EDTA AND COMPLEX FORMATION A Demonstration Lecture M. B. JOHNSTON and A. J. BARNARD, Jr. J. T. Baker Chemical Company, Phillipsburg, New Jersey
H. A. FLASCHKA Georgia Institute of Technology, Atlanta, Georgia
BYTHEIR nature general textbooks often lag by many years in their inclusion of salient advances in analytical reagents. Teachers are aware of this lag when confronted with the necessity of up-dating lectures and courses. Often this requires tedious collection and digestion of data scattered through the world literature. Such is the case with EDTA. Only a few current textbooks devote space to this versatile titrant, masking agent, and chromogenic agent. A few books confine their remarks to the EDTA titration of water hardness, but give only limited background information even for this special application. It is the purpose of this paper to aid in resolving this problem by presenting some integrated experiments designed to serve as a general introduction to watersoluble chelates, and especially to EDTA and its use in analytical chemistry. These experiments also delineate some special aspects of metal complexes that are sometimes a source of practical difficulty, and that, on the other hand, can sometimes be applied to advantage in chemical analysis. The experiments are scaled to serve as lecture demonstrations; however, they can he readily adapted to student use by employing the same reagent solutions and reducing the volumes employed appropriately (with the exception of Experiment I X which can he performed as given). All of the reagent solutions, with two exceptions, can be readily prepared from reagent chemicals availahle
(After Martell)
VOLUME 35, NO. 12, DECEMBER, 1958
in most instructional laboratories. The two special reagents, disodium ethylenediaminetetraacetate dihydrate and Eriochrome Black T (Erie T for short), are readily available through laboratory supply houses. The intention is not merely t o give 'Lcook-book" recipes but also to provide some background information by discussion of the experimental phenomena. This information can only he presented briefly and in summary form. Hence, a selected bibliography is appended that cites some relevant monographs (1-4), reviews (5-11), trade literature (13-id), and a few papers of historic interest (10,15-17). EDTA is the alphabetic designation for ethylenediaminetetraacetic acid or its anion form ethylenediaminetetraacetate. HOOC-H,C )N.cH~cH,N< HOOGHC
CHrCOOR CHrCOOH
Ethylenediaminetetraacetic acid
NaOOC.H&
CHrCOONa
[ HWC CH&OOH
Disodium ethylenediaminetetrrtacetilte dihydrate
Because of the very low water solubility of the free acid, a solution of the disodium salt is usually employed as the reagent. The usefulness of EDTA stems from the presence of six ligand groups which permits honding with an equal number of coordination positions of a metal ion. I n almost all cases a 1:l metal-EDTA complex (chelate) is formed and in a single step. The presence of carboxylic acid groups confers water solubility on such metal complexes. The steric configuration (18) of a metal-EDTA complex is shown in the figure. A fuller treatment of the theoretical background for the use of EDTA will be found in works cited in the bibliography (1-4, 10-11,16-18). The following sections of this paper present the directions for the preparation of the required reagents, followed by the experiments. The experiments are arranged so that the principles or phenomena to be studied are briefly summarized followed by working directions into which are interpolated brief statements concerning the observations. The final experiment (IX) is an actual EDTA titration. Rather than a study of the simple titration of a
single metal ion in pure solution, a slightly more complicated example has been selected which familarizes the student with such phenomena as indicator-blocking and masking. This titration can be performed either as a lecture experiment or as an actual student laboratory determination (using, if desired, "unknown" solutio& or mixtures of magnesium and iron(II1) salts). REAGENTS
The following directions yield amounts sufficient to conduct Experiments I through VIII twice, and Experiment IX about twenty times if it is to be conducted quantitatively as a student exercise. All chemicals employed should be of reagent grade. EDTA, 0.1 M . Dissolve 18.6 g. of disodium ethylenediaminp tetraacetate dihydrate; dilute to 500 ml. with H20. Stock water. Use exclusively water free of metal ions, either carefully redistilled or deionized via cation exchange. (Often "distilled" water contains interfering amounts of metals.) For demonstration purposes (is., except in Expt. IIID or in Expt. IX, when conducted quantitrttively), ordinary distilled water may be used if interferences are excluded by addition of 5 drops of 0.1 M EDTA per liter. Erio T solution. Dissolve 0.1 g. of Eriochrome Black T in 25 ml. of methanol (must be freshly prepared); alternatively, use a 1: 100 ground mixture with NaCl (indefinitely stable). Buffer, pH 10. Dissolve 35.0 g. of NHCI in 200 ml. aqueous NH, (taken directly from supplier's battle); dilute to 500 ml. with H20. Buffer, pH 6. Dissolve 21.8 g. of sodium acetate trihydrate and 5.2 ml. of glacial acetic acid in H20; dilute to 250 ml. with HZO. Buffer,pH 5. Dissolve 2.7 g. of sodium acetate trihydrate in 19.2 ml. of 1 M HC1; dilute to 100 ml. with H20. Metal salt solutions. Prepare 100 ml. of 0.1 M aqueous salutions of the following salts: BaCL, Cr(NO,),, Ni(NO&, K2CreOi, Zn(NO&, and 200 ml. of the following: Mg(NOa)., FeCl.. Bismuth nitrate, 0.1 M . Dissolve 4.9 g. of Bi(NOa)r5H?0in 3 ml. of coned. HNOs and 10 ml. of H20. When solution is complete, dilute to 100 ml. with H20. (Do not employ heat to effect &lution.) Other s~lutionsinclude methyl orange (0.1 g./lW ml. of H20), 100 ml. 15% KI, 10 ml. 10% (NHJaSO,, 100 ml. 1 M NaOH, 100 ml. 1 M HC1,3% H202. Solid reagenk include KCN, tartaric acid, ascorbic acid (reagent ar U.S.P.). Standard Solutions for Ezpt. I X . Use only deionized or redistilled water in the preparation and comparison of the following standard solutions. EDTA, 0.0100 M . Dry about 4 g. of reagent grade disodium ethylenediaminetetraitcetate at 80°C. for about 2 hours. Weigh accuratelv 3.722 e. of this dried material, dissolve in and dilute d o n e h e r in a volumetric flabk. Transfer the saluwith ~ . to tion at once to a polyethylene bottle for storage. Mapesium nit~afe standardized soh. Dilute 50 ml. of the 0.1 M Mg(NOa)% to 500 ml. with H.O. Compare this solution with the 0.0100 M EDTA as follows: Pipet 25 ml. of this Mg(NO& solution into a 100-ml. beaker, add 10 ml. of pH 10 buffer, and dilute to 50 ml. with H?O. Add a few crystals of KCN (to avoid any possible indicetor blocking by trace heavy metals), a few drops of Erio T solution, and titrete with the 0.0100 M EDTA until the last tint of red has just disappeared. Preferably warm the solution slightly during the titration because the rate of reaction between the megnesium-indicator complex and EDTA is somewhat slow st room temperature. (One ml. of 0.0100 M EDTA = 0.2432 mg, of magnesium.) EXPERIMENTS
with various metals that differ markedly in color from that of the free, "unmetallized" indicator. Metal indicators in general also show more or less pronounced color changes with changes in pH. The present experiment demonstrates the metal indicator activity of Erio T as well as its pH-sensitivity. This dye shows a wine red color in strongly acidic solution, which passes to a dark blue color in the pH interval 6 to 7, and this to a red a t 11 to 12. The metal complexes formed by the dye in the pH range of about 6.5 to 11, are usually red. In each of five vessels place 90 ml. of stock H,O and 3 drops of Erio T solution and stir. Use these solutions in the following exoeriments. Ezpl. IA. To the first vessel add 10 ml. of pH 10 buffer. Stir. Note the blue solution oolor. Ezpt. IB. To the second vessel add 10 ml. of 1M HCI. Stir. Note the red color. Ezpt. ZC. To the third vessel add 10 ml. of 1 M NaOH. Stir. Note the red color and contrast the shade developed in Expt. IB. (If desired, a single solution of the dye may be carried through the chanees in color: however. the ~roeressivedilution creates a
Before performing the following experiments, explain that the use of a buffered medium assures that no change in hydrogen ion concentration occurs, and that therefore color changes observed must be due to a reaction of the metal ion with the dye (metal indicator). The sensitivity of the reaction of metal ions with Erio T is amazingly high. Even or 10-%M Mg(NO8)? solutions give a faint red tint (see also Expt. IIID). Ezpl. ID. To the fourth vessel, add 10 ml. of pH 10 buffer. Stir. Note the color is the same as that in Expt. IA. Now add one drop of 0.1 M Mg(NO&. Note the immediate color change from blue to red. Ezpt. IE. Repeat Expt. I D substituting one drop of 0.1 M Zn(NO.)l. Note the color change from blue to a red differing somewhat in shade from that obtained in Expt. ID.
Before performing the following experiment, it may be explained that a metal and hydrogen ion (both cations) compete for the metal indicator. Hence, a pH range must be selected in which the metal can compete successfully if the dye is to serve for the detection of a metal. Ezpt. I F . Add one drop of 0.1 M Zn(N0a)s to the solution from Expt. IB (i.e., the eecond vessel). Note that the solution color is unchanged. Experiment 11. Metal Complex Formation
EDTA reacts with many polyvalent metal ions to form water-aoluble metal complexes. EDTA is a tetrabasic acid. Two hydrogens have strongly acidic properties: the other two are only weakly acidic. The anion species predominating in about neutral solution may therefore be written as HzY-$. If complexation occurs, this anion is engaged completely and hydrogen ion is liberated. This can be expressed for a divalent metal ion, M+?,in the form of the following equation:
Experiment I. The Metal Indicator
Metal indicators are sensitive to changes in pM, that is, the negative logarithm of the concentration of the "free" metal ion. (The "free" metal ion, of course, may be in the form of such solution complexes as acetato, chloro, aquo, etc.) Metal indicators form complexes
I t may be explained that the further this reaction proceeds to the right, the greater is the stability of the metal complex, and the greater the acidity produced. Under certain circumstances this increase in acidity JOURNAL OF CHEMICAL EDUCATION
can serve as a basis of calculating the stability constant of the metal-EDTA complex. The phenomenon of complexation is demonstrated in the followinn exneriments in two wavs: The increase in acidity on the mixing of neutralized solutions of a metal salt and of EDTA is visualized by use of an acid-base indicator. Further, the formation of a colored metalEDTA complex is shown.
- .
Ezpt. IIA. Add 10 ml. of 0.1 M Zn(NO& and 3 drops of methyl orange solution to 90 ml. of stoek H20. Prepare a. second solution containing 10 ml. of 0.1 M EDTA, 3 drops of methyl orange solution, and 90 ml. of H?O. Nate that both solutions sre yellow (if not, add aqueous NH8 drapwise rts necessary). Now pour bath solutions simultaneously into a large beaker or a large cylinder. Note the color change from yellow to red. (If desired, in a psrsllel experiment the Zn(NO& and EDTA solutions can he mixed without addition of methyl orange: thus ascertaining that the zinc-EDTA complex is colorless.)
The color of the EDTA complex of some metals is different from that of the "free" metal ion. This can serve to demonstrat,e complexation. (Reference to the classical complexation of copper(I1) with ammonia may he expedient.) Ezpt. IIB. To each of two vessels, add 10 ml. of 0.1 M Ni(NO&, 20 ml. of pH 5 buffer and 60 ml. of stock H.O. Nate the green solution color of the "free" nickel ion. To one vessel add 10 ml. of 0.1 M EDTA. Note the color change from green to blue indicating formation of the nickel(I1)-EDTA complex. To the second vessel add 10 ml. of stock H 2 0 so that the final volumes are identical in both vessels. Save both salutians for comparison purposes in Expt. IV. Ezpt. IIC. To a third vessel add 10 ml. of 0.1 M EDTA, 20 ml. of pH 5 buffer, 60 ml. of H.0, and then 10 ml. of 0.1 M NifNO.1,. Note the formation of the hlue color of the nickel(II1E d ~ ~ - & n p l e thus x , indioating that the order of the a d d k i n of reagents has no effect. Complexation is also demonstrated hy formation of a colored cbromium(I1I)-EDTA complex in Expt. V. If an additional experiment is desired, the pink cohalt(I1)-EDTA complex may be developed in stock HzO buffered to p H 10.
Experiment 111. Complexation and the Metal Indicator
The formation of a metal complex by the metal indicator Erio T as seen in Experiment I is an example of compl?xation. The following experiments demonstrate how EDTA, by forming a stronger complex, can abstract a metal from its complex with thie indicator. It may be explained that. this is the basis of a visual detection of the end point in the EDTA titration. This titration was introduced by G. Schwarzenbach (15-I?), and Expt. I I I D was originally suggested by him (17). Ezpt. IIIA. Add 5 ml. of 0.1 M Mg(NO& to 90 ml. of stock H.0 and then 10 ml. of p H 10 buffer. Now add 10 ml. of 0.1 M EDTA. Note the solution remains colorless; hence, the magnesium-EDTA complex is colorless. Ezpt. IIIB. Add 3 drops of Erio T solution to 90 ml. of stock water and then 10 ml. of pH buffer. Note the blue eolor of the "free" dye (as already seen in Expt. IA). Now add 10 ml. of 0.1 M EDTA. Note the eolor is not changed; hence, there is no "complication" due to a color reaction of the dye with EDTA. Ezpt. IIIC. Add 3 drops of Erio T solution to 90 ml. of stack H20and then 10 ml. of pH 10 buffer. Now add 5 ml. of 0.1 M Mg(N0d1. Note the red color of the metal-dye complex (its already seen in Expt. ID). Now add an amount of 0.1 M EDTA greater than 5 ml. Note that the color reverts to that of the "free" (unmetallized) dye. (If desired, the mobile character of the system can he demonstrated by alternate additions of the metal salt and EDTA solutions.)
The release of metal ions (mainly calcium) from soft glass is an impressive demonstration of the sensitivity VOLUME 35, NO. 12, DECEMBER, 1958
of a metal indicator and also emphasizes the principles illustrated in Expt. IIIC. The following experiment also delineates that soft glass vessels should be avoided in analytical work with EDTA and metal indicators, especially on the micro scale and with dilute solutions. Ezpt. IUD. Take a soft glass container of capacity greater than 100 ml., preferably s, new bottle not yet placed in use. Rinse thoroughly with redistilled or deionized H*O( not with the prepared stock 8 0 containing EDTA). Place 80 ml. of redistilled or deionized H 9 0 in the bottle, and add 10 ml. of buffer pH 10 and 2 drops of Erio T solution. The solution should develop a hlue color. If not, add a. very dilute EDTA solution (obtained by 1:5 dilution of 0.1 M EDTA) dropwise until the blue is just reached. Allow to stand for a. short time. Observe the reappearance of the red color due to release of calcium ions from the glass. Add a few more drops of diluted EDTA solution and repeat the process. Take a soft glass rod and scratch the inside wall of the container. Nate the rapid reappearance of the red eolor because the scratched surface delivers large amounts of ions.
Experiment IV.
Stability of Complexes
Experiment IIIC demonstrated the difference in the stability of the complex of a single cation, namely magnesium ion, mit,h two different complexing agents, namely a mntnl indicator and EDTA. The following experiment demonstrates the difference in the stability of the complex formed by a single complexing agent, namely EDTA, with two different metals. The experiment is based on the addition of nickel(I1) to a bismuthEDTA solution; the color of the nickel(I1)-EDTA complex does not appear indicating the greater stability of the bismuth-EDTA complex. Ezpt. IVA. Add 10 ml. of 0.1 M EDTA to 10 ml. of 0.1 M Bi(NO&. Dilute with 50 ml. of H 2 0 and add 20 ml. of pH 5 buffer. Note that the solution is colorless and hence that the hismubh-EDTA eomnlex is colorless. Now add 10 ml. of 0.1 M
experiment use rertsanable care in measuring the volumes of bitmuth nitrate and EDTA solutions to assure that the bismuth ion is not in excess; otherwise a turbidity may appear.)
The following experiments demonstrate how the relative stability of the complexes formed by nickel with Erio T, EDTA and cyanide can he ranked on the basis of simple color reactions. Attention may he called to the fact that stability constants relate to equilibrium conditions and that false impressions may be reached if reaction rates are not appreciated (see also Expt. V). If desired, it may be explained that Expt. IVB is an example of so-called indicator "blocking" and that the nickel-Erio T complex dissociates only very slowly. Ezpt. IVB. Add 5 ml. of pH 10 buffer and 3 drops of Erio T solution to 95 ml. of stock H 2 0 Add 2 drops of 0.1 M Ni(NO&. Note the red color indicating the formation of the nickd(I1)-dye complex. Add 0.1 M EDTA dropwise until an appreciable excess (more than 15 drops) is present. Nate that the color of the solution is unchanged. Allow to stand for 15 to 30 minutes, note the blue color of the free dye is slowly attained. (A similar solution, containing only 4 drops of 0.1 M EDTA will undergo the color change in 3 to 4 minutes when warmed to about 70°C., due to tho increase of reaction rate with temperature.) Ezpt. IVC. Add 5 ml. of pH 10 buffer and 3 drops of 0.1 M EDTA to 95 ml. of stack H20. Add 2 drops of 0.1 M Ni(N0s)s. Now add 3 drops of Erio T solution. Note the blue color of the "free" dye. Ezpt. IVD. Add 5 ml. of pH 10 buffer and 3 drops of Eriochrome Black T indicator solution to 95 ml. of stock H.0. Add
Expt. VIIIB. Repeat Expt. VA or employ the violet solution saved from that experiment. To that solution, add 20 drops of 1M NaOH. Note the color change from violet to blue. Experiment IX. The EDTA Titration
An E D T A titration t o a visual end point is demonstrated in principle in Experiment 111. An applied example t h a t may be performed either as a lecture demonstration or as a student exercise is t h e titration of magnesium in t h e presence of large amounts of iron. T h e latter is masked by conversion t o t h e ferrocyanide ion. This example is particularly instructive a s most of the steps involve color changes. I n t h e masking of iron(III), cyanide must be added in alkaline solution so t h a t formation of hydrocyanic acid is avoided; however, iron(II1) hydroxide would then precipitate. This difficulty is avoided by first adding tartrate ion t o form t h e soluble iron(II1)tartrate complex. This complex is then reacted with cyanide ion in alkaline medium t o form ferricyanide ion. However, this latter ion will oxidize Erio T and thereby destroy its indicator action. Ferricyanide is therefore reduced t o ferrocyanide b y the action of ascorbic acid. This is a slow reaction; hence, warming is necessary. (If applied materials are being analyzed, ferrocyanide ion m a y form insoluble salts with some cations; in t h a t case, dilute t h e solution sufficiently or employ a back-titration, t h a t is,add a n excess of E D T A and back-titrate with a metal salt solution such a s magnesium nitrate.) As t h e rate of reaction of magnesium-indicator complex and E D T A is somewhat slow a t room temperature, i t is preferable t o titrate the solution while i t is warm. Calcium, manganese(II), lead, or zinc can be titrated in exactly t h e same way (and, if present, vill be co-titrated with magnesium). For a quantitative experiment, employ the standard EDTA and magnesium nitrate solutions given under "Reagents." For a semiquantitative demonstration, merely dilute t h e 0.1 M E D T A and Mg(NO& solutions tenfold. To a me~wredamount (10 to 30 m1.l of standard Me(N0.I.. ~ add some 0.1 M F~cI, (2 to 4 ml.) &d dilute to a v&m"".,.
no 707-19 11457) A-
(8) FLASCHM,H., A. J. BARNARD, JR., AND W. C. BROAD, "The EDTA Titration: Applications," Chemisl-Analyst, 46. 10612 (1957): . . . ibid... 47.. 22-28. 52-56. 78-84. in press (1958). (9) PRIRIL,R., "Recent Advances in Chelatometry," Analyst, 83, 188-95 (1958). (10) SCHWARZENBACH, G., "Chelate Formation as a Basis for Titration Processes," Anal. Chim. Acta, 7, 141-55
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(11) SCHWARZENBACH, G., 'The Complexones and Their Analytical Application," Analyst, 80, 713-29 (1955). (12) J. T. BAKERCEEMICALCO., "The EDTA Titration: Nature and Methods of End Point Detection." Phiui~sburg, New Jersey, 1957,32 pp. (13) THE Dow CEIEMICAL Co., "Keys to Chelation," Midland, Michiian, 1957.16 pp. CEIEMIC~L Co., "Sequestrene," Providence, Rhode (14) ALROSE Island, 1952, 54 pp. [available from Geigy Industrial Chemicals, P. 0. Box 430, Yonkers, New York].
(15) SCKW~RZENBACE, G., "Nouvelles mdthodes de dosage de cehins cations mPtaIliques," Helv. Chim. Acfa, 29, 1338 (1946); U. S. patents 2,583,890 and 2,583,891 (1952). (16) SCHWARZENBACH, G., W. BIEDERMANN, AND F. BANQERTER, "Neue einfache Titriermethoden mr Beatimmung der Wasserhiirte," Heb'. Chim. Acta, 29, 811418 (1946).
(17) BIEDERMANN, W., AND G. S ~ A R Z E N B A C"Die H , komplexometrisohe Titration der Erdalkelien und eiuiger anderer Metalle mit Eriochromschwarz T," Chimia, 2 , 5S59 (1948). (18) MARTELL, A. E., "The Behavior of Metal Complexes in Aqueous Solutions,"J. CHEM.EDUC.,29, 270-80 (1952).
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