibrium relations. Reversibility of the apparent equilibrium or congruence in dissolution or precipitation can often not be assumed. 

Vaslow, F., and G. E. Boyd (1952), "Thermodynamics of Coprecipitation Dilute Solid Solutions of AgBr in AgCI1", J. Amer. Chem. Soc. 74, 4691. 

Sol id Sol utions. The aqueous concentrations of trace elements in natural waters are frequently much lower than would be expected on the basis of equilibrium solubility calculations or of supply to the water from various sources. It is often assumed that adsorption of the element on mineral surfaces is the cause for the depleted aqueous concentration of the trace element (97). However, Sposito (Chapter 11) shows that the methods commonly used to distinguish between solubility or adsorption controls are conceptually flawed. One of the important problems illustrated in Chapter 11 is the evaluation of the state of saturation of natural waters with respect to solid phases. Generally, the conclusion that a trace element is undersaturated is based on a comparison of ion activity products with known pure solid phases that contain the trace element. If a solid phase is pure, then its activity is equal to one by thermodynamic convention. However, when a trace cation is coprecipitated with another cation, the activity of the solid phase end member containing the trace cation in the coprecipitate wil 1 be less than one. If the aqueous phase is at equil ibrium with the coprecipitate, then the ion activity product wi 1 1 be 1 ess than the sol ubi 1 ity constant of the pure sol id phase containing the trace element. This condition could then lead to the conclusion that a natural water was undersaturated with respect to the pure solid phase and that the aqueous concentration of the trace cation was controlled by adsorption on mineral surfaces. While this might be true, Sposito points out that the ion activity product comparison with the solubility product does not provide any conclusive evidence as to whether an adsorption or coprecipitation process controls the aqueous concentration. 

Figure 3. Percent coprecipitation of AX vs. that of BX from an aqueous solution under the assumption that AX and BX form ideal solid solutions. Modified from Gordon et al. (10).

Equilibrium between simple salts and aqueous solutions is often relatively easily demonstrated in the laboratory when the composition of the solid is invariant, such as occurs in the KCI-H2O system. However, when an additional component which coprecipitates is added to the system, the solid composition is no longer invariant. Very long times may be required to reach equilibrium when the reaction path requires shifts in the composition of both the solution and solid. Equilibrium is not established until the solid composition is homogeneous and the chemical potentials of all components between solid and aqueous phases are equivalent. As a result, equilibrium is rarely demonstrated with a solid solution series. 

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