Silver chloride precipitate, illustration

Suppose we have a solution that contains lead(II), mercury(I), silver, copper(II), and zinc ions. The method is outlined in Fig. 11.20, which includes additional cations, and is illustrated in Fig. 11.21. Most chlorides are soluble so, when hydrochloric acid is added to a mixture of salts, only certain chlorides precipitate (see Table 11.4). Silver and mercury(I) chlorides have such small values of Ksp that, even with low concentrations of Cl ions, the chlorides precipitate. Lead(II) chloride, which is slightly soluble, will precipitate if the chloride ion concentration is... 

The methods developed in the previous section for deriving titration curves can be extended to mixtures that form precipitates of different solubilities. To illustrate, consider the titration of 50.00 mL of a solution that is 0.0500 M in iodide ion and 0.0800 M in chloride ion with 0.1000 M silver nitrate. The curve for the initial stages of this titration is identical to the curve shown for iodide in Figure 13-5 because silver chloride, with its much larger solubility product, does not begin to precipitate until well into the titration. 

The equilibrium formation and dissolution of precipitates can be illustrated with silver chloride. Ag+forms a series of chloro complexes, at least up to AgCI4 v. In the presence of solid AgCl the formal description is most readily characterized by the following formalism ... 

The availability of equipment to measure molar conductivity of solutions was turned to good use. It is interesting to note that coordination chemists still make use of physical methods heavily in their quest to assign structures - it s just that the extent and sophistication of instrumentation has grown enormously in a century. What conductivity could tell the early coordination chemist was some further information about the apparently ionic species inferred to exist through the silver ion precipitation reactions. This is best illustrated for a series of platinum(IV) complexes with various amounts of chloride ion and ammonia present (Figure 3.1). From comparison of measured molar conductivity with conductivities of known compounds, the number of ions present in each of the complexes could be determined. We now understand these results in terms of modern formulation of the complexes as octahedral platinum(IV) compounds with coordinated ammonia, where coordinated chloride ions make up any shortfall in the fixed coordination number of six. This leaves in most cases some free ionic chloride ions to balance the charge on the complex cation. 

The solubility product principle enables simple calculations to be made of the effect of other species on the solubility of a given substance and may be used to determine the species that will precipitate in an electrolyte mixture. One simple result of applying the solubility product principle is the common ion effect. This is the effect caused by the addition of an ionic species that has an ion in common with the species of interest. Since the solubility of a species is given by the product of the concentration of its ions, when the concentration of one type of ion increases, the concentration of the other must decline, or the overall concentration of that compound must decline. We can illustrate this simply by using our previous example of silver chloride. The solubility product of silver chloride at 25°C is 1.56 x lO". This means that at saturation we can dissolve 1.25 x 10 mol of AgCl/lOOOg of water. If, however, we were to start with a solution that has a coneentration of 1 X 10 molal NaCl (hence 1 x 10 molal CP) the solubility product equation can be written in the form... 

This scheme is actually much more comphcated than the scientific model of silver ions sticking to chloride ions because of their opposite charges (Figure 4.7) and illustrates just how tenacious some misconceptions can be once they have a hold of a student s imagination. You can download a diagnostic task to identify common alternative conceptions students may hold about how bonds form in a precipitation reaction ( Reaction to form silver chloride ) from the Royal Society of Chemistry website (see the Other resources section at the end of this chapter). 

The procedure may be illustrated by the following simple experiment, which is a modification of the Gay Lussac-Stas method. The sodium chloride solution is added to the silver solution in the presence of free nitric acid and a small quantity of pure barium nitrate (the latter to assist coagulation of the precipitate). 

In their inorganic compounds the three halogens chlorine, bromine and iodine are in this order in regard to their aflSnity for other elements. This seems to hold also in their organic compounds as shown by the replacement of iodine by bromine or chlorine, as given previously, and by the fact that alkyl iodides are the least stable or hemostreactive, while the chlorides are the most stable or the least reactive. This is illustrated by their action upon silver nitrate. Ethyl iodide acts with silver nitrate in alcoholic solution precipitating silver iodide even in the cold. 

Fig. 12 illustrates the titration of sodium chloride with silver nitrate. After all chloride is precipitated, the addition of excess silver nitrate causes a rapid increase in conductivity. The slope of the initial portion of the curve may be either downward or upward, depending on the relative conductance of the ion being determined and the ion of like charge in the reagent that replaces it. Slow reactions and coprecipitation are sources of difficulty with precipitation and complex-formation titrations. 

One laboratory use of precipitation reactions is to determine the presence of certain ions in a solution. An illustrative example is given below. The four colorless solutions containing chloride (Cl-), iodide (I-), sulfide (S2-), and nitrate (NOp ions can be identified by using only a silver nitrate solution. A more detailed ion analysis is given in Appendix B. 

The Effect of Reaction Completeness on Titration Curves Figure 13-5 illustrates the effect of solubility product on the sharpness of the end point in titrations with 0.1 M silver nitrate. Clearly, the change in pAg at the equivalence point becomes greater as the solubility products become smaller, that is, as the reaction between the analyte and silver nitrate becomes more complete. By careful choice of indicator—one that changes color in the region of pAg from 4 to 6—titration of chloride ion should be possible with a minimal titration error. Note that ions forming precipitates with solubility products much larger than about 10 do not yield satisfactory end points. 

Worked example 6.5 illustrates the use of the common-ion effect in gravimetric analysis AgCl is always precipitated from a solution containing a slight excess of a common ion, CP or Ag , in the determination of silver or chloride respectively. 

Various approaches can be adopted to perturb the equilibrium shown in Scheme 3.3 and to favor formation of [Ru(CD3CN)(PR3)2Cp]+. For example, the interaction of [RuCl(PR3)2Cp] and chloride ion acceptors, such as silver, sodium, or ammonium hexafluorophosphates, results in precipitation of AgCl, NaCl, or NH4CI and 100% conversion of the starting compxmnd into the solvento-complex. The reaction with NaPFg is illustrated in Scheme 3.4. 

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