Complexation Key to a pH Locked Redox Reaction - ACS Publications

Laboratory Experiment pubs.acs.org/jchemeduc

Complexation Key to a pH Locked Redox Reaction Masood Ahmad Rizvi,*,† Yuvraj Dangat,‡ Tahir Shams,† and Khaliquz Zaman Khan† †

Department of Chemistry, University of Kashmir, Hazratbal, 190006 Jammu and Kashmir, India Physical Chemistry Division, CSIR-National Chemical Laboratory, Pune 411008, India

‡

S Supporting Information *

ABSTRACT: An unfavorable pH can block a feasible electron transfer for a pH dependent redox reaction. In this experiment, a series of potentiometric titrations demonstrate the sequential loss in feasibility of iron(II) dichromate redox reaction over a pH range of 0−4. The pH at which this reaction failed to occur was termed as a pH locked reaction. The comparative ability of 10 selected iron binding ligands with varied propensity for the redox potential modification of Fe(III)/Fe(II) redox couple to restore/unlock the pH locked redox reaction is shown using potentiometric titrations. The spectrophotometric speciation analysis of Fe(III) Tiron complexation with pH was carried out to explain the differing ability of EDTA and Tiron to unlock the reaction under different pH conditions. The experiment illustrates how environmental, biological redox reactions avoid severe laboratory conditions to occur and can be explored in the design of novel redox systems for natural attenuation of environmental toxins to their non- or lesser-toxic forms. The experiment also demonstrates prudent laboratory practice for safe waste disposal. KEYWORDS: Second-Year Undergraduate, Upper-Division Undergraduate, Physical Chemistry, Analytical Chemistry, Coordination Compounds, Aqueous Solution Chemistry, pH, Oxidation/Reduction, Thermodynamics, Potentiometry, Titration/Volumetric Analysis

T

The effect of pH on the reduction potential of the Cr(VI)− Cr(III) redox couple can be estimated using eq 1 (see Results and Discussion):

he redox potential (E) is the primary thermodynamic parameter guiding electron transfer reactions between chemical entities.1 However, redox reactions involving H+ or OH− ions are under the influence of an additional factor, the pH.2 The E−pH or Pourbaix diagram connects these interdependent quantities for a particular redox system.3,4 The pH dependent redox reaction works well under an optimum pH. A favorable pH can significantly enhance, while an unfavorable pH may block a feasible electron transfer reaction. Studying pH influence on a pH dependent redox system illustrates the pH−potential correlations build up from Nernst equation.2,5 Depending on the number of protons and electrons involved, pH dependent redox reactions undergo a defined potential modification per unit change of pH.2 The titration of a particular redox system under different pH conditions is comparable to the titration of a series of redox systems with different redox potentials.6 Potentiometric titrations7,8 can be pragmatic and pedagogic methods to visualize the pH effect on a pH dependent redox reaction. This experiment through a series of potentiometric titrations over a pH range (0−4) demonstrates the effect of pH on Fe(II)− Cr(VI) redox reaction. The Fe(II)−Cr(VI) redox reaction occurs favorably around pH 0.

© XXXX American Chemical Society and Division of Chemical Education, Inc.

E(pH) = E 0 ′ − 0.138pH

(1)

where E 0 ′ = E 0Cr(VI)/Cr(III) −

RT [Cr 3 +]2 ln nF [Cr2O7 2 −]

From eq 1, it is clear that the reduction potential of Cr(VI)− Cr(III) redox couple decreases by 138 mV for every unit increase in pH. Hence, at pH 4, the reduction potential of Cr(VI)−Cr(III) redox couple turns out to be 0.77 V. The cell potential of Fe(II)−Cr(VI) redox reaction at pH 4 can be estimated using eq 2:

The pH 4 at which this otherwise feasible redox reaction fails to occur has been termed as a pH locked reaction. Natural attenuation refers to utilization of benign natural processes for site-specific pollution treatment purpose, and Fe(II)−Cr(VI) redox reaction is a highly desirable environ-

A

DOI: 10.1021/acs.jchemed.5b00499 J. Chem. Educ. XXXX, XXX, XXX−XXX

Journal of Chemical Education

Laboratory Experiment

mental reaction9,10 for the natural attenuation of toxic Cr(VI) to comparatively less toxic Cr(III) as Fe(II) is a ubiquitous primary electron donor in various types of soils, aquatic systems, and subsurface environments.11 Cr(VI) is a hazardous chemical contaminant that must be removed in the water treatment process10,12 due to its carcinogenic and mutagenic effects.9 The amazingly lesser mobility, poor bioavailability, and reduced toxicity of Cr(III) in contrast to Cr(VI) make the Cr(VI)−Cr(III) reduction a vastly advantageous reaction from toxicity point of view. However, the necessity of a strong acidic medium (pH < 1) for Fe(II)−Cr(VI) redox reaction is a limiting factor in the natural water systems under anaerobic conditions. Reducing toxic Cr(VI) to relatively nontoxic Cr(III) in natural water systems around neutral pH is hence

a fascinating concept. An interesting observation about the Fe(II)−Cr(VI) redox reaction is that the carboxylates and phenolates are known to synergistically enhance the reduction of Cr(VI) by Fe(II) in natural water systems even at neutral pH.11 The plausible explanation for this interesting ability is the modulation of the Fe(III)−Fe(II) redox potential on complexation to these organic moieties.13−15 Fe2 + + x Ln − ↔ [Fe(L)x ](2 − xn) −

Where x is the number of ligand molecules attached. For example, complexation of iron with EDTA decreases reduction potential and with 1,10-phenanthroline raises the redox potential:16

[Fe(OH 2)6 ]3 + + e− → [Fe(OH 2)6 ]2 + (E 0 = 0.771 V) [Fe(EDTA)]− + e− → [Fe(EDTA)]2 − (E 0 = 0.08 V) [Fe(phen)3 ]3 + + e− → [Fe(phen)3 ]2 + (E 0 = 1.14 V)

The complexation modified redox potential of Fe(III)−Fe(II) redox couple can be utilized in the restoration of lost feasibility of Fe(II)−Cr(VI) redox reaction at pH 4. This ability of complexation has been referred to as unlocking of pH locked Fe(II)−Cr(VI) redox reaction and can be depicted as

previous work.16 Potentiometric measurements (E and pH) were carried under “nitrogen” environment on a thermostatic, magnetically stirred solution using Eutech PC5500 ion analyzer. Spectrophotometric measurements were carried on Shimadzu 1650 UV−visible spectrophotometer with thermostatic control. Procedures

The effect of hydrogen ion concentration (pH) on the potentiometric behavior of Fe(II)−Cr(VI) redox reaction was established by titrating 0.1 N Fe(II) with 0.1 N Cr(VI) in the presence of decreasing concentration (1.0, 0.5, 0.2, 0.1, 0.01, 0.001, 0.0001 M of H2SO4. In a typical titration set, 20.0 mL of H2SO4 (given concentration) was added to 12.5 mL of Fe(II) and titrated with Cr(VI). The cell potential was measured on each addition of Cr(VI) and plotted against the added volume of Cr(VI). The ligand effect on the Fe(II)−Cr(VI) redox reaction was established around pH 4.0 using 1 × 10−4 M H2SO4. The ligand-to-metal molar ratio was appropriately adjusted for an octahedral composition of iron (3:1 for bidentate and 1:1 for hexadentate). In a typical titration set, 12.0 mL of 0.05 M EDTA solution was added to 10.5 mL of 0.05 N Fe(II) solution and titrated against 0.05 N dichromate solution. On addition of ligand, an immediate fall in the cell potential was observed, which slowly began to increase on addition of Cr(VI). The cell potential was measured after each addition of Cr(VI) when the change in potential was less than ±2 mV. The pH of the titration mixture was measured at the start and the completion of the titration. For EDTA concentration profile, a sequential addition (10.0, 15.0, 20.0, 25.0, 30.0, and 35.0 mL) of 2.5 × 10−2 M EDTA solution was made to 15.0 mL of 5 × 10−2 N Fe(II) to which 20.0 mL of 1 × 10−4 M H2SO4 was added and titrated against 6 × 10−2 N dichromate solution. The Tiron−Fe(III) complex speciation analysis under two pH conditions (4 and 8) was determined through molar ratio method in which 1.0 mL of 1 × 10−3 M Fe(III) was taken in nine labeled 5.0 mL volumetric flasks. To each labeled flask, an increasing amount of 1 × 10−3 M tiron ligand solution in a sequential manner (0.0, 0.2, 0.4, 0.6, 0.8, 1.0, 1.2, 1.4, 1.6) was added, and their pH was adjusted around 3.5−4.0 using dilute H2SO4 solution over a pH meter. Their volume was raised to

In continuation of our work on coordination controlled redox behavior of transition metal ions,16−20 this experiment is a simple demonstration of pH effect on the redox potential of systems involving H+ ions in the redox reaction, pH locking, and unlocking of Fe(II)−Cr(VI) redox reaction by a selected group of ligands having varied propensity for the redox potential modification of Fe(III)−Fe(II) redox couple (see Pedagogy section).

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EXPERIMENTAL SECTION

Reagents

All solutions were prepared in double distilled water using analytical-grade chemicals purchased from Merck India. The freshly prepared 100.0 mL stock solutions of 0.1 M H2SO4, 0.5 N Fe(II), and 0.5 N Cr(VI) were prepared using 19.6070 g of (NH4)2SO4·FeSO4·6H2O and 2.4515 g of K2Cr2O7, respectively. The 0.1 M stock solutions of ligands were prepared by dissolving the calculated amount of oxalic acid (H2C2O4· 2H2O), malic acid (H6C4O5), glutamic acid (H9C5O4N), citric acid (H 8 C 6 O 7 ), iminodiacetic acid (IDA,H 7 C 4 O 4 N), diethylenetriaminepentaacetic acid (DPTA), 4,5-dihydroxy1,3-benzenedisulfonic acid disodium salt (Tiron), nitrilotriacetic acid (NTA), diphosphate (Na 2 H 2 P 2 O 7 ), and ethylenediaminetetraacetic acid disodium salt dihydrate (EDTA-Na2·2H2O). The 1 × 10−3 M Fe(III) solution was prepared from (NH4)Fe(SO4)2·12H2O. Apparatus

The cell potentials were measured using platinum indicator20 and calomel reference electrode at 25 °C under nitrogen environment in a self-fabricated assembly described in our B




Cr(VI)/Cr(III) redox couple. This lowering of Cr(VI)/Cr(III) redox potential with increase in pH makes the Fe(II)−Cr(VI) redox reaction decrease its cell potential and consequently free energy. A sequential decrease in the hydrogen ion concentration brings successive loss of free energy (feasibility), which can be witnessed from the receding potential change at the inflection point in the titration curve of Fe(II)−Cr(VI) redox reaction under different hydrogen ion concentrations Figure 1.

5.0 mL with distilled water maintaining the pH. In another set of experiments, 1.0 mL of 1 × 10−3 M Fe(III) was also taken in nine labeled 5.0 mL volumetric flasks, to which a sequential addition of (0.0, 0.5, 1.0, 1.5, 2.0, 2.5, 3.0, 3.5, 4.0) of 1 × 10−3 M Tiron ligand solution was added and their pH was adjusted around 8.5−9.0 using aqueous ammonia over a pH meter. The absorbance of both sets of solutions was recorded over the 400−800 nm wavelength range.

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HAZARDS Protective gear (goggles, lab coat, and safety gloves) should be worn while handling corrosive sulfuric acid in the fume hood. Contact with sulfuric acid cause burns to the skin and eyes. Fe(II) is a mild irritant. All chemicals used in experiment can be toxic if swallowed. Potassium dichromate is harmful by inhalation, irritating to skin, and a strong carcinogen; contact with it in every form should be avoided. Since chromium compounds are hazardous, the direct release of reaction system to the water streams (environment) should be avoided. The relatively safe disposal would be to add NaOH pellets to the titration solution on completion of experiment to precipitate metal hydroxides and separate the precipitate by decantation and pack the sludge for safe landfill disposal (see Supporting Information).

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RESULTS AND DISCUSSION The hydrogen ion is an electroactive species in the reduction of chromium(VI) to chromium(III); hence, dependence of Cr(VI)/Cr(III) redox potential on pH is obvious. Fe(II) oxidation to Fe(III) does not involve protons and as such has no direct dependence on pH, though the susceptibility and rate of Fe(II) oxidation to Fe(III) are enormously increased in alkaline pH.21 In this experiment, a pH study in acid range (0− 4) was attempted on Fe(II)−Cr(VI) redox reaction. The 0−4 pH range was selected keeping in view pH-dependent transformation of dichromate to chromate above pH 57 and increased susceptibility of Fe(II) oxidation to Fe(III) at higher pH. The formal potential of Cr(VI)/Cr(III) under different pH conditions can be worked out using eqs 3 and 4: Cr2O7 2 − + 14H+ + 6e− → 2Cr 3 + + 7H 2O E = E 0Cr(VI)/Cr(III) −

⎫ RT ⎧ [Cr 3 +]2 ⎬ ⎨ln 2− + 14 nF ⎩ [Cr2O7 ][H ] ⎭

E = E 0Cr(VI)/Cr(III) −

RT [Cr 3 +]2 ln nF [Cr2O7 2 −]

−

Figure 1. Effect of decreasing [H+] concentration on potentiometric response of Fe(II)−Cr(VI) redox reaction (each 1 × 10−1 N). Data from Table S1.

From Figure 1, it can be seen that at a lower hydrogen ion concentration of 10−4 (pH 4), the Fe(II) and Cr(VI) redox reaction becomes infeasible (as seen from absence of potential jump). This condition in which an otherwise feasible redox reaction loses its feasibility due to pH modification of formal potential has been referred to as a pH locked redox reaction. The two oxidation states of a transition metal redox couple are differently stabilized by the same ligand, for example, the formation constant of Fe(III) EDTA is nearly 1010-times larger than that of Fe(II) EDTA, whereas as formation constant of Fe(II), phenanthroline is approximately 105-times larger than that of Fe(III) phenanthroline complex.16 This differential stabilization of two oxidation states by the same ligand leads to the change in redox potential of transition metal redox couple. Thus, complexation with an appropriate ligand can be used to modulate the redox potential of transition metals redox couples from their aqua or free state value.16−20 Using this concept of complexation modulated redox potential, we attempted to “unlock” the pH locked Fe(II)−Cr(VI) redox reaction with complexation reaction of Fe(II) as a “key”. In our attempt to unlock the pH locked reaction, we screened a series of ten Fe(III) stabilizing ligands (Figure 2) with the objective to decrease the Fe(III)−Fe(II) redox potential so as to match it with the pH reduced formal redox potential of Cr(VI)−Cr(III) couple at pH 4 for a feasible redox reaction (Figure 3). From Figure 3, it can be seen that the ligands 3−8 are not capable of unlocking the reaction as no potential jump at the end point can be seen in their case. The magnitude of potential jump at the end point in Figure 3 depicts the efficacy of ligands to unlock the pH locked reaction. It can be concluded that the propensity of ligands {1, 2, 9, 10, 11} to unlock the pH locked

(3)

2.303RT × 14( −log[H+]) nF

RT [Cr 3 +]2 ln nF [Cr2O7 2 −] 2.303 × 14 × 8.314 × 298 pH − 6 × 96485



RT [Cr 3 +]2 ln − 0.138pH nF [Cr2O7 2 −]

E(pH ) = E 0 ′ − 0.138pH

(4)

From eq 4, it is clear that a decrease in formal potential by around 138 mV occurs for each unit increase in pH in case of C




Figure 2. Structures of ligands used.

Fe(III) relatively more than Fe(II) lowered the redox potential22 and were effective in the unlocking of Fe(II)− Cr(VI) redox reaction. The comparative ligand effect on the redox potential of Fe(III)−Fe(II) redox reaction can be explained using eq 5:16,23 Ecomplex = Eaqua −

β III RT ln II nF β

(5)

Where βIII and βII are the formation constants of Fe(III) and Fe(II) with the given ligand. An interesting feature of the experiment can be identification of electroactive species in titration. By varying the amount of EDTA added to the fixed molar quantity of Fe(II) taken for titration, the potential jump (end point)1 got shifted in proportion to the molar amount of {[Fe(EDTA)]2−} present in the reaction system and not the total Fe(II) (Figure 4). This observation made the distinction of EDTA complexed form of Fe(II) {[Fe(EDTA)]2−} to be an actual reductant (electroactive species) in place of free Fe(II) (Figure 4). This can be attributed to the reduction potential difference of Fe(II) EDTA complex (+ 0.08 V) and free Fe(II) (0.77 V vs NHE).1

Figure 3. Relative propensity of ligands to unlock the pH locked redox reaction of Fe(II) and Cr(VI) (each 5 × 10−2 N). Data from Table S2.

reaction is in accordance with their ability to decrease the redox potential of Fe(III)−Fe(II) redox couple. Ligands that stabilize D




values. We attempted to explore the origin of this discrepancy through a parallel experiment of speciation analysis. Adjusting pH near to 6.0 gave the anticipated larger potential change for Tiron than EDTA ligand (Figure 5). The probable explanation can be change of the stability constant and corresponding reduction potential with the change in number of Tiron chelate rings and water molecules in the Fe(II) coordination sphere under different pH.25 It was concluded that the pH 4.0 is not adequate for the formation of tris complex of Tiron ligand with Fe(II). To validate our point further, we attempted Fe(III) Tiron complex speciation analysis with pH in another parallel experiment using spectrophotometry. The experiment resulted in a distribution diagram (the fraction of predominant forms as a function of pH) depicting the predominance of different compositions for Fe(III) Tiron complex under different pH conditions (Figure 6) with their corresponding solution colors in Figure 7.

Figure 4. Effect of 2.5 × 10−2 M EDTA concentration on redox reaction of 15.0 mL of 5 × 10−2 N Fe(II) with 6.0 × 10−2 N Cr(VI) at pH 4. Data from Table S3.

This experiment also illustrates the concept of chemical speciation (predominance of different chemical forms of a given chemical moiety under different conditions), which can be assessed from the comparison of [FeII EDTA ]2− and [FeII(Tiron)2]6− redox potential (+ 0.08 V and −0.509 V vs NHE, respectively).24 These redox potentials indicate that [FeII(Tiron)2]6− to be more efficient reductant over [FeII EDTA]2− for the pH locked Fe(II)−Cr(VI) redox reaction; consequently, a larger potential change at end point is anticipated with Tiron than EDTA ligand. However, a remarkable observation was made on comparing the EDTA and Tiron ligand effect on the Fe(II)−Cr(VI)) redox reaction at two different pH (4.0 and 6.1) and is shown in Figure 5. From Figure 5, it is clear that potential jump at the end point with EDTA ligand is larger than in case of Tiron at pH 4.0 suggesting EDTA to be more efficient over Tiron for unlocking the reaction, which was disagreeing to their redox potential

Figure 6. Speciation diagram of Fe(III)−Tiron complexation at different pH (4.0, 6.0, 9.0).

Figure 7. Color changes of Fe(III)−Tiron mixture under different pH.

The composition of Fe (III) Tiron complexes under two pH values (4 and 8.5) was verified through spectrophotometric titration of Fe(III) Tiron system having a variable molar ratio of Fe(III) and Tiron ligand (Figures 8 and 9). Figures 8 and 9 verify the 1:1 and 1:3 molar ratios of Fe(III) and Tiron at pH 4.0 and pH 8.5, respectively, as reported in literature.26 The pH study reveals formation of [Fe(Tiron)]− and [Fe(Tiron)3]9− at pH 4 and 8.5, respectively. The increase of chelation (from 1 to 3) at higher pH brings enhanced stabilization of Fe (III) in tris complex with concomitant decrease in the redox potential, thus favoring effective Cr(VI) reduction at higher pH.15

Figure 5. Comparison of EDTA and Tiron ligand effect on Cr(VI)− Fe(II) redox reaction at pH 4.0 and 6.1. Data from Table S4. E




complexation under different pH conditions was done spectrophotometrically and corroborated with ligand propensity to unlock Fe(II)−Cr(VI) redox reaction at two different pH values (4 and 8.5). This work explains how a particular redox reaction can occur under environmental and biological conditions that are not analogous to the severe laboratory conditions. The work offers a mechanism for the synergistic effect of Fe(II) and organic moieties (polyphenolates and carboxylates) toward Cr(VI) reduction and can be explored for the natural attenuation of other redox active environmental toxins to their non- or lesser-toxic forms. The experiment also describes safe practice of laboratory waste disposal for toxic chromium species. Pedagogy

The experiment has three main components: pH locking of reaction, complexation effect (using EDTA ligand) for unlocking of pH locked reaction, and identification of electroactive species in the titration. The experiment gets successfully completed in two consecutive laboratory sessions of 3-h duration (each student in a group of five students of advanced/senior level inorganic laboratory course performs one potentiometric titration and shares the data with others in the group to complete the study of an experimental component, complete lab details in Supporting Information). This laboratory exercise is aimed to demonstrate the pH effect on redox potential of Cr(VI)−Cr(III) redox couple and its influence on Fe(II)−Cr(VI) redox reaction. The students are expected to make a distinction between standard reduction potential and conditional or formal reduction potential of a redox couple so as to broaden the vision of a student to think beyond electrochemical series in the design of redox systems for analytical applications. The students are expected to learn the complexation effect on the redox potential of iron(III)− iron(II) redox couple and correlate the impact of Fe(III) complex stability with different ligands on the redox potential of a Fe(III)−Fe(II) redox couple. This experiment can serve as a model for student conception of natural redox processes, action of metaloenzymes, and metal extraction by siderophores, besides offering insights to redox reactions occurring in environmental and biological conditions that are not analogous to severe laboratory conditions. The experiment has an interesting environmental context in the form of concept of natural attenuation of toxic contaminants and prudent practices for disposal of toxic waste from laboratories. This will help build the thinking ability of the students for converting environmentally friendly laboratory practices into a real-world scenario. Besides these salient features, this laboratory exercise illustrates some basic chemistry concepts to the students: (i) pH dependence of the redox potential in case of systems involving H+ ions in the redox reaction. (ii) Nernst equation prediction of the cell potential change with pH, for example, the redox potential of Cr(VI)/ Cr(III) decreases by 138 mV with every unit increase in pH. (iii) A feasible reaction can become blocked under unfavorable reaction conditions. (iv) Differential stabilization of the two oxidation states in a transition metal redox couple by same ligand. (v) Complexation shifts redox potential of transition metal redox couple due to differential stabilization of two oxidation states in a redox couple.

Figure 8. Spectrophotometric titration at pH 4.0 of 1 × 10−3 M each of Fe(III) and Tiron showing a 1:1 mol ratio at its equivalence point. Main graph: spectra of various Fe(III)−Tiron mixtures. Insert: maximum absorbance as a function of mole ratio for each curve.

Figure 9. Spectrophotometric titration at pH 8.0 of 1 × 10−3 M each of Fe(III) and Tiron showing a 1:3 mol ratio at its equivalence point. Main graph: spectra of various Fe(III)−Tiron mixtures. Insert: maximum absorbance as a function of mole ratio for each curve.

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CONCLUSION The experiment describes dependence of redox potential on pH through potentiometric titration of Fe(II)−Cr(VI) redox reaction over a pH range of 0−4. The successive diminishment of the potential jump with increasing pH depicts the feasibility loss of Fe(II)−Cr(VI) redox reaction above pH 4. Adding a ligand that preferentially complexes Fe(III) reduces the Fe(III)−Fe(II) redox potential and restores the feasibility of the titration locked by pH. This effect was referred to as unlocking the pH locked redox reaction. The efficacy of the selected group of ligands toward unlocking was found to be in accordance with their capacity to decrease the Fe(III)−Fe(II) reduction potential. The speciation analysis of Fe(III) tiron F




(10) Barrera-Díaz, C. E.; Lugo-Lugo, V.; Bilyeu, B. A review of chemical, electrochemical and biological methods for aqueous Cr(VI) reduction. J. Hazard. Mater. 2012, 223−224, 1−12. (11) Nancharaiah, Y. V.; Dodge, C.; Venugopalan, V. P.; Narasimhan, S. V.; Francis, A. J. Immobilization of Cr(VI) and Its Reduction to Cr(III) Phosphate by Granular Biofilms Comprising a Mixture of Microbes. Appl. Environ. Microbiol. 2010, 76 (8), 2433−2438. (12) Dermatas, D.; Chrysochoou, M.; Moon, D. H.; Grubb, D. G.; Wazne, M.; Christodoulatos, C. Ettringite-Induced Heave in Chromite Ore Processing Residue (COPR) upon Ferrous Sulfate Treatment. Environ. Sci. Technol. 2006, 40 (18), 5786−5792. (13) Strathmann, T. J.; Stone, A. T. Reduction of oxamyl and related pesticides by FeII: Influence of organic ligands and natural organic. Environ. Sci. Technol. 2002, 36, 5172−5183. (14) Buerge, I. J.; Hug, S. J. Influence of organic ligands on chromium(VI) reduction by iron(II). Environ. Sci. Technol. 1998, 32 (14), 2092−2099. (15) Naka, D.; Kim, D.; Strathmann, T. J. Abiotic Reduction of Nitroaromatic Compounds by Aqueous Iron(II)-Catechol Complexes. Environ. Sci. Technol. 2006, 40, 3006−3012. (16) Rizvi, M. A.; Syed, R. M.; Khan, B. U. Complexation Effect on Redox Potential of Iron(III)-Iron(II) Couple: A Simple Potentiometric Experiment. J. Chem. Educ. 2011, 88, 220−222. (17) Rizvi, M. A.; Teshima, N.; Peerzada, G. M. 1,10-Phenanthroline Modulated Redox Potentials Explored for Benign Iron Speciation Analysis. Croat. Chem. Acta 2013, 86 (3), 345−350. (18) Rizvi, M. A. Complexation Modulated Redox Behavior of Transition Metal Systems. Russ. J. Gen. Chem. 2015, 85 (4), 959−973. (19) Raashid, S.; Rizvi, M. A.; Khan, B. U. Coordination inspired redox behaviour of Fe (II) and Co (II) explored for simultaneous Iron oxidation state analysis. J. Pharm. Res. 2012, 5 (5), 2715−2720. (20) Peerzada, G. M.; Rizvi, M. A.; Teshima, N. Utilizing Fe(III)/ (II)-EDTA Couple for Estimation of Transition Metal Ion Mixture Over Platinum Electrode. Asian J. Chem. 2013, 25 (9), 4776−4778. (21) Morgan, B.; Lahav, O. The effect of pH on the kinetics of spontaneous Fe(II) oxidation by O2 in aqueous solution − basic principles and a simple heuristic description. Chemosphere 2007, 68 (11), 2080−2084. (22) Harrington, J. M.; Crumbliss, A. L. The redox hypothesis in siderophore mediated iron uptake. BioMetals 2009, 22 (4), 679−689. (23) Kosman, D. J. Iron metabolism in aerobes: Managing ferric iron hydrolysis and ferrous iron autoxidation. Coord. Chem. Rev. 2013, 257, 210−217. (24) Naka, D.; Kim, D.; Carbonaro, R. F.; Strathmann, T. J. Abiotic reduction of nitroaromatic contaminants by iron(II) complexes with organothiol ligands. Environ. Toxicol. Chem. 2008, 27, 1257−1266. (25) Cuculic, V.; Pizeta, I.; Branica, M. Voltammetry of Dissolved Iron(III)−Nitrilotriacetate−Hydroxide System in Water Solution. Electroanalysis 2005, 17 (23), 2129−2136. (26) Liu, J.; Ma, H. Investigation of the composition of complexes and the stoichiometry of non-complex reactions by flow injection method. Talanta 1993, 40 (7), 969−974.

(vi) The very high formation constant of Fe(III) with EDTA lowers the redox potential of iron redox couple. (vii) Ligands can unlock the pH locked Fe(II)−Cr(VI) redox reaction as per their ability to lower iron redox potential. For a successful completion of the experiment and stoichiometric end points, titrations should be performed with fresh solutions under nitrogen atmosphere with constant stirring. EMF values stabilize within a minute with EDTA ligand but may stabilize up to 2 min for electrode equilibration with other weaker ligands. A molar ratio of Fe(II)-to-ligand slightly higher than required for octahedral composition produces good results.

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ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available on the ACS Publications website at DOI: 10.1021/acs.jchemed.5b00499. Student instructions, experimental data from which graphs of manuscript have been plotted, derivations, calculations (PDF, DOCX) Report form, instructor notes (PDF, DOCX)

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AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected] or [email protected]. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS The authors wish to thank Professor Alvin L. Crumbliss, Department of Chemistry, Duke University, Durham, North Carolina 27708, United States for helpful discussions. The constructive comments of anonymous reviewers are respectfully acknowledged.

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REFERENCES

(1) Harris, D. C. Quantitative Chemical Analysis, 5th ed.; W. H.Freeman and Company: New York, 1999; pp 161, 313−315, 421. (2) Walczak, M. M.; Dryer, D. A.; Jacobson, D. D.; Foss, M. G.; Flynn, N. T. pH Dependent Redox Couple: An Illustration of the Nernst Equation. J. Chem. Educ. 1997, 74 (10), 1195. (3) Pesterfield, L. L.; Maddox, J. B.; Crocker, M. S.; Schweitzer, G. K. Pourbaix (E−pH-M) Diagrams in Three Dimensions. J. Chem. Educ. 2012, 89 (7), 891−899. (4) Shriver, D. F.; Atkins, P. W.; Overton, T. L.; Rourke, J. P.; Weller, M. T.; Armstrong, F. A. Inorganic Chemistry, 5th ed.; Oxford University Press: Oxford, 2010; pp 156, 168. (5) Vidal-Iglesias, F. J.; Solla-Gullon, J.; Rodes, A.; Herrero, E.; Aldaz, A. Understanding the Nernst Equation and Other Electrochemical Concepts: An Easy Experimental Approach for Students. J. Chem. Educ. 2012, 89 (7), 936−939. (6) Serjeant, E. P. Potentiometry and Potentiometric Titrations. In A Series of Monographs on Analytical Chemistry and Its Applications; Elving, P. J., Winefordner, J. D., Eds.; John Wiley & Sons: New York, 1984; Vol. 69, p 38. (7) Kalbus, L. H.; Petrucci, R. H.; Forman, J. E.; Kalbus, G. E. Titration of chromate-dichromate mixtures. J. Chem. Educ. 1991, 68 (8), 677. (8) Barreto, M. S.; Medeiros, L. L.; Furtado, P. C. Indirect Potentiometric Titration of Fe(III) with Ce(IV) by Gran’s Method. J. Chem. Educ. 2001, 78 (1), 91. (9) Young, J. A. J. Chem. Educ. 2003, 80 (8), 874. G


Complexation Key to a pH Locked Redox Reaction - ACS Publications

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