QUANTITATIVE DETERMINATION OF METALS IONS USING Fe (III)/ Fe (II) REDOX TITRATION SYSTEM WITH A PLATINUM ELECTRODE

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1 Journal of the University of M. Chemical M. El Jamal, Technology H. H. Hammud and Metallurgy, 42, 1, 27, QUANTITATIVE DETERMINATION OF METALS IONS USING Fe (III/ Fe (II REDOX TITRATION SYSTEM WITH A PLATINUM ELECTRODE M. M. El Jamal 1, H. H. Hammud 2 1 Faculty of Science, Chemistry Department, Beirut Arab University, Beirut Lebanon, Box: mjamal@ul.edu.lb 2 Faculty of Science, Chemistry Department, Received 18 January 27 Accepted 2 February 27 Lebanese University, El Hadath, Lebanon ABSTRACT The determination of metal ions using Pt electrode to detect the equivalence point was performed by classical potentiometry. Cu (II, Cd (II, Zn (II, Al (III, Cr (III, Pb (II, Ni (II, and Co (II having their formation constant between that of FeY 2- and FeY - can be determined quantitatively by this method. The behavior of some metal ions (such as Ni (II, Cr (III and Al (I is affected by kinetic problems where Fe (III does not uptakes EDTA from their complexes with EDTA. Other metal ions like Zn (II, Co (II, and Cd (II partially deliver their EDTA to Fe (III. However, the metal ions likes Ba (II, Ca (II and Mn (II having their formation constant less than that of FeY 2- can not be studied by this method since Fe (III uptakes the EDTA from the metal complexes before up taking from FeY 2-. Keywords: potentiometry, Pt electrode, EDTA titration, Fe (III/ Fe (II system. INTRODUCTION Complexometry titration of metal ions by direct or back titration with EDTA has been well studied [1, 2]. The determination of concentration of metal ions can be evaluated by ph-metry titration with Na 2 EDTA [] and also by UV/Vis absorption spectrometric titration using special ligands [, 4]. Twenty years ago, intensive research was done to elaborate selective electrodes to measure the concentration of metal ions [5]. However, these selective electrodes need a special treatment to be active and any defect affects their response. Anodic stripping voltammetry has been widely used for the determination of lead, zinc and cadmium. Amperometry titration using dropping Hg electrode was also used to determine metal ions [6, 7]. We propose the determination of metal ions using a Pt electrode. This electrode is easy to use and do not need special treatment before use. In the literature, we only found the determination of Ni (II by the suggested method [8], we try to apply this method to other metal ions, like Co (II, Cu (II, Cd (II, Zn (II, Pb (II, Al (III, Cr (III. The iron system has the advantage that both Fe (III and Fe (II are soluble in water and forms stable complexes with EDTA, with different constant formation, in contrast to Cerium system where only Ce (III forms complex with EDTA [1, 2]. The Pt electrode is good sensor to follow the concentrations of Fe (III and Fe (II. EDTA is the best ligand for this study because EDTA forms 1:1 stable complex with the major metal ions especially with Fe (III and Fe (II ions, in contrast 97

2 Journal of the University of Chemical Technology and Metallurgy, 42, 1, 27 to o-phenantroline (o-phen or other ligand which forms MLn complexes (n = 2 or with several metal ions [9]. The formation constant of M (o-phen n to be studied are smaller than Fe III (o-phen and Fe II (O-phen [9]. Thus these metals can not be detected by this method. EXPERIMENTAL The potentiometric measurements were carried out using computerized potentiometer with Nico 2 Elit 8-channels ions analyser (reading ±.1 ph or mv unit. The combined glass electrode was calibrated from time to time using two standard buffer solutions 4. and 7.. The response of the Pt combined electrode and its slope was verified with Fe (III/ Fe (II in 2. M H 2 SO 4. All the reagents used were purchased from BDH products. All solutions were prepared in deionized water. Ferrous solution was prepared from Mohr salt in.1 M H 2 SO 4 to prevent its oxidation. Ferric solution was prepared from hydrated ferric sulfate in.1 M H 2 SO 4 to prevent its precipitation. The Na 2 EDTA solution was prepared as described in [1]. Metallic ions solutions were standardized by complexometry with.1 M Na 2 EDTA solution [1, 2]. Potentiometric procedure In the first step 1. ml of EDTA (.1 M was added to 1. ml of Fe (II.1 M and 1. ml of acetic buffer 2 M ( ph = 4.2 into a beaker of 1 ml and diluted to 4. ml with deionized water, then titrated with Fe (III (.1 M. The acetic buffer is used to fix the ionic force and the ph of the solution. In the second step, we add to the initial mixture a variable volume (between. and 1. ml of the metal ion (.5 M to be titrated so that the total volume of the solution is 4. ml. The concentration of the metal ion, Fe (II and M n in the beaker must be less than the initial concentration of EDTA. The ph and the potential were taken after each addition of Fe (III solution (.1 M. The ph measurement is not necessary; it is only to be sure that the ph remained constant during titration. Only in the case of Cr (III, the titration was done at 45 o C to accelerate the complex formation. The titration of Cr (III at room temperature is not accurate because of non spontaneous formation of complex, since the violet colour of CrY - develops with time. THEORY Variation of the potential of the solution during titration EDTA forms complexes with Fe (III and Fe (II. To simplify the potential relations, we do not take into account the formation of complexes between Fe (III, Fe (II and acetate (Fe (II-acetate: logâ 2 = 5., Fe (III- acetate: logâ 2 = 1. [9]. Before adding Fe (III there is in solution the complex FeY 2- and excess of EDTA: Fe( II H = Y 2CHCOO FeY 2CHCOOH K 1 The Pt electrode measures a mixed potential of the FeY 2- (reductant and CH COOH (oxidant in CH COOH/ H 2. After adding Fe (III the reaction will be: Fe( III H Y CH COO FeY CH COOH K = And the potential E (V of the solution will be as predicted by Nernst equation at 25 o C: 2 E = E ( FeY / FeY.59 log[ FeY ] /[ FeY 2 ( ( / (.59log /.59log[ ]/[ ] E = E Fe III Fe II 2 β1 β2 FeY FeY =.1.59 log[ FeY /[ FeY β1 and β 2 : Formation constant of FeY 2- and FeY - and equal to and , respectively [9]. β1 and β 2 are both multiplied by the same factor 4- (á Y, their ratio remains the same. E ( Fe ( III / Fe ( II =. 77 E = E The last equation can be written as follow: ( FeY / FeY 2 V 2.59logV (Linear relation where V is the volume of Fe (III added. When the free EDTA are complexed, Fe (III added will react now with FeY 2- to form the more stable complex Fe Y - : ] ] 98

3 M. M. El Jamal, H. H. Hammud 2 FeY Fe( III FeY Fe( II K = β / β = ( ( 2 FeY Fe III FeY Fe II 2 2 ( if MY is more stable than FeY (2 During the reaction [FeY 2- ] decreases and [FeY - ] continues to increase. But as the.59 log [FeY - ]/ [FeY 2- ] slightly varies, we will not observe an appreciable change in potential. At the equivalent point the ox-red system in solution will be Fe (III/ Fe (II and a sudden jump in potential will occur. The relation between potential E (V and the added volume of Fe (III after the equivalent point will be: E = E ( Fe( E = E.59log[ Fe( I/[ Fe( ( Fe( III / Fe( II.59 log( V V1 V 1 is the volume of Fe (III to complex the total EDTA. The titration curve of Fe (III with EDTA shows only one equivalent point (Fig E ( mv ml Ni (II theoritical curve of1 ml of Ni (II Fe (II alone V2 volume of Fe (III Experimental curve Theoritical curve Fig. 1. Comparison between the theoretical and experimental titration curves of 1 ml of Ni (II,.5 M with Fe (III in presence of 1ml of Fe (II,.1 M and 1 ml of EDTA,.1 M. V1 2 MY Fe( III FeY M ( II with K = β / β ( β : the formation constant of MY 1 ( Theoretically, during the reactions (1 and (2, the potential of the solution will be the same as in absence of M (II and the potentiometric curve will be overlaid. At the end of the reaction (2, the potential increases suddenly, the first equivalent point corresponds to the complexation of free EDTA with Fe(III and destruction of FeY 2- (call V 2 this volume, then additional Fe (III attacks the more stable complex MY 2-. This reaction ( gives a second plateau and prevents the potential to reach the higher value of Fe (III/Fe (II system. At the end of reaction ( a second equivalent point is obtained corresponding to the total destruction of MY 2-. The theoretical titration curve of M (II such as Ni (II thus should show two equivalent points as seen in (Fig. 1. The concentration of Fe (III during the reaction ( will be: [ FeY ][ Ni( [ Fe( I = [ NiY β [ FeY ][ Ni( = 1 2 ] β1 [ Ni( As the Pt electrode is sensitive to the system present in solution the potential will be: E = E ( Fe(.59log[ Fe( I/[ Fe( Variation of the potential of the solution during titration of M (II in presence of excess of EDTA and FeY 2- Before adding Fe (III, the composition of the solution is MY 2-, FeY 2- and excess of EDTA, after adding Fe (III the successive chemical reactions will be in order: EDTA Fe ( III FeY (1 E = E Fe(.59 log( β / β ( 1 [ FeY ][ Ni(.59 log [ Ni( [ Fe( From the titration curve we can calculate the concentration of the metal ion and the constant formation of MY 2-. The volume of EDTA reacted with x ml of Ni (II is equal to the difference in Fe (III volumes (V 1 -V 2 99

4 Journal of the University of Chemical Technology and Metallurgy, 42, 1, 27 and the concentration of the unknown solution will be: x. [Ni ( =.1 (V 1 V 2. When MY 2- is totally destroyed, free Fe (III will be in solution, the potential will be another time equal to: E = E ( Fe(.59log[ Fe( I/[ Fe( shape of the curve, the potential measured, the height of the jump, and the volume of the equivalent point. We conclude that the best ph for titration of Fe (III with EDTA is between. and E = E ( Fe(.59log( V V ph =1 ph =.5 ph= 4. ph=1 But some time for kinetic problem like the case of Al (III, Ni (II and Cr (III, we see only one equivalent point corresponding to the titration of the remaining EDTA in solution after adding M (II. RESULTS AND DISCUSSION E (mv ph=5.2 ph=.5 25 ph 6.2 ph= ph=5.2 1 ph= Determination of the optimum conditions for the titration of the metal ion with Fe (III Effect of the initial volume of Fe (II added to solution The titration curves of Fe (III with EDTA in presence of different volume of Fe (II (.5, 1, and 5 ml show that the volume of Fe (II added affects the E measured, but does not affect the shape of the curve, neither the volume of the equivalent point, but the sudden change in E is obtained at low volume (.5 and 1 ml. We prefer to add 1 ml of Fe (II for convenience. Stabilities of the complexes as a function of the ph The conditional constant formation of FeY - and FeY 2- varies with ph. The complex FeY 2- is not stable at ph lower than. due to the weakness of its conditional constant value, but the constant formation of FeY - is very big. The reaction (1 occurred partially in very acidic medium and it is confirmed by the appearance of the yellow colour of FeY - after adding of a small quantity of Fe (III. On the other hand in the presence of EDTA, Fe (III hydrolysis (formation of Fe(OH 2 and Fe (OH 2 occurred at ph higher than 6., so we must choose a compromise about the selection of the ph of the solution. The curve E in function of ph of the system FeY - / FeY 2- shows a plateau between and 7 so we conclude that these complexes are stable in this zone. Effect of the ph of the buffer The titration curves of Fe (III with EDTA in presence of Fe (II at different ph between 1. and 6. (Fig. 2 show that the ph of the solution affects the Volume of Fe (III Fig. 2. Effect of the ph of the solution on the titration curve of Fe (III with EDTA. Also, the conditional constant formation of MY n-4 varies with ph, so we must choose the best ph for the titration of Fe (III and M n with EDTA. We select Ni (II for this study. The titration curves of Ni (II with Fe (III in presence of Fe (II and excess of EDTA at different ph between 1 and 5 show that the complexes NiY 2- and FeY - are stable at lower ph, but the titration at ph 1. is not accurate. The titration curves at ph higher than show the typical S- shape. Since the complex stability increases with ph, we have chosen the ph 4.2 as optimal ph for the titration of all metal ions. Apparent standard potential of Fe (III/ Fe (II in acetic acid medium 2 M In order to interpret the titration curve of the metal ion, we calculated experimentally the apparent standard potential in acetic buffer for Fe (III/Fe (II system and FeY - /FeY 2- systems. The titration curve of a mixture of 5. ml of each of Fe (III and Fe (II (.1 M with EDTA shows two plateaus. When the volume of EDTA added is less than 5. ml, EDTA complexes only with Fe (III ion. Fe (II EDTA complex begins to form for addition of more than 5 ml EDTA. The measured potential in the first plateau corresponds to the system Fe (III/Fe (II in acetic buffer, whereas the second one corresponds to the system FeY - / FeY 2-. We obtained.5 V and.14 vs. S.H.E, respectively. The second value is near the published value [9]. 1

5 M. M. El Jamal, H. H. Hammud Table 1. Volume of.1 M Fe (III, at the first equivalent point as a function of the volume of Ni (II added.5 M (2, 5, 7, 1 ml. Volume of Ni (II to be titrated Volume of Fe (III at the 1st eq.pt V 1 - V [ Ni ( Titration of metal ions with Fe (III in presence of Fe (II and excess EDTA The metal ions studied were divided into two groups with respect to the behaviour of their complexes with free Fe (III. The first group contained four metal ions (Ni (II, Cu (II, Al (III and Cr (III and the second contained four metal ions (Zn (II, Co (II, Cd (II and Pb (. Titration of Ni (II, Cu (II, Al (III and Cr (III The potentiometric titrations of Ni (II show a shift of the equivalent point to the left with respect to the titration curve in absence of Ni (II (Fig.. The shift is proportional to the volume of Ni (II added. The volume of the first equivalent point is a function of the volume of Ni (II added as shown in Table 1. The obtained titration curves for Al (III, Ni (II, and Cr (III are very similar. The titration curve of Cu (II is also similar, but deviation occurred at higher potential : 1 mv vs AgCl /Ag electrode. After total complexation of Cu (II, the copper curve deviates slightly down with respect to the nickel curve (Fig. 4. This difference in behaviour between Ni (II and Cu (II can be explained with a partial release of EDTA from CuY 2- to Fe (III. This deviation is important at higher concentration of Cu (II (greater than 5. ml copper as predicted by Le Chatelier law. The experimental titration curves of Ni (II are in contradiction with the theoretical curve, which must present two equivalent points, one corresponding to the titration of free EDTA and destruction of FeY 2- and the second - to the destruction of MY 2-. This can be explained with a kinetic problem, since the reaction ( is thermodynamically possible (the equilibrium constant >1 6, but kinetically unfavorable (the activation energy of the slow step is very high. Spectroscopy study confirms that the absorbance at E (m V E (m V nm (ë max of NiY 2- remains constant during the addition of Fe (III. For excess Fe (III, ë max shift to 56 nm due to the formation of Fe (OH n -n species (n = 1 or 2. Several attempts were done to over come this problem, 5 ml Ni (II 1 Ni (II 2 ml Ni (II 7 ml Ni Fe (II alone Volume of Fe (III Fig.. Titration curves of x ml of Ni (II.5 M with Fe (III,.1 M, in presence of Fe (II and excess of EDTA. 5 ml Ni (II delat E/5 ml Cu 1 Ni (II delat E/1"ml Cu V olume of Fe (III Fig. 4. Comparison between the titration curves of Ni (II and Cu (II with Fe (III at similar conditions. 11

6 Journal of the University of Chemical Technology and Metallurgy, 42, 1, 27 E (mv to accelerate the exchange between MY 2- and Fe (III. Three factors were studied on Ni (II in order to displace the equilibrium to the right: a The titration at higher temperature (5 and 45 o C; b The titration at different ph; c The titration in presence of dimethylglyoxime (DMG. No change in the measured potential of the titration curve was obtained. The addition of DMG does not cause the precipitation of Ni (DMG 2. We observe only at ph 1. a partial release of EDTA from NiY 2- to Fe (III, but at ph 1. the titration of Ni (II is not accurate as mentioned above. Titration curves of Zn (II, Co (II, Cd (II and Pb (II The titration curves of Zn (II with Fe (III in presence of Fe (II in acetic buffer, show two equivalent points, the first equivalent point corresponds to the titration of free EDTA and destruction of FeY 2- and the second equivalent point corresponds to a partial release of EDTA from ZnY 2- to Fe (III (reaction ( (Fig.5. We use the derivative method to find the volumes of the equivalent points in contrast to the titration of Ni (II where we use the tangential method. The shift to the left in the first equivalent point is proportional to the volume of Zn (II added as in the case of Ni (II. The release of EDTA is more important than in the case of Cu (II. No correlation exists between the stability of the complexes and the exchange of EDTA between MY n-4 and Fe (III. The exchange of EDTA between Fe (III and Zn (II is slow. Since we are interested only in the determination of the concentration of Zn (II ml Zn (II1 m l Zn (II 5 ml Zn (II 2 ml volume of F e (III Fig. 5. Titration curves of x ml of Zn (II,.5 M with Fe (III.1 M in presence of Fe (II and excess of EDTA. E (mv ml of Co (II at 5 C 1 ml of Co (II at 64 C 1 ml of Co (II at 29 C Fe (II alone at 29 C Volume o f Fe (III Fig. 6. Titration curves 1 ml of Co (II at different temperature (65, 5 and 29 º C. the metal ion, we can stop the titration a little after the first equivalent point to reduce the time of the experiment. The second equivalent point is not involved in calculation. The potentiometric titration of 1 ml of.5 M Co (II with.1 M Fe (III in presence of excess EDTA at different temperature (29, 5, and 65 o C shows that these curves are nearly overlaid but deviates to the right near the second equivalent point (Fig. 6. The deviation increases with the increase in temperature. This deviation can be due to the increase in the exchange of EDTA between CoY 2- and Fe (III. At higher temperature (65 o C the exchange is not complete since the volume of the second equivalence point is experimentally 9. while theoretically is 9.8. The potential during reaction ( increases after each addition of Fe (III then decreases slowly, due to the slow formation of FeY -. The titration curves of Co (II and Cd (II are similar to those of Zn (II. The decrease in potential in the case of Zn (II is faster than that of Co (II which means that the exchange between Zn (II and Fe (III is faster than for Co (II and Fe (III. Spectroscopy study confirms the exchange of EDTA between CoY 2- and Fe (III, according to the reaction (. The absorbance at 5 nm (ë max of CoY 2- increases suddenly after adding Fe (III to CoY 2- solution due to the formation of Fe(OH n -n (n = 1 or 2. The absorbance then decreases with time due to the change of Fe(OH n -n to FeY -. The titration curves of a mixture of 5. ml of Ni(II.5 M with 5. ml of Al (III.5 M, or a 29 C 64C 12

7 M. M. El Jamal, H. H. Hammud Fig. 7. Titration curves of x ml of Pb (II,.5 M with Fe (III.1 M in presence of Fe (II and excess of EDTA. mixture of 5 ml of Ni (II.5 M and 5 ml of Zn (II.5 M indicate that the volume of the first equivalent point is similar to the volume of the equivalent point of 1. ml of Ni (II.5 M. So, we can only calculate the total concentration of metal ions. Thus the method is not selective. The titration curves of Pb (II show two equivalent points. The first corresponds to the total complexation of Pb (II with EDTA while the second one corresponds to the nearly total exchange of EDTA between Pb (II and Fe (III (Fig. 7. The presence of sulfate in solution coming from Fe (III and Fe (II solutions causes the precipitation of PbSO 4. In this way it accelerates the exchange between Pb (II and Fe (III: 2 2 ( III SO4 PbY FeY PbSO4 Fe The titrations of metal ions likes Hg (II and Ce (IV are not possible with this method, because the standard potential of these ions is higher than the standard potential of Fe (I/ Fe (II. This causes oxidation of Fe (II and further none functioning of the indicator electrode. Also, the formation of gray mercury colloidal precipitate glue on the hole of the combined Pt electrode leads to incorrect measured potential. CONCLUSIONS The titration curves E versus volume Fe (I show that the metal ions studied can be divided into three categories (Fig. 8: The metal ions having constant formation MY n-4 less than the constant formation of FeY 2- (1 14 like Mn (II, Ca (II, Ba (II; The metal ions having constant formation MY n-4 between that of FeY 2- and FeY - (1 14 < â < 1 25 like Co (II, Cu (II, Al (III, Cr(III, Cd (II, Zn(II, Pb (II and Ni (II; The metal ions which oxidize Fe (II in the present conditions, like Hg (II. The first and the third categories can not be studied by this method. In the first case Fe (III added will attack MY n-4 before FeY 2-. The reaction between MY n-4 and Fe (III does not change the concentration of FeY 2-, so there is no big change in E. The curves without M n and in presence of M n are similar. The presence of M n does not affect the potential of the solution. The metal ions in the third category react with Fe (II to give Fe (III. So the indicator electrode is eliminated from the solution. The addition of Fe (III will increase the concentration of FeY -, causing a slow increase in the measured potential. The second category includes eight metal ions Ni (II, Zn (II, Co (II, Cd (II, Cu (II, Al (III, Cr(III, Pb (II easily detected by the proposed method. The titration curves show that the behaviour of ions is different and can be divided into two groups. The first group includes Al (III, Cu (II, Ni (II and Cr (III where no exchange of EDTA occurs between MY n-4 and Fe (III. The second group includes Pb (II, Zn (II and Co (II where a partial release of EDTA occurs between delta E /5 ml Pb delta E /Fe (II alone delta E /1 m l Pb delta E /7 ml Pb delta E/1 Ni (II E (mv ml Co Fe (II alone 1 Ni (II 1ml Cu volume of Fe (III Fig. 8. Titration curves of 1 ml of three metal ions (Cu, Ni, Co with Fe (III, in presence of Fe (II and excess of EDTA. Cu Co 1

8 Journal of the University of Chemical Technology and Metallurgy, 42, 1, 27 MY n-4 and Fe (III. The detection limit of the first group is better than that of the second group since greater jump is obtained at the equivalent point. REFERENCES 1. A. Skoog, M. Donald, F. West, J. Holler, Chimie analytique, 7 éme édition, Bruxelles, De boek, D. C. Harris, Quantitative analysis, Fourth edition, F. B. Audat, F. Raffegeau, D. Prevoteau, Chimie Inorganique et générale, Dunod, Paris, Second édition, 2, G. D. Christian, Analytical chemistry, Fourth edition, A Abbaspour, B. Khajeh, Analytical sciences, Sep. 18, G. Charlot, Méthodes électrochimiques, v.iv, D. T. Sawyer, W R Heineman, J M Beebe, Chemistry experiments for instrumental methods, J.Wiley & Sons, H. A. Flaschka, E.D.T.A. titrations, Pergamon press, M. Roche, J. DesBarres, C. Cohin, A. Jardy, D. Bauer, Chimie de solutions (Langage et informatique, 199, p.179, p

Complex formation metal ion=lewis acid=electron-pair acceptor ligand = Lewis base= electron-pair donor

Complex formation titrations Complex formation Chelate effect EDTA Conditional formation constant EDTA titration Indicators Types of titrations Auxiliary complexing agents Complex formation metal ion=lewis