D. Harvey - Modern Analytical Chemistry (794078), страница 94
Текст из файла (страница 94)
Finally, a smooth curve is drawn to connect the threestraight-line segments (Figure 9.28f). Additional examples of Cd2+–EDTAtitration curves are shown in Figure 9.29. Note in particular, that the change inpCd near the equivalence point is not as great when the titration is carried outat a lower pH or in the presence of a higher concentration of NH3.18.0(a) (c)16.014.0(b)12.0pCd32210.0(c)8.06.04.0(a)2.0(b)0.00.0010.0020.0030.00Volume titrant (mL)40.0050.00Figure 9.29Effect of pH and [NH3] on the titration curve for 50.0 mL of 5.00 × 10–3 MCd2+ with 0.0100 M EDTA.
(a) pH = 10, [NH3] = 0; (b) pH = 7, [NH3] = 0;(c) pH = 10, [NH3] = 0.5 M.9C.3 Selecting and Evaluating the End PointThe equivalence point of a complexation titration occurs when stoichiometrically equivalent amounts of analyte and titrant have reacted.
For titrations involving metal ions and EDTA, the equivalence point occurs when CM and CEDTAare equal and may be located visually by looking for the titration curve’s inflection point.As with acid–base titrations, the equivalence point of a complexation titrationis estimated by an experimental end point. A variety of methods have been used tofind the end point, including visual indicators and sensors that respond to a changein the solution conditions. Typical examples of sensors include recording a potentiometric titration curve using an ion-selective electrode (analogous to measuringpH with a pH electrode),* monitoring the temperature of the titration mixture, andmonitoring the absorbance of electromagnetic radiation by the titration mixture.*See Chapter 11 for a further discussion of ion-selective electrodes.1400-CH09 9/9/99 2:13 PM Page 323Chapter 9 Titrimetric Methods of Analysis323The first two sensors were discussed in Section 9B.3 for acid–base titrations and arenot considered further in this section.Finding the End Point with a Visual Indicator Most indicators for complexationtitrations are organic dyes that form stable complexes with metal ions.
These dyesare known as metallochromic indicators. To function as an indicator for anEDTA titration, the metal–indicator complex must possess a color different fromthat of the uncomplexed indicator. Furthermore, the formation constant for themetal–indicator complex must be less favorable than that for the metal–EDTAcomplex.The indicator, Inm–, is added to the solution of analyte, forming a coloredmetal–indicator complex, MInn-m. As EDTA is added, it reacts first with the free analyte, and then displaces the analyte from the metal–indicator complex, affecting achange in the solution’s color.
The accuracy of the end point depends on thestrength of the metal–indicator complex relative to that of the metal–EDTA complex. If the metal–indicator complex is too strong, the color change occurs after theequivalence point. If the metal–indicator complex is too weak, however, the endpoint is signaled before reaching the equivalence point.Most metallochromic indicators also are weak acids or bases.
The conditional formation constant for the metal–indicator complex, therefore, depends onthe solution’s pH. This provides some control over the indicator’s titration error.The apparent strength of a metal–indicator complex can be adjusted by controllingthe pH at which the titration is carried out. Unfortunately, because they also areacid–base indicators, the color of the uncomplexed indicator changes with pH. Forexample, calmagite, which we may represent as H3In, undergoes a change in colorfrom the red of H2In– to the blue of HIn2– at a pH of approximately 8.1, and fromthe blue of HIn2– to the red-orange of In3– at a pH of approximately 12.4. Since thecolor of calmagite’s metal–indicator complexes are red, it is only useful as a metallochromic indicator in the pH range of 9–11, at which almost all the indicator ispresent as HIn2–.A partial list of metallochromic indicators, and the metal ions and pH conditions for which they are useful, is given in Table 9.16.
Even when a suitable indicator does not exist, it is often possible to conduct an EDTA titration by introducing asmall amount of a secondary metal–EDTA complex, provided that the secondarymetal ion forms a stronger complex with the indicator and a weaker complex withEDTA than the analyte. For example, calmagite can be used in the determination ofTable 9.16Selected Metallochromic IndicatorsIndicatorcalmagiteEriochrome Black TEriochrome Blue Black RmurexidePANsalicylic acidUseful pH Range9–117.5–10.58–126–132–112–3Useful forBa, Ca, Mg, ZnBa, Ca, Mg, ZnCa, Mg, Zn, CuCa, Ni, CuCd, Cu, ZnFemetallochromic indicatorA visual indicator used to signal the endpoint in a complexation titration.1400-CH09 9/9/99 2:13 PM Page 324324Modern Analytical ChemistryCa2+ if a small amount of Mg2+–EDTA is added to the solution containing the analyte.
The Mg2+ is displaced from the EDTA by Ca2+, freeing the Mg2+ to form thered Mg2+–indicator complex. After all the Ca2+ has been titrated, Mg2+ is displacedfrom the Mg2+–indicator complex by EDTA, signaling the end point by the presence of the uncomplexed indicator’s blue form.Finding the End Point by Monitoring Absorbance. An important limitationwhen using a visual indicator is the need to observe the change in color signaling the end point.
This may be difficult when the solution is already colored.For example, ammonia is used to adjust the pH of solutions containing Cu2+before its titration with EDTA. The presence of the intensely coloredCu(NH3)42+ complex obscures the indicator’s color, making an accurate determination of the end point difficult.
Other absorbing species present within thesample matrix may also interfere in a similar fashion. This is often a problemwhen analyzing clinical samples such as blood or environmental samples suchas natural waters.As long as at least one species in a complexation titration absorbs electromagnetic radiation, the equivalence point can be located by monitoring the absorbance of the analytical solution at a carefully selected wavelength.* For example, the equivalence point for the titration of Cu 2+ with EDTA, in thepresence of NH3, can be located by monitoring the absorbance at a wavelengthof 745 nm, where the Cu(NH3)42+ complex absorbs strongly. At the beginningof the titration the absorbance is at a maximum. As EDTA is added, however,the reactionCu(NH3)42+(aq) + Y4–(aq) → CuY2–(aq) + 4NH3(aq)occurs, decreasing both the concentration of Cu(NH3)42+ and the absorbance.The absorbance reaches a minimum at the equivalence point and remains essentially unchanged as EDTA is added in excess. The resulting spectrophotometrictitration curve is shown in Figure 9.30a.
In order to keep the individual segmentsof the titration curve linear, the measured absorbance, Ameas, is corrected fordilutionAcorr = Ameas ×VEDTA + VCuVCuwhere Acorr is the corrected absorbance, and VEDTA and VCu are, respectively, thevolumes of EDTA and Cu. The equivalence point is given by the intersection of thelinear segments, which are extrapolated if necessary to correct for any curvature inthe titration curve. Other common spectrophotometric titration curves are shownin Figures 9.30b–f.9C.4 Representative MethodAlthough every complexation titrimetric method has its own unique considerations, the following description for determining the hardness of water provides aninstructive example of a typical procedure.*See Chapter 10 for a further discussion of absorbance and spectroscopy.1400-CH09 9/9/99 2:13 PM Page 325Volume of titrant(a)Volume of titrantCorrected absorbanceCorrected absorbance(b)Volume of titrantVolume of titrant(d)orCorrected absorbanceCorrected absorbance(c)Volume of titrant(e)325Corrected absorbanceCorrected absorbanceChapter 9 Titrimetric Methods of AnalysisorVolume of titrant(f)Figure 9.30Spectrophotometric titration curves for thetitration of an analyte, A, with a titrant, T, toform a product, P, in the presence of a visualindicator.
Titration curves are shown forcases where (a) only A absorbs; (b) only Tabsorbs; (c) only P absorbs; (d) A and Tabsorb; (e) P and T absorb; and (f) only thevisual indicator absorbs.1400-CH09 9/9/99 2:13 PM Page 326Method 9.2Determination of Hardness of Water and Wastewater12Description of the Method.
The operational definition of water hardness is thetotal concentration of cations in a sample capable of forming insoluble complexeswith soap. Although most divalent and trivalent metal ions contribute to hardness,the most important are Ca2+ and Mg2+.
![](https://s.studizba.com/z.php?f=/templates/si/i/noavatar.png&w=100&h=100&t=1"){kind=avatar}
