D. Harvey - Modern Analytical Chemistry (794078), страница 95
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Hardness is determined by titrating withEDTA at a buffered pH of 10. Eriochrome Black T or calmagite is used as a visualindicator. Hardness is reported in parts per million CaCO3.Procedure. Select a volume of sample requiring less than 15 mL of titrant to keepthe analysis time under 5 min and, if necessary, dilute the sample to 50 mL withdistilled water. Adjust the pH by adding 1–2 mL of a pH 10 buffer containing a smallamount of Mg2+–EDTA. Add 1–2 drops of indicator, and titrate with a standardsolution of EDTA until the red-to-blue end point is reached.Questions1. Why is the sample buffered to a pH of 10? What problems might be expectedat higher or lower pHs?Of the cations contributing to hardness, Mg2+ forms the weakest complex withEDTA and is the last cation to be titrated. Calmagite was selected as theindicator because it gives a distinct end point with Mg2+. Because of calmagite’sacid–base properties the indicator is only useful in the pH range of 9–11 (seeTable 9.16).
Figure 9.31 shows the titration curve for a solution of 10–3 M Mg2+with 10–2 M EDTA at pHs of 9, 10, and 11. Superimposed on each titration curveis the range of conditions in which the average analyst will find the end point.At a pH of 9 an early end point is possible, leading to a negative determinateerror, and at a pH of 11 there is a chance of a late end point and a positivedeterminate error.pMg10.09.08.07.06.05.04.03.02.01.00.0Early end point10.09.08.07.06.05.04.03.02.01.00.00.00 2.00 4.00 6.00 8.00 10.00Volume titrant (mL)0.00 2.00 4.00 6.00 8.00 10.00Volume titrant (mL)(a)(b)pMgThe photo in Colorplate 8b shows theindicator’s color change for this titration.Representative MethodsModern Analytical ChemistrypMg326Figure 9.3110–3Mg2+10–2Titration curves forMwithMEDTA using calmagite as an indicator at(a) pH = 9, (b) pH = 10, and (c) pH = 11.
Therange of pMg and volume of titrant overwhich the indicator is expected to changecolor is shown for each titration curve.10.09.08.07.06.05.04.03.02.01.00.0Late end point0.00 2.00 4.00 6.00 8.00 10.00Volume titrant (mL)(c)—Continued1400-CH09 9/9/99 2:13 PM Page 327Chapter 9 Titrimetric Methods of Analysis2. Why is a small amount of Mg2+–EDTA complex added to the buffer?The titration’s end point is signaled by the indicator calmagite, which gives agood end point with magnesium, but a poor end point with other cations suchas calcium.
If the sample does not contain any Mg2+ as a source of hardness,then the titration will have a poorly defined end point and inaccurate resultswill be obtained. By adding a small amount of Mg2+–EDTA to the buffer, asource of Mg2+ is ensured. When the buffer is added to the sample, the Mg2+ isdisplaced by Ca2+, because Ca2+ forms a stronger complex with EDTA.
Since thedisplacement is stoichiometric, the total concentration of hardness cationsremains unchanged, and there is no change in the amount of EDTA needed toreach the equivalence point.3. Why does the procedure specify that the titration take no longer than 5 min?The presence of a time limitation suggests that there must be a kineticallycontrolled interference, possibly arising from a competing chemical reaction. Inthis case the interference is the possible precipitation of CaCO3.9C.5 Quantitative ApplicationsWith a few exceptions, most quantitative applications of complexation titrimetryhave been replaced by other analytical methods.
In this section we review the general application of complexation titrimetry with an emphasis on selected applications from the analysis of water and wastewater. We begin, however, with a discussion of the selection and standardization of complexation titrants.Selection and Standardization of Titrants EDTA is a versatile titrant that can beused for the analysis of virtually all metal ions. Although EDTA is the most commonly employed titrant for complexation titrations involving metal ions, it cannotbe used for the direct analysis of anions or neutral ligands.
In the latter case, standard solutions of Ag+ or Hg2+ are used as the titrant.Solutions of EDTA are prepared from the soluble disodium salt,Na2H2Y • 2H2O. Concentrations can be determined directly from the known massof EDTA; however, for more accurate work, standardization is accomplished bytitrating against a solution made from the primary standard CaCO3. Solutions ofAg+ and Hg2+ are prepared from AgNO3 and Hg(NO3)2, both of which are secondary standards. Standardization is accomplished by titrating against a solutionprepared from primary standard grade NaCl.Inorganic Analysis Complexation titrimetry continues to be listed as a standardmethod for the determination of hardness, Ca2+, CN–, and Cl– in water and wastewater analysis. The evaluation of hardness was described earlier in Method 9.2.
Thedetermination of Ca2+ is complicated by the presence of Mg2+, which also reactswith EDTA. To prevent an interference from Mg2+, the pH is adjusted to 12–13,precipitating any Mg2+ as Mg(OH)2. Titrating with EDTA using murexide or Eriochrome Blue Black R as a visual indicator gives the concentration of Ca2+.Cyanide is determined at concentrations greater than 1 ppm by making thesample alkaline with NaOH and titrating with a standard solution of AgNO3,forming the soluble Ag(CN) 2 – complex. The end point is determined usingp-dimethylaminobenzalrhodamine as a visual indicator, with the solution turning from yellow to a salmon color in the presence of excess Ag+.3271400-CH09 9/9/99 2:13 PM Page 328328Modern Analytical ChemistryChloride is determined by titrating with Hg(NO3)2, forming soluble HgCl2.The sample is acidified to within the pH range of 2.3–3.8 where diphenylcarbazone,which forms a colored complex with excess Hg2+, serves as the visual indicator.
Xylene cyanol FF is added as a pH indicator to ensure that the pH is within the desiredrange. The initial solution is a greenish blue, and the titration is carried out to apurple end point.Quantitative Calculations The stoichiometry of complexation reactions is given bythe conservation of electron pairs between the ligand, which is an electron-pairdonor, and the metal, which is an electron-pair acceptor (see Section 2C); thusmoles of electron pairs donated× moles ligand =mole ligandmoles of electron pairs accepted× moles metalmole metalThis is simplified for titrations involving EDTA where the stoichiometry is always1:1 regardless of how many electron pairs are involved in the formation of themetal–ligand complex.EXAMPLE 9.8The concentration of a solution of EDTA was determined by standardizingagainst a solution of Ca 2+ prepared from the primary standard CaCO 3.
A0.4071-g sample of CaCO3 was transferred to a 500-mL volumetric flask,dissolved using a minimum of 6 M HCl, and diluted to volume. A 50.00-mLportion of this solution was transferred into a 250-mL Erlenmeyer flask and thepH adjusted by adding 5 mL of a pH 10 NH3–NH4Cl buffer containing a smallamount of Mg 2+–EDTA. After adding calmagite as a visual indicator, thesolution was titrated with the EDTA, requiring 42.63 mL to reach the endpoint. Report the molar concentration of the titrant.SOLUTIONConservation of electron pairs for the titration reaction requires thatMoles EDTA = moles Ca2+Making appropriate substitution for the moles of EDTA and Ca2+ givesMEDTA × VEDTA = MCa × VCaThe molarity of theCa2+solution ismoles CaCO3g CaCO30.4071 g==VflaskFW CaCO3 × Vflask100.09 g/mol × 0.5000 L= 8.135 × 10 −3 M Ca 2+Substituting known values and solving for MEDTA givesM CaVCa(8.135 × 10 −3 M)(50.00 mL)== 9.541 × 10 −3 M EDTAVEDTA42.63 mL1400-CH09 9/9/99 2:13 PM Page 329Chapter 9 Titrimetric Methods of AnalysisThe principle of the conservation of electron pairs is easily extended to other complexation reactions, as shown in the following example.EXAMPLE 9.9The concentration of Cl– in a 100.0-mL sample of water drawn from a freshwater acquifer suffering from encroachment of sea water, was determined bytitrating with 0.0516 M Hg(NO3)2.
The sample was acidified and titrated to thediphenylcarbazone end point, requiring 6.18 mL of the titrant. Report theconcentration of Cl– in parts per million.SOLUTIONConservation of electron pairs requires thatMoles Cl– = 2 × moles Hg2+Making appropriate substitutions for the moles of Cl– and Hg2+g Cl −= 2 × M Hg × VHgAW Cland rearranging leaves us withg Cl– = 2 × MHg × VHg × AW ClSubstituting known values and solving gives2 × 0.0516 M × 0.00618 L × 35.453 g/mol = 0.0226 g Cl–The concentration of Cl– in parts per million, therefore, ismg Cl −22.6 mg== 226 ppmliter0.1000 LFinally, quantitative problems involving multiple analytes and back titrations alsocan be solved by applying the principle of conservation of electron pairs.EXAMPLE 9.10An alloy of chromel containing Ni, Fe, and Cr was analyzed by acomplexation titration using EDTA as the titrant.
A 0.7176-g sample of thealloy was dissolved in HNO3 and diluted to 250 mL in a volumetric flask. A50.00-mL aliquot of the sample, treated with pyrophosphate to mask the Feand Cr, required 26.14 mL of 0.05831 M EDTA to reach the murexide endpoint. A second 50.00-mL aliquot was treated with hexamethylenetetramineto mask the Cr. Titrating with 0.05831 M EDTA required 35.43 mL to reachthe murexide end point.
Finally, a third 50.00-mL aliquot was treated with50.00 mL of 0.05831 M EDTA, and back titrated to the murexide end pointwith 6.21 mL of 0.06316 M Cu2+. Report the weight percents of Ni, Fe, andCr in the alloy.3291400-CH09 9/9/99 2:13 PM Page 330330Modern Analytical ChemistrySOLUTIONConservation of electron pairs for the three titrations requires that forTitration 1:moles Ni = moles EDTA1Titration 2:moles Ni + moles Fe = moles EDTA2Titration 3:moles Ni + moles Fe + moles Cr + moles Cu = moles EDTA3(Fe, Cr masked)(Cr masked)Note that the third titration is a back titration. Titration 1 can be used todetermine the amount of Ni in the alloy.
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