Spectrophotometric Titration of a Mixture of Calcium and Magnesium Roberl Fulton, Michael Ross, and Karl Schroeder Saint John's University, Collegeville, M N 56321 T h e advantages of spectrophotometric titrations have been well established (1,2) and many experiments using the (3-6). technioue have been oublished in THIS JOURNAL these experiments have generally been developed for an instrumental analvsis lahoratorv a n d often require special instrumentation. This experiment describes a manual titration m i n e a Soectronic 20 and a manually operated . for 2-mL buret. T h e use of standard equipment access bv t h e large numbers of students usually found i n introductory quantitative analysis courses. T h e importance of including KIY1'A titrations in the undergraduate chemistry cttrriru&n has long been rrcognized (7) and many experiments have been developed to help illustrate t h r chemistry of such systrms (8). T h e spectrophotometric titration described in this paper does an effective ~ o b of demonstratine the orincioles underlvine EDTA f i r rations and helps s t u d e i t s t;visuaiize t h e chekical changes taking d a c e durine" the course of t h e titration. Bv interrelating- t h e optical properties of the various chemical species with their chemical properties. students gain a much deeper under. . standing of concepts'rrlated to hasking, choice ($indicator, choice of nH, and choice of wavelength. Additionally. because the conditions chosen for the titrations are n i t well suited t o visual endpoints (i.e., dilute solutions), students become aware of t h e unique advantage of spectrophotometric titrations.
ow ever,
Apparatus The spectrophotometer used in this study was a Bausch and Lomb Spectronic 20 adapted to accommodate 1-in. optical cells. The titration cell was a 1-in. optical cell with a %-in. magnetic stirring bar in the bottom (because the stirring bar stays below the light pathway, it can be left in the eell during the optical measurement). The buret used was a canventional 2-mL glass microburet (e.g., Sargent-Welch Cat. No. S-10945-50B),but we havealso useda Gilmont 2-mL micrometer buret (e.g., American Scientific Products Cat. No. P5166-2). After each addition of an increment of titrant, mixing of the sample was accomplished with a conventional magnetic stirrer (e.g., Sargent-Welch Cat. No. S-76491-10). The titration eell was removed from the Soectronic 20 for this mixing - and the eell was carefully realigned when placed back into the instrument.'
salt in disrrlled water. The reaultmg sdution was standardi,ed againat the EDTA titranr. Uommal concenrrarion. 0.008.41. Calmagite-lndirntor d u t i m r wrrr prcpnred from purified mlmagite (9) to be ahout 0.1%. Eriochrome Blue SE-The indicator was powdered and shaken with NaClk100). The solid suspension was added as needed. ~ u f f epH r 10--6.8 g of ammo"ium chloride and 57 mLof concentrated ammonium hydroxide were added to enough distilled water to make 100 mL. Buffer pH 12.8-A 0.5 M solution of sodium hydroxide was used to adjust the pH to about 12.8for the titrations of calcium. Procedures Selection of the Wavelength The most sensitive wavelength to use for the titration is determined by obtaining the absorption spectra of the free indicator (no metal ion present) and the compleaed indicator at the appropriate pH. For example, in the titration of calcium and magnesium at pH 10 using Calmagite, the wavelength used was 605 nm, the wavelength where the absorption difference was the greatest (see Fig. 1). Other conditions used are shown in Table 1. Titration of the Sample The appropriate amount of calcium or magnesium (see Tables 2 and 3) was obtained by pipetting the proper amount of the carresponding stock (or diluted stack) solutions into a l-in. spectrophotameter cell. Buffer or 0.5MNaOH (see Table 1) was added and the volume was adjusted to about 30 ml with distilled water. In order to prevent loss of NH3 from the pH 10 buffer or absorption of COI by the pH 12.8 buffer, the cell should he stoppered or covered wlth Parafilm while the spectraphotmetric measurements are made. The cell was inserted into the spectrophotometer and the absorbance adjusted to zero. The appropriate indicator was added until the absorbance at the chosen wavelength (see Table 1)was about 0.32 and the sample was titrated with EDTA.
Reagents Unless specified otherwise, all reagents were analytical reagent grade. EDTA-Stock titrant solutions were prepared from primary standard disodium salt (99.7%) by dissolving the salt in distilled water. Nominal concentration, 0.005 M. Calcium-Stock solutions of calcium ion were prepared from primary standard calcium carbonate (100%)by dissolving the salt in a minimum of 0.5 M hydrochloric acid and diluting with distilled water. Nominal concentration, 0.W8 M. Magnesium-Stock solutions of magnesium ion were prepared from magnesium chloride by dissolving the appropriate amount of
'
In some cases, the Spectronic 20 can be raised and the magnetic stirrer placed under the cell compartment. This greatly facilitates the mixing operation because the cell need not be removed. For - best orecision the absorbance should be adiusted to the value that will provide an absorbance of about 0.9 at the ind of the titration. This is best found empirically using a practice titration. ~~~
7
~~
~
WAVELENGTH, N M
Figure 1. Absorption spectra of calmagite (- 3.3 X free dye, 6-with excess MgZ+. Volume 63 Number 8
10C Mat pH
August 1986
= 10. A-
721
Table 1. THratlon Condltlons for Calclum Only and Calclum Plus Magneslum Sample
Nominal pH
Buffer Reagent
Indicator
Wavelengm (nm)
Ca only Ca Mg
12.8 10.0
2 mL 0.5 MNaOH l o mL NHrNH,Cl
Eriachmme Blue SE Calmaglla
605. 605
+
'The abmbbance pssk fa me t e e dye occurs at ebout 610 nm, but at 605 nm
Table 2.
Expt.
pH
indicator
I I1 111 IV
12.8 12.8 12.2 13.2
EBSEa EBSE EBSE EBSE
15.66 16.13 16.13 16.13
Expt. V Vi VII Vlll IX X a
pH
indicator
pmol Cataken
10.0 10.0 12.8 12.8 12.8 12.8
CALe CAL EBSEb EBSE EWE EBSE
15.78 16.13 15.78 16.13 16.13 16.13
calclumcomplexeddye Is a maximum.
Tltratlon ol Sample$ Contalnlng Calclum Only
pmoi Ca taken
Table 3.
aboorbanu, dlfferarca batween W hn, dye and
Average pmoi
Ca found 15.68 16.06 16.07 15.97
Absolute Experimental Error3 0.02 -0.07 -0.06 -0.16
Number of Measurements
Relative Standard DBviation of Measurement (%)
6 5 9 2
0.1 0.1 0.2 0.4
( 0.1) (-0.4) (-0.4) (-1.0)
THratlon of Samples Contalnlng Calclum and Magneslum
am01 Mgtaken 2.61 2.61 2.61 2.61 17.40 87.00
Absolute Experimental ErrorC
Average pmol found 18.35 18.85 15.79 15.92 16.00 15.27
(Ca+Mg) (Ca+Mg) (Ca only) (Caonly) (Ca only) (Ca only)
,
-0.04 0.11 0.01 -0.21 -0.13 -0.86
Number of Measurements
Relative Standard Deviationof Measurement (%)
3 4 2 6 5 2
0.1 0.5 0.0 1.0 0.2 0.1
(-0.2) ( 0.6) ( 0.1) (-1.3) (-0.8) (-5.3)
CAL-Calmsglfe
'EBSE--Eriochrome Blw SE " R~Ia1Iveexperimentalarmr (percent1In perenmeslr
The ahsorbance was measured after the addition of each increment of titrant. As the titration approached the endpoint, as indicated by arapid increase in the ahsorbance, the size of the increment was reduced (by immersing the tip of the buret, increments as small as about 0.01-mL can be added). The titration is continued for about 0.5-1.0mL beyond the endpoint or until a well-definedlinear region is established (see Fig. 2). The endpoint is obtained by plotting the absorbance against the corresponding volume of titrant as shown in Figure 2. The endpoint volume is that which corresponds to the intersection of the rapidly rising part of the curve before the endpoint with the linear region of the curve beyond the endpoint. Because the precision of the endpoint depends on the definition of the curve in the vicinity of the endpoint, care must be taken to insure that sufficientdata is taken in that region. Results
Tables 2 and 3 summarize the results of a number of different experiments involving the titration of dilute solutions of calcium alone. mixtures of calcium and mamesium. and calcium in the presence of magnesium. The results indi-. cate that withcareful workit i s ~ o s s i h ltotitrate e these kinds of samples to a precision of a few parts per thousand. Because the visual endpoints are not sharply defined (pronounced curvature in the curve prior to the endpoint, see Fig. 2), visual titration of these s a m ~ l e as t com~arahleconcentration levels can be done only with great difficulty. In addition, the data show that calcium may he titrated in the presence of magnesium as long as the magnesium concentration is not significantly larger than the calcium concentration and the p H is held above 12.2. ~~
722
Journal of Chemical Education
~~
--------. -
I ENDPOINT
I 1 i
I 1.2
I
11
1.6
2.0
I 2.4
VOLUME, M I Figure 2. Typical endpoint region of the tihation curve for calcium with EDTA using an Eriochrome Blue SE indicator at pH = 12.8 and wavelengm = 605.
Calcium Only
The results of Experiments I-IV indicate that calcium may he titrated satisfactorily in the pH range of about 12.2 to 12.8 using Eriochrome Blue S E as the indicator. At pH's greater than about 12.9, for example, 13.2 (Experiment IV), significant amounts of calcium are lost due to the probable precipitation of c a ( 0 H ) ~ . Calclum and Magnesium
Exoeriments V and VI (Table 3) summarize the results for the titration of a mixture'of calc&m and magnesium a t pH 10.0 using calmagite as the indicator. At this pH hoth are titrated and the experimental results for the sum of the two agree well with the theoretical amounts taken. Calcium in the Presence of Magnesium
The final four Experiments in Table 3 (Experiments V I E X) show the results of titrating calcium in the presence of magnesium. At the pH used in ihese experiments (12.8),the magnesium is masked bv ~recipitntionas MdOHh. When themagnesium to calcium ratio is lower th& about one to one, the results for calcium are very good. However, when the ratio gets as high as five t o one (Experiment X), the results for calcium are low, indicating that significant Z apparently co-precipitated with amounts of C ~ ( O H )have the Mg(OHI2.
dilute and not easy to titrate visually using standard procedures.3 In addition t o helping students understand the advantaaes of suectrouhotometric titrations. this exoeriment demoistrate; the Eoncept of masking and the heed for choosing a metal indicator that is matched t o the pH of the solution being titrated. Finally, the experiment illustrates how the s ~ e c t r aorooerties l of the indicator snecies and the titration [eadions are related t o the corresponding spectrophotometric titration curves. We have used this experiment as a part of theintroductory quantitative analvsis lahoratorv for a number of vears and &e pleased with the results and level of understadding that s students demonstrate on the written laboratory r e ~ o r t that they submit. Acknowledgment
The authors wish t o thank the many students who have oarticioated in these exoeriments. S u ~ ~ obvr the t Research eorpoiation, ~ c ~ h e r s Research in ~ e l l o w s h iFund, ~ and Saint John's University is also gratefully acknowledged. (11 U n d e d . A. J. Chem. Edue. 19S4,31.394. (21 Kolthoff, I. M.; Eluing, P. J. '"Reatlae on Analytical Chemistry: Wilni: New York, 1961: Pt I. Yo1 5, p 2795. (3) Sellers, D. E.:Tang, S. S. N.: VanAtb.R. E. J. Chem. Educ. 1962.39, 4m. (4) Beilhy, A. L.:Landomki, C. A. J. Chm. Edue. 1970.47, 238. (5) Olsen. E. D.;Forehack. C. C. J Chem. Edue. 1972.49,206. (61 Arndur.S.;Levene, W . J. J ChemEdue. 1974,51,136. (I1 Johnaton, M. B.;Bamard,A. J. J. Chem. Edur. 1958,3S, 601. ( 8 ) Ramasy.C. G.J. Chem. Educ. 1917.54.714. (9) Kratochvil, B.: Nolan, J.:Cantwell, F. F.; Fulton, R. B. Con. J. Chem 198L.59.2539.
Conclusrom
The experiments described in this paper can be used to titrate calcium and calcium-magnesium mixtures that are
We generally have students titrate these samples using standard visual techniques in order to compare the results of the two methods.
Volume 63
Number 8
August 1986
723
