D. Harvey - Modern Analytical Chemistry (794078), страница 72
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Interferents forming precipitates that areless soluble than the analyte may be precipitated and removed by filtration, leavingthe analyte behind in solution. Alternatively, either the analyte or the interferentcan be masked using a suitable complexing agent, preventing its precipitation.occlusionA coprecipitated impurity trappedwithin a precipitate as it forms.digestionThe process by which a precipitate isgiven time to form larger, purerparticles.adsorbateA coprecipitated impurity that adsorbsto the surface of a precipitate.CACACACACACACACACACAACACACACACACACACACACCACACACAMACACACACACAACACACACACACAMACACACCACACACACACACACACACAAMACACACACACACACACACCACACACACACACACACACA(a)CACACACACACACACACACAACACACACACACACACACACCACACACACACACACACACAACACACACACACACACACACCACACACACACACACACACAACACACACACACACACACACCACACACACACACACACACA(b)CACACCCCC CACACACAC CCACACACACACACACCACACACACACACACACACAC(c)Figure 8.4Example of coprecipitation: (a) schematic ofa chemically adsorbed inclusion or aphysically adsorbed occlusion in a crystallattice, where C and A represent thecation–anion pair comprising the analyteand the precipitant, and M is the impurity;(b) schematic of an occlusion by entrapmentof supernatant solution; (c) surfaceadsorption of excess C.1400-CH08 9/9/99 2:17 PM Page 240240Modern Analytical ChemistryBoth of the above-mentioned approaches are illustrated in Fresenius’s analytical method for determining Ni and Co in ores containing Pb2+, Cu2+, and Fe3+ aspotential interfering ions (see Figure 1.1 in Chapter 1).
The ore is dissolved in a solution containing H2SO4, selectively precipitating Pb2+ as PbSO4. After filtering, thesupernatant solution is treated with H2S. Because the solution is strongly acidic,however, only CuS precipitates. After removing the CuS by filtration, the solution ismade basic with ammonia until Fe(OH)3 precipitates. Cobalt and nickel, whichform soluble amine complexes, remain in solution.In some situations the rate at which a precipitate forms can be used to separatean analyte from a potential interferent.
For example, due to similarities in theirchemistry, a gravimetric analysis for Ca2+ may be adversely affected by the presenceof Mg2+. Precipitates of Ca(OH)2, however, form more rapidly than precipitates ofMg(OH)2. If Ca(OH)2 is filtered before Mg(OH)2 begins to precipitate, then aquantitative analysis for Ca2+ is feasible.Finally, in some cases it is easier to isolate and weigh both the analyte and theinterferent. After recording its weight, the mixed precipitate is treated to convert atleast one of the two precipitates to a new chemical form. This new mixed precipitateis also isolated and weighed.
For example, a mixture containing Ca2+ and Mg2+ canbe analyzed for both cations by first isolating a mixed precipitate of CaCO3 andMgCO3. After weighing, the mixed precipitate is heated, converting it to a mixtureof CaO and MgO. ThusGrams of mixed precipitate 1 = grams CaCO3 + grams MgCO3Grams of mixed precipitate 2 = grams CaO + grams MgOAlthough these equations contain four unknowns (grams CaCO3, grams MgCO3,grams CaO, and grams MgO), the stoichiometric relationships between CaCO3 andCaOMoles CaCO3 = moles CaOand between MgCO3 and MgOMoles MgCO3 = moles MgOprovide enough additional information to determine the amounts of both Ca2+ andMg2+ in the sample.*Controlling Particle Size Following precipitation and digestion, the precipitatemust be separated from the supernatant solution and freed of any remaining impurities, including residual solvent. These tasks are accomplished by filtering, rinsing,and drying the precipitate.
The size of the precipitate’s particles determines the easeand success of filtration. Smaller, colloidal particles are difficult to filter becausethey may readily pass through the pores of the filtering device. Large, crystallineparticles, however, are easily filtered.By carefully controlling the precipitation reaction we can significantly increasea precipitate’s average particle size. Precipitation consists of two distinct events: nucleation, or the initial formation of smaller stable particles of precipitate, and thesubsequent growth of these particles. Larger particles form when the rate of particlegrowth exceeds the rate of nucleation.*Example 8.2 shows how to solve this type of problem.1400-CH08 9/9/99 2:17 PM Page 241Chapter 8 Gravimetric Methods of AnalysisA solute’s relative supersaturation, RSS, can be expressed asQ−SRSS =S8.12where Q is the solute’s actual concentration, S is the solute’s expected concentration at equilibrium, and Q – S is a measure of the solute’s supersaturation whenprecipitation begins.3 A large, positive value of RSS indicates that a solution ishighly supersaturated.
Such solutions are unstable and show high rates of nucleation, producing a precipitate consisting of numerous small particles. WhenRSS is small, precipitation is more likely to occur by particle growth than bynucleation.Examining equation 8.12 shows that we can minimize RSS by either decreasingthe solute’s concentration or increasing the precipitate’s solubility. A precipitate’ssolubility usually increases at higher temperatures, and adjusting pH may affect aprecipitate’s solubility if it contains an acidic or basic anion.
Temperature and pH,therefore, are useful ways to increase the value of S. Conducting the precipitation ina dilute solution of analyte, or adding the precipitant slowly and with vigorous stirring are ways to decrease the value of Q.There are, however, practical limitations to minimizing RSS. Precipitates thatare extremely insoluble, such as Fe(OH)3 and PbS, have such small solubilities thata large RSS cannot be avoided.
Such solutes inevitably form small particles. In addition, conditions that yield a small RSS may lead to a relatively stable supersaturatedsolution that requires a long time to fully precipitate. For example, almost a monthis required to form a visible precipitate of BaSO4 under conditions in which the initial RSS is 5.4An increase in the time required to form a visible precipitate under conditionsof low RSS is a consequence of both a slow rate of nucleation and a steady decreasein RSS as the precipitate forms.
One solution to the latter problem is to chemicallygenerate the precipitant in solution as the product of a slow chemical reaction. Thismaintains the RSS at an effectively constant level. The precipitate initially formsunder conditions of low RSS, leading to the nucleation of a limited number of particles. As additional precipitant is created, nucleation is eventually superseded by particle growth. This process is called homogeneous precipitation.5Two general methods are used for homogeneous precipitation. If the precipitate’s solubility is pH-dependent, then the analyte and precipitant can be mixedunder conditions in which precipitation does not occur.
The pH is then raised orlowered as needed by chemically generating OH– or H3O+. For example, the hydrolysis of urea can be used as a source of OH–.CO(NH2)2(aq) + H2O(l)NH3(aq) + H2O(l)NH2SO3H(aq) + 2H2O(l)relative supersaturationA measure of the extent to which asolution, or a localized region ofsolution, contains more dissolved solutethan that expected at equilibrium (RSS).homogeneous precipitationA precipitation in which the precipitantis generated in situ by a chemicalreaction.t CO2(g) + 2NH3(aq)t NH4+(aq) + OH–(aq)The hydrolysis of urea is strongly temperature-dependent, with the rate being negligible at room temperature.
The rate of hydrolysis, and thus the rate of precipitateformation, can be controlled by adjusting the solution’s temperature. Precipitates ofBaCrO4, for example, have been produced in this manner.In the second method of homogeneous precipitation, the precipitant itself isgenerated by a chemical reaction. For example, Ba2+ can be homogeneously precipitated as BaSO4 by hydrolyzing sulphamic acid to produce SO42–.t NH4+(aq) + H3O+(aq) + SO42–(aq)241Color Plate 5 shows the differencebetween a precipitate formed by directprecipitation and a precipitate formedby a homogeneous precipitation.1400-CH08 9/9/99 2:17 PM Page 242242Modern Analytical ChemistryNO3– OH–Ag+NO3–NO3–PrimaryadsorptionlayerPrecipitate–NO3Ag++AgClAgClAgClCl–Ag+NO3–AgClAgCl+AgClAgClAgClOH–AgClAgClAgClAg+ClAgClAgClAgNO3–AgClAgClAgClAg+ClAgClAgClAg+NO3–AgClAgClAgClAgClAgClAg+NO3–AgClAgClOH–Ag+Ag+Figure 8.5Schematic model of the solid–solutioninterface at a particle of AgCl in a solutioncontaining excess AgNO3.coagulationThe process of smaller particles ofprecipitate clumping together to formlarger particles.–NO3– NO3OH–NO3–NO3–Bulk solutionSecondaryadsorptionlayerHomogeneous precipitation affords the dual advantages of producing largeparticles of precipitate that are relatively free from impurities.
These advantages,however, may be offset by increasing the time needed to produce the precipitate,and a tendency for the precipitate to deposit as a thin film on the container’s walls.The latter problem is particularly severe for hydroxide precipitates generated usingurea.An additional method for increasing particle size deserves mention. When aprecipitate’s particles are electrically neutral, they tend to coagulate into larger particles. Surface adsorption of excess lattice ions, however, provides the precipitate’sparticles with a net positive or negative surface charge. Electrostatic repulsion between the particles prevents them from coagulating into larger particles.Consider, for instance, the precipitation of AgCl from a solution of AgNO3,using NaCl as a precipitant.
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