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edta y complexometria

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  • EDTA AND COMPLEX FORMATION

    BY THEIR nature general textbooks often lag by many years in their inclusion of salient advances in analytical reagents. Teachers are aware of this lag when con- fronted 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, mask- ing 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 water- soluble chelates, and especially to EDTA and its use in analytical chemistry. These experiments also deline- ate 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 IX 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

    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

    in most instructional laboratories. The two special reagents, disodium ethylenediaminetetraacetate dihy- drate and Eriochrome Black T (Erie T for short), are readily available through laboratory supply houses.

    The intention is not merely to give 'Lcook-book" recipes but also to provide some background informa- tion by discussion of the experimental phenomena. This information can only he presented briefly and in summary form. Hence, a selected bibliography is ap- pended that cites some relevant monographs (1-4), re- views (5-11), trade literature (13-id), and a few papers of historic interest (10,15-17).

    EDTA is the alphabetic designation for ethylene- diaminetetraacetic acid or its anion form ethylene- diaminetetraacetate.

    HOOC-H,C CHrCOOR )N.cH~cH,N<

    HOOGHC CHrCOOH Ethylenediaminetetraacetic acid

    NaOOC.H& CHrCOONa [ HWCH1y< 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 hond- ing with an equal number of coordination positions of a metal ion. In 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 solu- bility on such metal complexes. The steric configura- tion (18) of a metal-EDTA complex is shown in the figure. A fuller treatment of the theoretical back- ground 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 direc- tions 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 titra- tion. Rather than a study of the simple titration of a

  • single metal ion in pure solution, a slightly more com- plicated 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 Ex- periment 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 salu- tions 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?0 in 3 ml. of coned. HNOs and 10 ml. of H20. When solution is com- plete, dilute to 100 ml. with H20. (Do not employ heat to effect &lution.)

    Other s~lutions include 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 (re-

    agent ar U.S.P.). Standard Solutions for Ezpt. I X . Use only deionized or redis-

    tilled 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 80C. for about 2 hours. Weigh accuratelv 3.722 e. of this dried material, dissolve in and dilute with ~ . d to o n e h e r in a volumetric flabk. Transfer the salu- tion at once to a polyethylene bottle for storage.

    Mapesium ni t~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 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

    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 ex- periment 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 1 M 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 ~roeressive dilution creates a

    Before performing the following experiments, ex- plain 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 re- action 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 suc