Development of novel techniques to determine hydrolysis constants of the rare earths, yttrium and iron

Citation
Development of novel techniques to determine hydrolysis constants of the rare earths, yttrium and iron

Material Information

Title:
Development of novel techniques to determine hydrolysis constants of the rare earths, yttrium and iron
Creator:
Klungness, Greta D. S.
Place of Publication:
Tampa, Florida
Publisher:
University of South Florida
Publication Date:
Language:
English
Physical Description:
vii, 50 leaves : ill. ; 29 cm.

Subjects

Subjects / Keywords:
Hydrolysis ( lcsh )
Yttrium ( lcsh )
Rare earth metals ( lcsh )
Dissertations, Academic -- Marine science -- Masters -- USF ( FTS )

Notes

General Note:
Thesis (M.S.)--University of South Florida, 2000. Includes bibliographical references (leaves 41-44).

Record Information

Source Institution:
University of South Florida
Holding Location:
Universtity of South Florida
Rights Management:
All applicable rights reserved by the source institution and holding location.
Resource Identifier:
028176741 ( ALEPH )
48123880 ( OCLC )
F51-00151 ( USFLDC DOI )
f51.151 ( USFLDC Handle )

Postcard Information

Format:
Book

Downloads

This item is only available as the following downloads:


Full Text

PAGE 1

DEVELOPMENT OF NOVEL TECHNIQUES TO DETERMINE HYDROLYSIS CONSTANTS OF THE RARE EARTHS, YTTRTIJM AND IRON by GRETAD.S. KLUNGNESS v A thesis submitted in partial fulfillment of the requirements for the degree of Master of Science College ofMarine Science University of South Florida December 2000 Major Professor: Robert H. Byrne, Ph.D.

PAGE 2

Office of Graduate Studies University of South Florida Tampa, Florida CERTIFICATE OF APPROVAL Master's Thesis This is to certify that the Master's Thesis of GRETA D.S. KLUNGNESS with a major in Marine Science has been approved for the thesis requirement on October 3, 2000 for the Master of Science degree Examining Committee : Major Professor: Robert H. Byrne Ph D. .R. Betze& Ph D Member: Paula G. Coble, Ph.D.

PAGE 3

TABLE OF CONTENTS LIST OF TABLES .................... ............................................................... iii LIST OF FIGURE S ................... ... ... . ......... ............................................. iv ABSTRACT .............. . ............ .......................................................... ..... vi CHAPTER 1. COMPARATIVE HYDROLYSIS BEHAVIOR OF THE RARE EARTHS AND YTTRIUM : THE INFLUENCE OF TEMPERATURE AND IONIC STRENGTH .. Introduct ion................. . .. .. ...... . . ... .. .. . . .. . ............. ... .. .. . ....... . .. 1 Ana l ytical Procedures....................... ...................... ...... .. ......... . ...... 4 Theory ................... .......... . . ... .. .. .. ........... . .. . .................... 4 Met ho ds . . ...... . ... . ...... ....... .......... .. ...................... .. ... ... .... 6 Materials. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 9 Results ............................... .. ................... ...... ... .. ... .. .. .................. 10 YREE hydroly sis Patterns ........ . ............. ...................... ... . ........ 10 Influence of ionic strength on l og 1 (M) .. .................. .. ................. 15 Temperature dependence ofYREE hydrolysis constants . ..................... 16 Conclusions .. ...... ........ .. . . .. . . ...... .. .............. . . .. . .. ......... . ...... . 19 Acknowledgements ........................ ..... . .. ...... .. .. ... . . . . ........ 20 CHAPTER 2 DEVELOPMENT OF NOVEL SPECTROPHOTOMETRIC PROCEDURES FOR EXAMINATION OF FERRIC ION HYDROLYSIS . ........... . .. 21 Introduction .................................................................................... 21 Theory ..................... .. . .. ... . . . ... ..................................... ........... 23 Potentiometric Experiments ........... .......... .. .. ...... ..... .. . .. . . . ......... 24 Sp ectrophotometric Experiments ......... ... ...... ................... . ....... . . .. 25 Metho ds and Materials ................. . .. ........... ........................................ 27 Materials ..... . ........... . . . ..................................... ... . ... .. .. ... 27 Methods-Potent iome try .. .... ......................................................... 28 Methods Spectrophotometry ....... ............... ....... .. ......... .. . . .. ....... 28 Results and Discussion ............... . ........... .. . .. ................................... 31 REFERENCES .................. ............. ... . ...... .. ...... . .. . .. ... ... .. .. ......... . .... 41 APPENDICES ............. .. .. . . ........ .. ..... . ... .. .. .......... .. . .. .. ...... .. ... . ... .. . . 45 Appendix A: Previous 1 data (from Figure 1).. .. .. .. . .. .. .. . .. . .. .. . . . . 46 Appendix B: data at 55C, 40C and 25C .. ...... .. ... . ......... 47

PAGE 4

Appendix C : data at 25C and 1=0.7 ........ ................. 47 Appendix D : Absorbance and [NaSCN] data used to calculate and at constant pH ....... ........................ .. ..................... 48 Appendix E: Thiocyanate absorbance data at 460 nm from Experiment 1 .. ... . .... 49 Appendi x F: Thiocyanate absorba nce data at 460 nm from Experiment 2 .... ....... 4 9 Appendix G: Thiocyanate absorbance data at 460 nm from Experi ment 3 . . . . ... 50 Appendi x H: Thiocyanate absorbance data at 460 nm from Experimen t 4 . . . .. . 50 11

PAGE 5

LIST OF TABLES Table 1. Enthalpy Data (kcal mor1 ) for YREE H y drol ys is ............................. .... 18 Table 2 YREE Hydrolysis constants at 25C are shown for I = 0 7 mola r 1 (M) S .E.) and I= 0 0 ......... .. . ....... . ..... ........ . .. ... .. ...... . 1 8 Table 3. Results from potentio metric experiments .... ..... ... ................. . .......... 32 Table 4. R es ults offive experimental d e t e rminations of and ...... . ... 33 Table 5. Summary of critically sel ect ed constants ....................................... . ... 34 Table 6. Summary res ults using e quation (27) and pH 3 7 ................... 37 Table 7. Summa t y of 1* r es ults using equation (27) with [Fe111]r adjusted to g ive a z ero intercept ..................... ..................... ................. ........... 37 Table 8. Previous d e terminations of ............................................ ........... 38 lll

PAGE 6

LIST OF FIGURES Figure 1. Hydrolysis constant estimates at 25 and zero ionic strength. Ionic strength corrections were performed using the hydrolysis behavior ofFem as an appropriate analog..................................... ...... . .. .. ................... 3 Figure 2. Representative Potentiometric data vs time . ................................... ... 1 1 Figure 3. Representative Spectrophotometric data vs. time ................................. 12 Figure 4. logp 1 *(M) determinations at I=0 .7 molar and 25C: Average of5 spectrophotometric experiments wit h standard errors plus the results of two potentiometric experiments ..................................... .. .................... 13 Fig u re 5. Average logp 1 *(M) determinations for all experiments at 1=0.7 molar and 25C plus the results of two potentiometric expe rim ents at 25C and I= 0.1 molar. The standard errors shown in this figure refer to the seven experiments performed at 0.7 molar ionic strength ................................ 14 Figure 6. Ionic Strength Dependence oflogp1 *(Ho) at 25C. Note that the difference (:: 0.05) between logp 1 *(M) for heavy REEs at 0.7 mo l ar and 0.1 molar ionic strengths (Figure 5) is consis tent with the observations shown in this figure ....... ............................................... ............................. 16 Figure 7. Potentiometric l ogp 1 *(M) results (0 7 molar) at 25C, 40 C and 55C. Enor lim its shown in this figure indicate the range in logp1 *(M) for two measurements at each temperature ................................................ ... 1 7 Figure 8. Absorbance spectra showing influence of pH on FeSCN 2+ and FeOH2+ absorbance ...................................... ..... ..... ....... . .. . ......... . . .. 30 Figure 9. Absorbance spectra ofFe111 in 0.7 NaC104 between pH 2.7 and 4 ............... 31 Figure 10. Potentiometric data used for determ i nation of scNPJ and scNP2 ............... .. 32 Figure 11. Data used for determination ofE1scNP 1 and E2scNP2 ............................... 33 Figure 12a Data used for determination of P 1 ................................................... 35 IV

PAGE 7

Figure 12b. Data u sed for determination ofP1* ................................................... 36 Figure 13. FeSCN absorbance in NaCI04 NaN03 and NaC l. .. .. ..................... . . 40 v

PAGE 8

DEVELOPMENT OF NOVEL TECHNIQUES TO DETERMINE HYDROLYSIS CONSTANTS OF THE RARE EARTHS, YTTRIUM AND IRON by GRETA D S. KLUNGNESS An Abstract of a thesis submit ted in partial fulfillment of the r e quir e m ents for the degree of M as ter of Scienc e College of Marine Science Univ e rsity of South Florida December 2000 Major Professor: Robert H Byrne, Ph. D VI

PAGE 9

In the first chapter of this work potent iometric and spectrophotometric techniques are used to determine hydrolysis constants ofYttrium and the rare earths (YREEs), where = [MOH2+ ][H+][M3+r1 Although they were conducted on different time scales (as long as fifteen minutes for potentiometry vs. less than 500 milliseconds for spectrophotometry), the stability constant patterns produced by both procedures are in excellent agreement. Potentiometric results obtained over a temperature range between 25C and 55C indicated that enthalpies appropriate to YREE h ydrolysis can be reasonably approximated as m0(M) = 11.3 kcal mor1 for all YREEs. The dependence ofYREE hydrolysis constants on temperature (T), ionic strength (I), and YREE identity (M) can be expressed as (M) = (M)-{ 2.044I}i /(1 + 5.52I }i ) + 1.84 x 10-31}2446K/ T + 8.204 In the second chapter of this work, sodium thiocyanate was used in potentiometric and spectrophotometric experiments to determine the first hydrolysis constant of iron (p1*(Fe)) Sodium thiocyanate reacts with iron to form a complex (FeSCN2+ ) that absorbs strongly in the visible ( 460nm). The results obtained in this work indicate that observations ofFeSCN2+ absorbance in long pathlength cells can be used to extend previous spectrophotometric examinations ofFe3+ hydrolysis to lower total iron concentrations and higher pH Abstract Approved: -------------,.----r-------Major Professor: Robert H.1Byme, Ph.D Professor, College of Marine Science Date Approved: __:::D::::.....(!...CJ=' __ Vll

PAGE 10

CHAPTER 1 COMPARATIVE HYDROLYSIS BEHAVIOR OF THE RARE EARTHS AND YTTRIUM: THE INFLUENCE OF TEMPERATURE AND IONIC STRENGTH llltroduction Solut ion equilibria strongl y influence the environmental behavior of the trivalent rare earths and yttrium (YREE). hnportant environmental YREE species include carbonate hydroxide, sulfate fluoride and chloride comple x es. Among all the equilibri a curr e ntl y known to influ e nc e YREE solution speciation in natural waters hydrolysis equilibria are probably the least well understood The first hydrolysis step for trivalent YREEs, written as M3+ +H 2 0 <=> MOH 2 + +H+ can be quantitativel y cha racteri z ed with hydrol y sis constants written in the fonn (M) = [MOH 2 + ][H + ][M3+rl (1) (2) where M3+ is a trivalent YREE and brackets denote concentrations of each indicated species. Previous experimental determinations 1 (M) do not provide a coherent representation ofYREE hydrol y sis behavior The log *(M) data shown in Figure 1 at z ero ionic str e ngth e x hibit an ov e r a ll range of approximatel y two orders of magnitude. These data show no discemable pattern in the magnitudes of log 1 (M) across the YREE series other than a general but poorl y defined, increase between the lightest rare e a rth 1

PAGE 11

elements (REEs) and the heaviest REEs. In addition to the very large range in previous log pi*(M) estimates for the YREEs, there are essentially no data relating YREE hydrol ysis constants to either temperature or ionic strength. In view of the extreme paucity of experimental log P 1 (M) data obtained using consistent procedures for all YREEs, and the absence of log p 1 (M) data obtained over a range of temperatures and ionic strengths, further experimental observations are required over a range of conditions appropriate to the natural environment. In this work two procedures hav e been used to determine the hydrolysis constants of yttrium and all rare earths (except Pm). Both methods involve measurement of pH in solutions buffered by M3+ /MOH2+ hydrogen ion exchange-couples In the first procedure pH is measured spectrophotometrically with meta cresol purple, a pH indicator calibrated previously (Clayton and Byrne, 1993) at ionic strengt h s appropriate to seawater. Spectrophotometric pH measurements in systems buffered by M3+ IMOH2+ couples involved the use of a stopped flow rapid scan spectroscopy system to observe pH via indicator absorbance ratios on time scales of ze ro to 500 milliseconds. This procedure was used to separate the first hydro l ysis step (equation 1) from slower hydrolysis steps leading to formation of solid M(OHh(s) The second measurement procedure involved e l ectro d es to measure pH in solu tion s with much higher total metal concentrations but lower pH (higher [M 3 + ] / [MOH2+ ] ratios) than those in our spectrophotometric experiments. These measurements at low pH substantially lowered the saturation state of the solutions with respect to formation of solid M(OH)3(s). The overall expe rimental design acknowledges the importance of obtaining consistent results using independent methode logies 2

PAGE 12

0 II c .... "0 .... Col 1-c 1-c c Col ----l< ..... C!l.. bJl c --3 4 -5 -6 -7 -8 -9 -10 -11 La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu REE Fi g ure 1. Hydrolysis constant estimates at 25 and zero ionic strength Ionic strength corrections w e re performed using the hydro l ysis behavior ofFe111 a s an appropriate analog (Milburn and Vosburgh, 1955) Ref erence Frolova et al. 1966 1 4 Caceci and Choppin 1983 2 Guillamont et al. 1971 15 Din Ngo and Burkov 1974 3 Kragten and Decnop-Weever, 1978 16 Davydov and Voronik 1983 4 Kragten and Decnop-Weever 1979 17 l v anov-Emin et al., 1970 5 Kragten and Decnop-Weever, 1980 18 Lope z -Gonzale z et al., 1997 6 Kragten and Decnop-Weever, 1982 19 Moeller 1946 7 Kragten and Decnop-Weever, 1 983a,b 20 Nair et al., 1982 8 Kragten and Decnop-Weever 1984 21 Schmidt et al. 1978 9 Kragten and Decnop-Weever, 1987 22 Tobias and Garrett 1958 10 Amaya eta!., 1 973 23 Usherenko and Skorik 1972 11 Burkov et al., 1975 24 Wheelwright et al. 195 3 12 Burkov et al., 1982 25 Baes and Mesmer, 1976 13 Biedermann and Ciavatta, 1961 26 Lee and Byrne, 1992 3

PAGE 13

Direct comparisons are obtained between potentiometrically and spectrophotometrically derived log pI. (M) results at 25C and ionic strength en equal to 0 7 molar. Additional potentiometric results at !=0.7 molar are used to determine log p1.(M) at 40C and 55C Finally, log P 1 (M) is determined at I =O.l molar and 25C for all YREEs, and the ionic strength dependence oflog P1.(Ho) is observed at 25 for ionic strengths between 0.04 and 5.4 molar. Analytical Procedures Theory In this work, the potentiometric and spectrophotometric procedures used to determine log P 1 (M), account for all significant acid-base equilibria. The metal ion buffers used in potentiometric experiments were prepared by combining acidic solutions, containing HCl04 and either dissolved M(Cl04 )3 or M(N03 ) 3 with solutions ofNaOH. These acid and base titrants were added to solutions containing NaC10 4 as the background electrolyte. The charge balance equation describing the final buffered solution at equilibrium is 3[M3+ ] + 2[MOH2+ ]+[H + ]+[Na + ]-[CI0 4 -]-[0H-] = 0 (3) The concentrations ofM3+ and MOH2+ in this buffer can be expressed as (4) where MT is the total concentration of a given YREE The (Na + ] and [CI04-] concentrations in equation (3) have four sources: NaCI04, HCI04, M(CI04h and NaOH. The Na+ and Cl04 added as NaCl04 are ba lanced ([Na + ] = [Cl04-]) and need not be considered in equation (4). As such the difference between [Na + ] and [Cl04-] in 4

PAGE 14

equation (4) equals the [Na+ ] added as NaOH minus [C10 4-] added as HC104 and M(Cl04)3. Thus we can write (5) where [.6.] can be conceptuali ze d as the concentration ofbase (NaOH) added in excess of acid (HCl04), and 3MT is the concentration of Cl04 in solution from added M(Cl04)3. Combining equations (3)-(5), equation (2) can be written as =pH+ log w { (D.)+[H+ ] K [H+ ]-1 } M T ( .6. ) + K w [ H + ] -I [ H + ] (6) Under the conditions of each experiment both [H + ] and [OH-] (i.e. K w[H+r1 ) are small(:::; 21-1mol dm-3 ) whereby the final term in equation (6) is nearly equal to the term log [MOH2+]/[M3+]. For each experiment, performed at constant temperature and ionic strength, th e terms .6. and MT-.6. were held essentially constant. In this case, the pattern of log 1 (M) values in each experiment was very nearly identical to the pattern of directly observed equilibrium pH values With one except ion The equations used in the spectrophotometric experiments are identical to those used in the potentiometric experiments. Spectrophotometric experiments contained two forms of m-cresol purple, L2 -and HL-. Consequently two additional tern1s must be added to equation (3). These terms -2[L2-]-[HL], are added to the left side of equation (3). Additionally, since the indicator solutions were prepared by dissolving Na2L in NaOH, the [Na +] term in equation (3) must include the additional Na + ad ded to solutions along with m-cresol purple. In spectrophotometric experiments, log 1 (M) values can also be calculated using equation (6). However, in this case [D.] includes the following term: 5

PAGE 15

2[L]T-2[L 2-]-[HL-] = [HL] (7) Consequently, whereas equation (6) is used to directly calculate log in potentiometric experiments as used in spectrophotometric experiments one additional term ([HL-]) is added to the term in equation (6) Methods Potentiometric log 1 (M) determinations were conducted by mixing two solutions and measuring solution pH over a fifteen minute period. One of these two solutions consisted of a rare earth element, total concent r ation 0 .01 molar, in NaC104 at [H + ] =1x10-4. The second solution was a 0.1 N volumetric standard NaOH solution. The mixture of these solutions produced a buffering ratio on the order of0. 11, whereby the measured solution pH was approximately one log unit below the log 1* (M) appropriate to a given element. Two experiments were conducted at 25C and 0.7 molar ionic strength (I) and two experiments at I= 0.7 molar were conducted at two higher temperatures (40C and 55C) Two additional experiments were conducted at 25C and 0.1 molar ionic strength. The procedures described above were conducted using all YREEs except Pm. Experiments at 25C were also conducted using a single element (Ho) over a range of ionic strengths between 0.04 and 5.4 molar. The glass electrode used in these experiments was calibrated on the free hydrogen ion concentration scale by titrating unbuffered NaCl04 with a standard HC104 solution which had been calibrated against a standard NaOH solution (0. IN NaOH). In all cases, the Orion Ross-type combination electrode used in this work exhibited Nemstian behavior. The acidic solutions used in each experiment were freshly prepared prior to 6

PAGE 16

each experiment. Each acidic solution was stored in a plastic beaker under an atmosphere of ultra high purity N2 and was thermostated at the desired experimental temperature prior to mixing. The basic solution was kept in its original container and placed in the temperature bath to maintain the appropriate temperature prior to mixing Experiments were initiated by combining a small volume of the standard NaOH solution with a large volume of each solution containing a 0 .01 molar YREE. The NaOH addition was made with a variable capacity Eppendorf pipette. All pipettes used in this work were calibrated gravimetrically. Replicate gravimetric / volumetric analyses indicated that volumetric additions were accurate to 1% or better In all potentiometric experiments, a Lauda K-4R thermocooler was used to provide constant thermal conditions (25C, 40C 55C). These temperatures were monitored directly in each experimental solution and were constant to 0.1 C Prior to additions of the NaOH titrant the combination electrode was equilibrated in the acidic YREE-containing so lution until the electrode potential was constant within 0 1mv for 5 minutes The small K w[H+ri terms in equation (6) were calculated using the ionic strength dependence of Kw in NaCl04 s ummarized by Baes and Mesmer (1976), and the temperature dependence of Hamed and Owen (1958) The procedures used in spectrophotometric experime nts were nearly identical to the procedur es used in the potentiometric work. In spectrophotometric experiments the pH term in equation (6) is calculated as (8) 7

PAGE 17

where pK2 refers to the HLdissociation constant, R is a measured absorbance ratio at the absorbance maxima ofL2 and HL(578 and 434nm), and the terms e1 e 2 and e3 are molar absorptivity ratios (Baes and Mesmer, 1976 ; Byrne, 1987; Robert-Baldo et al., 1985). The use of meta cresol purple and equation (8) to measure the pH of seawater on the total hydrogen ion concentration scale is described in Clayton and Byrne (1993) On the free hydrogen ion concentration scale pH (-log[W]) is higher than the total scale pH (expressed as pHr= -log[W]T where [H+ ]T= [W]+[HS04 ]) by 0.14log units (Breland and Byrne 1993). Consequently, pK2 = 8.1456 at 25C and I = 0.7 molar ionic strength Subsequent to calculating pH in each spectrophotometric experiment using absorbance ratios (R), the term [HL-] which is included in the [t:l] term of equation (8) is given as (9) Spectrophotometric experiments involved the use of an OLIS RSM1 000 to measure absorbance ratios at a rate of 1000 scans per second. Using absorbance observations at the isosbestic point of m-cresol purple the instrument was adjusted to mix two solutions in a one to one ratio. One solution cons i sted of 4x 1 o -5 molar m-cresol purple (25C) and either 1xl0-3 molar or 6x10-4 molar NaOH in 0.7 molar NaCl04 (25C). The second solution consisted of a 2x10-3 molar YREE in 0.7 molar NaCl04 (25C) at [H + ] =lx 10-4. Each of these solutions as well as the mixing and measurement chambers were submerged in a thetmostated solution whose temperature (25C) was constant within 0.1 C. Nine absorbance scans, recorded against an absorbance baseline obtained usin g pure 0.7 molar NaC10 4 were used to construct an average pH profile for each YREE. 8

PAGE 18

Maximum absorbance ratios, generally obtained between 70 and 80 milliseconds, were used to calculate the pH used for logp1.(M) calculations in equation (6). Materials Potentiometric analysis (pH determinations) were conducted usin g a Coming Model129 pH meter in the absolute millivolt mode. The pH electrode used in this work was an Orion Ross-type combination electrode (Orion No. 81 02BN). The volumetric pipettes us e d in this were Eppendorfvariable capacity 100-lOOOf..lL and 1000-5000f..lL pipettes. Yttrium and rare earth element oxides (99.99%) were obtained from Aldrich, Milwaukee, WI. Ce, Pr, and Tb were obtained as Ce(N03 ) 3 Pr(N03 ) 3 and Tb(N03 ) 3 (99.99%). YREE stock solutions were prepared by dissolving salts in perchloric acid to make solutions that were 0.1 molar in each YREE and 0 .01 molar in HC104 Solutions were stored in glass volumetric flasks (250mL) and the YREE concentration of each flask was determined via ICP-MS analysis using YREE s tandards obtained from SPEX (Metuchen, NJ). Anhydrous NaC104 was obtain e d from Sigma, (St. Louis, MO) NaC104 (4 molar) stock so lutions were prepared by (a) dissolving NaCl04 in MilliQ water (Millipore Corporation, Bedford, MA), (b) adding sufficient NaOH to raise th e pH to values between 10 and 11, (c) filtering these so lution s through clean 0.4f..l polycarbonate membrane filters (Poretics Products, Livermore CA) after storage in a sealed volumetric flask overnight. To avoid uptake of C02, each solution was subsequent l y acidified to pH 4.0 and was likewise stored in a glass volumetric flask Perchloric acid was obtained from Aldrich (Milwaukee WI). The concen tration s of concentrated HC104 standard so lutions were determined titrametrically using 9

PAGE 19

volumetric standard NaOH solutions obtained from Fisher Scientific (Pittsburgh, P A). Meta cresol purple salts were obtained from Eastman Kodak (Rochester NY) Ultra high Purity N2 was obtained from Air Products and Chemicals Inc. (Allentown, P A) Results YREE hydrolysis patterns Equilibrium pH values for each potentiometric experiment were based on millivolt readings recorded 5 minutes after mixing. Typical pH profiles through time are shown in Figure 2 Typical pH profiles for spectrophotometric experiments are shown in Figure 3. Due to polymerization/precipitation reactions, the decrease in solution pH after the observed maxima was anticipated at the high pH of these experiments. Figure 4 shows the results of two potentiometric log P1 *(M) determinations at 25C along with the average log p 1 (M) results for five spectrophotometric experiments at 25C. The standard errors shown in Figure 4 refer to the spectrophotometric average. The potentiometric results in Figure 4 were obtained using a buffer ratio ([MOH 2 + )/[M3+ ]) on the order of 0 .11, whereby the measured pH in each experiment was approximately one unit below the Figure 4 calculated log p 1* (M) values. Two buffer ratios were used in the spectrophotometric experiments, [MOH 2+)/ [M3+ ] = 0.43 and [MOH 2 + )/[M3+ ] = 1 These buffer ratios produced higher solution pH than the potentiometric experiments and thereby created conditions more favorable to formation of precipitates and polymers The concordance of the potentiometric and spectrophotometric results, obtained over a wide range ofM(OH)J(s) saturation conditions, indicates that polymerization reactions and precipitation are unimportant in these experiments While the absolute magnitudes of spectrophotometrically derived log P 1 (M) values varied significantly in these 10

PAGE 20

experiments (as reflected by the s t andard errors shown in the figure), the pattern of log 1 (M) values across the series of YREEs was highly consistent in every spectrophotometric and potentiometric experiment. This observation indicates that the log 1 (M) standard errors shown for each YREE are highly correlated. 8 7 8 7 7.7 y 7 .7 6 7 6.7 La =a. 5 7 :I: 5.7 c. 4 .7 4 .7 3 7 3.7 2 7 2 7 0 0.5 1.5 2 5 0 0.5 1.5 2 2.5 Time [log(s)J Time [ log(s)l 8 7 8.7 7 .7 Nd 7 7 Eu 6 .7 6 .7 =a. 5.7 =a. 5 .7 4 .7 4 7 3.7 3 7 2 7 2 .7 0 0 5 1.5 2 2 .5 0 0 .5 1.5 2.5 Time [log(s)J T i me [log(s)[ 8.7 8 7 7 7 Ho 7 7 Lu 6 7 v 6 7 =a. 5 7 =a. 5 7 / 4 7 4 7 3 7 3 7 2 7 2. 7 . 0 0 5 I 1.5 2 2 .5 3 0 0 .5 I 1.5 2 2 5 3 Time [log( s)J Time [l og(s)[ -Figure 2. Representative Potentiometri c data vs. time Notably, the averaged log results obtained using spectrophotometric procedures are in excellent agreement with the result of each potentiometric experiment. The only 11

PAGE 21

exception to this generalization is seen in the hydrolysis behavior of Ce. Spectrophotometric experiments, conducted on very short time scales, provide a sequence of (M) values for La, Ce and Pr that are closely consistent with the patterns of hydrolysis constants estimated for these elements using linear free energy relationships (Lee and Byrne, 1992). While it is possible that accessibility to th e tetravalent oxidation state, unique to Ce in aqueous solutions, may be responsib le for the observed Ce 7 .7 y 7.6 ::c c. 7.S 7 .4 so 200 Time (ms) 3SO soo 8.2 8 1 Nd 8 7.9 7.8 7.7 7.6 s o 200 T i me (ms) JSO soo 7 4 Ho 7 .3 ::c c. 7 2 7 1 so 200 Time (ms) 3so 50 0 8.7 8.S La ::c 8.3 c. 8. 1 7.9 so 200 Time (ms) 3S O s o o 7.7 .------------------, 7.S ::c c. 7.4 7.3 Eu 7. 2 !-, so zoo Time (ms) 3 S O 500 7.1 7 "'11"1' Lu ::c c. 6 9 1\ 6.8 50 200 Time (ms) 350 soo -Figure 3. Representative Spectrophotometric data vs. time 12

PAGE 22

anomaly, Ce oxidation is expected to be extremely slow under these experimental conditions As such, the origin of the Fig. 4 discrepancies in log p1.(Ce) values estimated from spectrophotometric and potentiometric procedures is uncertain -8 -9 potentiometric experiment I r;/-/ \ 1-1 . 2 1 potentwmetnc expenment I, :f I,' /' I I spectrophotometric average (five experiments) Y La Ce Pr Nd Pm Sm Eu Gd Tb D y Ho Er Tm Yb Lu REE Figure 4. logP1*(M) det e rminations at 1 = 0 .7 molar a nd 25C : Average of 5 s pectrophotometric experiments with standard errors plus the results of two potentiometric experiments. Fig. 5 compares the grand average of all potentiometric and spectrophotometric experiments performed at 25C and 0 7 molar ionic strength with the results of two potentiometric experiments performed at 25C and 0 1 molar ionic strength. The standard errors shown on this graph refer to the sev en experiments perfonned at an ionic strength of 0. 7 The potentiometrically derived log P 1 (M) patterns in this figure are in excellent agreement with the spectrophotometr i cally and potentiometrically derived 13

PAGE 23

patterns shown in Figure 4 and indicate that the log 1 (M) pattern across the YREE series is insensitive to changes in ionic strength. Constancy in YREE stability constant patterns over a range of ionic strengths has also been observed for YREE complexation by NT A3 -(Li and Byrne 1997) and F (Luo and Byrne 2000). -8 -9 I= 0 7 mola r (average of seven experiments) Y La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu REE Figure 5 Average determinations for all expe rim en t s at 1=0 .7 molar and 25C plu s th e results of two potenti o m etric experiments at 25C and I = 0 1 molar. The s tandard errors show n in thi s figur e refe r to th e seve n expe rim en t s performed at 0.7 molar ionic streng th The log *(M) observations obtained in this work (Figure 5) are broadly consistent with the observations ofFrolova (Figure 1) However, my YREE hydrolysis constants deviate significantly from the generally smooth patterns observed by Frolova [1]. In the present study prominent variations in slope *(M) plotted against atomic number) are observed between Gd and Lu. S i milar variations have also been observed 14

PAGE 24

for C032 complexation constants (Liu and Byrne, 1998) obtained via solvent exchange analysis and for F complexation (Luo and Byrne, 2000; Schij f and Byrne 1999) obtained via specific ion electrode and ion exchange analysis In spite of the close similarity in the ionic radii ofHo3+ and Y3+ (Shannon, 1976) the hydrolysis behav ior ofY3+ is intermediate to that observed for Sm3+ and Eu3+ (Figure 5) This is consistent with behavior expected for Y3+ and REE complexation by a highly covalent ligand (Hancock and Martell, 1989). Yttrium is a harder ion than any of the rare earths (Byrne and Lee, 1993). It is more weakly complexed than Ho in cases where REE-ligand interactions are significantly covalent (Kumok and Serebre1mikov, 1965), and is more strongly complexed than Ho3+ for bonding which is strongly electrostatic in character. Influence of ionic strength on log (M) Given the observed insensitivity of log P 1* (M) patterns to variations in ionic strength, my investigations of the influence of ionic strength on YREE hydrolysis were confined to a single element, Ho The log p 1* (M) results obtained for Ho between ionic strengths 0 04 molar and 5 .7 molar are shown graphically in Figure 6. As is the case for log 1* (M) results obtained potentiometrically at ionic strengths 0 7 and 0.1 molar the results shown in Figure 6 were obtained with a [Ho0H2+ ]/[Ho3+ ] ratio on the order of 0.11. Since a relatively high Ho3+ concentration ([Ho3+]) was used in these experiments, the ionic strength range in these measurements did not include values smaller than 0.04. The ionic strength dependence of the data shown in Figure 6 was quantitatively described using an equation identical in form to that employed in descriptions of YREE complexation by Fand NTA3 -. The best fit curve in Figure 6 is given by 15

PAGE 25

(10) where log0 1 (Ho) = -7.56 is the hydro l ysis constant appropriate to 25C and zero ionic strength -7.7 . . -.-7. 8 . 0 = . --"" '!""' c::l. bl) I 0 I . . : : -7.9 . . I : . 8.0 0 I 2 3 4 5 Ionic Strength (molar) Figure 6. lome Strength Dependence oflogp1*(Ho) at 25C. Note that the difference(:: 0.05) between logp 1 *(M) for heavy REEs at 0. 7 molar and 0 1 molar ionic strengths (Figure 5) is con sis tent with the observations shown in this figure Tem p erature de p endence ofYREE hydrolysis constants 6 The temperature dependence ofYREE hydrolysis constants was determined by conducting additional potentiometric experiments at 40C and 55C. The results, Figure 7, are the average of two potentiometric experiments at I= 0.7 molar for each temperature. Figure 7 shows that the patterns oflog 1 .. (M) values obta i ned at 40 C and 55C closely resemble the patterns obtained potentiometrically at 25C. 16

PAGE 26

Y La Ce Pr Nd Pm Sm Eu Gd Tb D y Ho Er Tm Yb Lu REE Figure 7 Potentiometric logp1*(M) results (0.7 molar) at 25C, 40C and 55C Error limit s shown in this figure indicate the range in logp1*(M) for two measurements at each temperature The log 1 (M) data in Figure 7 were fit using the van't Hoff equation. The enthalpy values obtained in these fits are given in Table 1 The average 6H0(M) value for all YREEs is 11.3 kcal mor1 In consideration of the standard errors given with each 6H0(M) estimate in Table 1, this enthalpy is in good agreement with values obtained for all elements except Ce. Enthalpies ranged from 13.6 kcal mor1 forCe to 10 kcal mor1 for Yb and Lu. The standard errors given with each 6H0(M) value in Table 2 indicate that reaction enthalpies appropriate to equation 1 are much better defined for the he avy rare earths than is the case for the lightest rare earths (La-Nd). 17

PAGE 27

REE AH+S.E. y 11.0.1 La 10.3 2.0 Ce 13.6 6.4 Pr 13.0 .7 Nd 9 9 .0 Pm Sm 10.2 1.3 Eu 10.8 1.3 Gd 12.6 1.6 Tb 12.0 1.3 Dy 11.3 0.8 Ho 11.7.7 Er 11.0 0.5 Tm 11.7.4 Yb 10.0 0.5 Lu 9.9 0.7 Table I. Enthalpy Data (kcal mon for YREE Hydrolysis. REE *(M) + S.E y -8.11 0 03 -7 .80 La -9.12 0 08 -8.81 Ce -8.65 0.10 -8.34 Pr -8.62 0.06 -8.32 Nd -8.49 0.07 -8.18 Pm Sm -8.15 0.08 -7.84 Eu -8 06 0 07 -7.76 Gd -8 .14 0 05 -7.83 Tb -7 95 0.05 -7.64 Dy -7 89 0.05 -7.59 Ho -7.87 0.05 -7.56 Er -7.82 0.05 -7.52 Tm -7.70 0.05 7.39 Yb -7.55 0.05 -7.24 Lu -7.57 0.05 -7.27 Table 2. YREE Hydrolys1s constants at 25C are shown for I= 0.7 molar S.E.) and I= 0.0 The results obtained in this work allow log P 1 (M) values to be calculated over a wide range of temperatures and ionic strengths. Table 2 provides the grand average YREE hydrolysis constants appropriate at 25C and the ionic strength of seawater (0. 7 molar). The log0 p1 *(M) values given in Table 2 were calculated using the ionic strength behavior observed for Ho, and the average log Pt *(M) results obtained for all experiments performed at 25C and 0 7 molar ionic strength. Using = 11.3 kcal mor1 as a reasonable approximation for all YREEs, hydrolysis constants can be estimated as a function of temperature and ion ic strength en wit h the following equation where T is the Kelvin temperature and log0p t (M) values are given in Table 2. 18

PAGE 28

Conclusions The log values obtained in this work are in general agreement with the previous measurements ofFrolova et al. (1966). However, observations obtained in the present study show much more structure than the Frolova et al. (1966) results The nearest neighbor log P 1 (M) relationships observed for YREEs in this work are broadly consistent with those observed in LFER characteri z ations ofYREE stability constants with a wide variety of organic ligands (Byrne and Lee, 1993; Lee and Byrne 1992) The nearest neighbor hydrolysis constants observed in this study also resemble YREE stability constant patterns observed for cot and F complexation (Liu and Byrne, 1998; Luo and Byrne, 2000; Schijfand Byrne, 1999). My hydrolysis constant observations indicate that for surface seawater at 25 C (pH= 8 2) the ratio [MOH 2 + ] / [M 3+] varies from values smaller than one for La to values on the order of five for Lu The enthalpy values obtained for the first YREE hydrolysis step are in good agreement with values predicted (Byrne et al. 1988) using the enthalpy estimation systematics ofBaes and Mesmer (1981) For the enthalpies observed in the present work (.0.H0 = 11. 3 kcal mor1 ) YREE h y drolysis constants are predicted to decrease by more than a factor of four between 25C and temperatures characteristic of the deep ocean (0C t 6C) Consequently MOH 2 + is a minor species for the YREEs in ocean waters relative to carbonate complexes, MC03 + and M(C03) 2 Hydrolysis constants have been advocated as proxies for metal ion adsorption constants appropriate to the surfaces of metal oxides and hydrous oxides and organic patiicles (Balistrieri et al., 1981; Dzombak and More l 1990; Huang and Stumm 1973; Johannessen et al. 1997; Stumm, 1987). The application of surface complexation 19

PAGE 29

models to YREEs has generally involved the use oflog p 1* (M) values obtained from linear free energy relationships (Lee and Byrne 1992) or the hydrolysis data recommended by Baes and Mesmer (1976). The log p 1* (M) data obtained in the present work provides a basis for refined YREE scavenging models (adsorptive transport) in seawater and freshwater systems. Acknowledgements This work was supported by contract OCE-95-22878 from the National Science Foundation. 20

PAGE 30

CHAPTER2 DEVELOPMENT OF NOVEL SPECTROPHOTOMETRIC PROCEDURES FOR EXAMINATION OF FERRIC ION HYDROLYSIS Introduction Iron is a vital trace nutrient, essential for all life. Due to its importance in primary production with consequent implica tions for g lobal warming, the occurrence of iron in natural waters has been studied for many yea rs It has been hypothesized that atmospheric C02 concentrations are strongly influenced by the supply of iron to the oceans. (Martin et al., 1994 ; Martin and Fitzwater, 1988; Martin et al., 1991 ) Iron is considered to be the limiting nutrient in the southern ocean and parts of the Pacific due to the weak input of iron from aeolian sources (Barber and Chavez, 1991 ; Berger and Wefer, 1991). The Iron Hypothesis" proposes that an increase in iron inputs to the ocean would boo s t primary production/ph yto plankton growth, thus removing carbon from the water column The resultingpC02 decrease in the surface ocean would allow further oceanic C02 uptake from the atmosphere. Iron bioavailability is related to sp eciation O f the total iron in the surface ocean, more than 99% is strongly complexed by iron-specific organ ics (van d e n Berg, 19 95; Wu and Luther, 1995) The remaining dissolved iron is inorganically complexed, dominantly as ferric hydroxides Ferric hydroxide species are thought to b e the principal 2 1

PAGE 31

form of inorganic iron that is bioavailable to phytoplankton (Wells et al., 1995). Consequently, iron hydrolysis has been the focus of a variety of studies in which diverse methods have been used to investigate equilibria and conjugate hydrolysis constants ) (12) where square brackets denote species concentrations. The generally poor accord between independent determinations of hydrolysis constants appropriate to Fe(OH)2+ formation can generally be attributed to analytical problems posed in work at very low Fe concentrations (e.g. adsorption, contamination, etc.) and the importance of polymeric species and precipitates at high concentrations. This work represents a feasibility study in which iron hydrolysis is studied at very low iron concentrations using solutions of sodium thiocyanate (NaSCN). Thiocyanate complexes with iron to form chemical species that absorb visible light. Solutions containing Fe(SCN)11 at sufficiently high concentrations appear reddish/burgundy to the naked eye. Ferric thiocyanate complexes have absorbance maxima near 460nm, a wavelength at which ferric hydroxide complexes (Fe(OH)113-n) do not absorb appreciably. Competitive complexation ofFe3+ by OHand SCNresults in visible absorbance spectra that are highly pH dependent. In this work the influence of pH on ferric thiocyanate spectra are interpreted in terms of the equilibria depicted in equation 12 Although the present study solely involved observations with a 1 Ocm path length cell, the procedures developed in this work can be extended to long pathlength analyses (Byrne et al., 2000b) whereby iron concentrations can be reduced by more than an order of magnitude 22

PAGE 32

Consequently, future studies using this methodology can be extended to higher pH without interference from precipitation and polynuclear species formation This will allow greater access to conditions whereby formation of species such as Fe(OH)2 +can be more accurately examined. Theory Under the conditions of this work the total iron concentration is the sum concentration of five species where square brackets [] denote molar concentration of each species. Using equation (12) the concentration ofFeOH2+ and Fe(OH)/ can be written as follows (14a) (14b) Likewise, the concentrations ofthe ferric thiocyanate complexes can be expressed in terms of ferric thiocyanate formation constants (s cNBn) as follows where (FeSCN2+ ) seN Bn -[Fe3+ JscN-]" (15a) (15b) (16) Substituting equations 14a), 14b ), 15a) and 15b) into equation 13 we obtain an expression relating the free ferric ion concentration to the total iron, pH and free thiocyanate ion concentration. 23

PAGE 33

Equation (17) was used in potentiometric experiments to d eterm ine and in spectrophotometric experime nts to determine . Potentiometric Experiments Potentiometric experiments were used to relate absolute millivolt readings of a platinum electrode to free ferric ion concentrations in solutions containing both ferric and ferrous iron. The platinum electrode responds to the ratio between free ferric and ferrous ion concentrations The equation describing this relationship at 25C is (18) where Eo is a constant, which incorporates the standard potential for the reaction Fe3 + + e B Fe2 + as well as activity coefficients for both Fe3+ and Fe2+ ions. E1.j is the liquid junction potential of the refer ence electrode in solution. Since this work was performed at constant temperature and ionic strength, and the medium composition was approximately constant the liqu id junction potential is expected to remain constant. Noting that the ratio of total ferric and ferrous concentrations in these experiments also remains constant equations (17) and (18) can be combined to yield an expression that directly relates observed oxidation/reduction potentials, thiocyanate concentrations, and pH: where is defined as 24

PAGE 34

(20) Equation (19) can also be written in the following form For determinations of the hydrogen ion concentration was 0 05 molar whereupon the *[H+ r2 term is negligible. Using 1 = -2 754 (Byrne et al. 2000a) equation (21) can be written as ( 1 oCE' o-E ) / s 9 .1 6 -1 035 'fscN] -I= A + A rscN] J1! SCN 1-'1 S C N 1-'2 (22) Due to the high iron concentration in these experiments, a significant fraction of the total thiocyanate i s complexed as Fe(SCN)0 Free thiocyanate concentrations were calculated with the following equation (23) Equation s (22) and (23) were us e d iteratively to calculate 1 and from observations of E as a function of [SCN-] in potentiometric experiments at const a nt pH The first measured electrode potential in each experiment was obtained with [SCN-] = 0 whereupon equation 19 allows calculation of as = E + 59 16log(1.035) Spectrophotometric Experiments Absorbance is related to species concentration and pathlength tlu ough Beer s Jaw (24) 25

PAGE 35

where Ei is the molar absorptivity of species "i" in solution, ci is the molar concentration of solution species "i" and e is spectrophotometric cell pathlength in centimeters. In th ese experiments, the Beer's law equation can then be written as follows (25) where Eo is the molar absorptivity ofFe3+ in the absence of SeN, E1 i s the molar absorptivity ofthe first ferric thiocyanate complex and e 2 is the molar absorptivity of the second ferric thiocyanate complex. Combining equations (17) and (25) yields an expression relating A, [Fem]T, thiocyanate concentration and pH A Eo +E) seN P1 [seN-]+ seN P z [seNY [ Fem ]TR = r + P;[H+ J-1 +P;[H+ r (26) Equation (25) can b e rearranged in a manner that allows direct determination of the coeffic ient s Eo, EtscNPt and EtscNP2 at low pH cp;[H+ r-0): All of the terms on the left of equation (27) are dir ect ly observed or mode l ed from potentiometric experiments. The coe f ficients (EiPn) on the right hand s ide of equation (27) are then obtained in quadratic fits of th e left side ofthe equation (Y t) as a function of [SCNl Having obtained estimates for Eo, EtscNPt and EtscNP 2 scNP t and scNP 2 equation (26) can be rearranged to yield an equation that allows estimation ofP 1 and P2 in experiments at high pH. 26

PAGE 36

(28) Spectrophotometric experiments were performed at constant thiocyanate concentration such that [FeSeN2+ ] / [Fe3+ ] = 1 Thus the term (1+scN [seNJ+scN [seNY) was approximately equal to 2 The left side of equation (28) (in which all parameters are observed or calculated from previous measurements) is seen to be a simple quadratic function of [H +rl. Equation (28) was used for determination 1 *and in experiments conducted between pH 2 7 and 4.5. Methods and Materials Materials The Nael04 solutions for all potentiometric and spectrophotometric experiments were prepared from anhydrous Nae104 obtained from Sigma, (St. Louis MO). A 6M Nael04 stock solution was prepared in the same manner as in the previous chapter (the stock solution was stored overnight at pH between 10 and 11, and w a s filtered through 0.4J..Lm polycarbonate membrane filters and subsequently acidified to pH 4) Sodium thiocyanate (99.98 % ) and anhydrous ferric chloride hexahydrate Feeb.6H20 (99 99%) were purchased from Fisher Scientific (Pittsburgh PA) Ferrous chloride hexahydrate, Fee}z.6H20 (99.99%) and perchloric acid were obtained from Aldrich (Milwaukee WI). 27

PAGE 37

Methods Potentiometry The ferric thiocyanat e comple x ation constants 1 and w e re determined potentiometrically in a series of 10 experiments. 150 mls of solution were placed in a jacketed beaker maintained at 25C 0 1 using a Lauda 6M thermocooler. The solution was composed of 5x104M Fe111, 10 -3M Feu, O .OSM HC104 and 0 65M NaCl04 (total ionic strength was 0.70 M). An Orion/Ross-type combination pH/reference electrode (Orion No. 8102BN) and a platinum metrohm electrode(#6-0351 100) in combination with an Orion 800500 reference electrode were equilibrated while the solution was sparged with ultra-high-purity nitrogen and st i rred constantly with a Teflon-coated stirring bar. The pH vs. reference and platinum vs. reference electrode combinations were calibrated on free ion concentration scales ([W] and [Fe3+]) prior to each experiment. Each electrode exhibited Nemstian behavior Sodium thiocyanate was added with an Eppendorf pipette, providin g solution concentrations between 0 002 and 0.05 molar The s odium thiocyanate titrant was compos e d of 0.65M NaSCN and O.OSM HC104 Equilibrat ion occurred within 30 seconds. This was evident from ancillary experiment s in which electrode pot e ntials wer e observed to vary by less than 0 1m v between equilibration periods of 6 0 seconds and 10 minutes. Solution pH varied by less than 0.002 pH units throughout each experiment. The electrode potentials generated from these experiments were used in Equation (22) to obtain I and 28

PAGE 38

Methods Spectrophotometry Two series of spectrophotometric experiments were conducted. In the first series pH and iron concentrations were constant and thiocyanate was varied between 0.004 and 0 048 M. Absorbance was observed with changing thiocyanate concentrations by adding a titrant solution (0.65 M NaSCN and 0 05 M HC10 4 ) to a background/baseline solution of0.65 M NaCl04 plus 0.05 M HC10 4 with an initial Fe111 concentration equal to 2x105 M. These experiments were conducted in an open-top, 1 Ocm pathlength, quartz spectrophotometric cell. The titrant solution was added using a Gilmont 2ml micro syringe with a 24-gauge Teflon syringe needle The solution mixture in the spectrophotometric cell was stirred thoroughly with a glass rod and absorbance was measured after each thiocyanate addition. Temperature was maintained at 25.0C 0.1 using a Neslab RTE-211 thermo-circulator. The absorbance results from this series of experiments plott e d as a function of [SCN-] were fit using equation (27) to obtain ErscNP 1 and E1scNP2 The constant Eo was experimenta ll y determined to be z ero. These data are shown in Appendix D. In the second series of spectrophotometric experiments both iron and thiocyanate concentrations were constant and pH was varied. The background/baseline solution in this case consisted of0.698 M NaC10 4 0.002 M HCl04 and 5.49x10 3 M NaSCN. Iron was then added to this baseline solution to yield a final concentration of2xl0-5 M Fe1 ll. At a NaSCN concentration equal to 5.5 x 1 o 3M, the ferric thiocyanate complex is approximately 50% ofthe total iron ([FeSCN2+ ] / [Fe 3+] = 1). Solution pH was lo wer e d in one experiment by adding 0.585 M HCl04 and was raised in subsequent experiments by adding a titrant so luti on of0.7M NaHC03 These titrants were added with a Gilmont 2ml 29

PAGE 39

micro-syringe connected to a 24-ga uge Teflon syringe n eed l e The so l u tion was mec h anically stirred after each addi t ion of acid or base using fol ded Teflon t ubing rotated with a CAFRAMO type RZR l (Wiarton, Ontario Canada) external stirrer. The same combination reference/pH electrode u sed in the poten t iometric experiments was u sed here for measurements of solution pH The electrode and the Teflon tubing were removed briefly during absorbance measurements. An examp l e of the a b so r bance s p ectra generated in one of these experiments is shown in Figure 8. The first ferric-thiocyanate comp l ex peak at 460nm is prominent at the lower pH and as pH increases the FeOH2+ peak becomes significant at 293nm There are two very distinct isosbestic points All spec t ra pass thro u g h t h ese points wit h the exception of that obtained at t h e highest pH, 4.52 Gl u = Cll .0 "" Q < 0.12 pH 0 .10 2 7 2 89 -3.01 0 08 3.20 3.55 0 06 -3.81 4 05 4 20 0 04 4 30 4 52 0.02 0 00 250 300 350 400 450 500 550 600 650 Wave l e ngth Figure 8 Absorbance s pectra showing influence of pH on FeSCN2 + and FeOH2 + absorban ce 30 7 00

PAGE 40

An additional experiment was carried out under identical conditions as just cited but with no added thiocyanate The results of this experiment (Figure 9) indicate that FeOH 2+ and Fe3+ absorbance contributions are negligible at 460nm. = = 0.1 0 .08 -e 0 .06 .c < 0 .04 0.02 250 300 350 400 450 500 550 600 Wavelength Figure 9. Absorbance spectra of Fe in 0 7 NaC104 between pH 2. 7 and 4 Results and Discussion 650 700 Linear and nonlinear data regressions were performed using SigmaPlot. Dilution factors, although minute, were taken into account in all analyses. Figure I 0 shows the data from the eight potentiometric determinations of and Table 3 shows ferric thiocyanate stability constant results with standard errors. 31

PAGE 41

200 -. z u 180 rn ........ ..-._ II) ("') 0 ....( 1 60 -o o; ;g_ o L:l 140 + e o -..._, 120 0 0 00 0 0 1 0.02 0.03 0 04 0.05 [NaSCN] FigurelO. P otentiometric data u se d for determination of sCNPt and scNP 2 SCN(31 SCNr32 1 119.39 0.48 1511.57 18.22 2 118.97 0.93 1505 63 35.48 3 118.38 0.87 1 5 34 60 33.15 4 121.25 0.52 1457.58 19.66 5 119.70 0.70 1488.66 26 92 6 1 20.30 0.81 1473 98 30 69 7 117 50 0.69 1494 86 26 .13 8 1 21.29 0 65 1 581.71 24 .65 Average 119. 60 0 .71 1 50 6.07 26.83 Table 3 Results from potentiom e tric experm1ents Using the potentiometrically derive d average values for scNP 1 and scNP2 (Table 3), equatio n (27) was used in determinations of g lscNP 1 and g2SCNP2 Figure 11 shows the data used for these determinations, and Table 4 shows results ob t ained in determ i nat i ons of the coefficien t s g1scNB1 and tscNP2 32

PAGE 42

70000 60000 50000 40000 30000 20000 10000 0 0 .00 0 .01 0 .02 0.03 [Na SCN ] 0 .04 F i gure 11. Data used for determination of E 1scNP1 and E2scNP2 E xp CISCNJ31 C2SCNJ32 1 5.63.+05 l.O. E+03 1 .336.+07 2 65.+04 2 5.69.+05 3.0.E +03 1.335.+07 7.75 .+04 3 5.62 +05 7.0.E+02 1.338.+07 1.81. + 04 5.85.E+05 6.7.E+02 1.317.+07 1.75. + 04 5 5.87 +05 9.0.E+02 1.337 +07 2 35 .+04 Average 5.73 +05 1.3.E+03 1.333.+07 3.26.+04 Table 4 Results of five expenmental detemunatwns of E1scNP1 and E2scNP2 0 .05 Table 5 shows a comparison of the summary (Table 3) scNP n resu l ts with critically reviewed recommended ferric thiocyana t e s t abi l ity constants (Bahta et al., 1 997; Beck, 1977). The results obtained (Table 3) in t his work are i n good overall agreement with previous published results obtained of a range of ionic strengths. 33

PAGE 43

Source I logscNl3t logscNl32 Macdonald et al. 1.8 2.09 0.004 3 84 0 03 Portanova et al. 1.0 2.10 0.007 3.14 0 .03 Mieling & Pardue 1.0 2.11 0.01 2.75 0.02 O z utsumi et al. 1.0 2.11 0.01(3cr) 3.34 0.02(3cr) Laurence 0 5 2.14 0 .005 3.45 0 02 Lister & Rivington 1.2 2.11 0.003 3.29 0 .03 This work 0 7 2.08 0 003 3.18.01 Table 5. Summary of critically se le cted const ants Finally th e potentiometrically derived scNl3t and scNl32 data along with spectrophotometrically derived t::1scNl3t and t::2 scNP2, data were used in eq u a tion (28) to obtain plots of Y 2 vs. [H +rt. The results of four sets of experiments are shown in Figur es 12a and 12b Linear behavior was observed up to a pH of 4. 34

PAGE 44

0.8 ,------------------Exp.l 0 6 0.4 0.2 0.0 0 100 400 500 25 20 Exp. 2 .. 15 10 5 ------0 0 5000 10000 15000 20000 25000 [H+r Figure 12a. Data used for determination 1 35

PAGE 45

15 Exp.3 10 > 5 0 0 2000 4000 6000 8000 10000 12000 14000 [ H+ r l 40 Exp.4 "" 30 .. 0 ... >cM20 0 10000 20000 30000 40000 50000 [H+rl Figure 12b. Data used for d e temunation of P 1 36

PAGE 46

Exp Exp 1-1 -2 83 0 006 3-1 -2.87 0.004 1-2 2 82 0 006 3-2 -2.88 0 006 1-3 2 82 0.003 3-3 -2.87 0.002 1-4 -2 82 0.005 4-1 -2.82 0.003 1-5 -2.79 0.009 4-2 -2.86 0.002 2-1 -2 90 0 005 4 3 2 .85 0 .001 2-2 2.86 0 003 4-4 -2 88 0 004 2-3 -2.88 0 002 4-5 -2 87 0.001 2-4 -2 89 0 007 Average -2.85 0.004 Tabl e 6. Summary of logp 1 r es ults using e qu a tion ( 27) and 3. 7 Linear regressions ofY2 v s [H+rl (for [Wr1 s; 50 0 0 ) produ c ed the res ult s shown in table 6 Equation (27) indicates that the intercepts of th e Y 2 vs. [H + T1 plots should be zero. Small (non z ero) intercepts were observed, po ss ibl y due to sm all differences between calcu l a t ed v s actual Felll concentrations in solution. Tabl e 7 s h o ws the 1* results obtained when the Felll concentrations u sed in equation (27) are adjusted to give a zero intercept. These adjustments were smaller than 5% and ge n e rally caus e d only minor changes in p 1 . Exp Exp l og Pt 1-1 -2.82 3-1 -2.93 1-2 -2 8 1 3-2 -2 97 1 3 -2.81 3-3 -2 93 1-4 -2 .81 4-1 -2.84 1-5 -2.79 4-2 -2.84 2-1 -2 94 4-3 -2.83 2-2 -2 87 4-4 -2 87 2-3 -2.92 4-5 -2.86 2-4 -2 88 Ave r age 2 .87 Table7. S umma ry of logP1 r es ult s u s m g e quation ( 2 7) wi th [Fe111]T adju s ted to g ive a z ero int e r c ept 37

PAGE 47

The negative deviations from linearity seen in Figures 12a and 12b may be attributable to scattering from precipitated (colloidal) iron. If so, in the future this problem could be mitigated through the use of long pathlength cells and lower iron concentrations. The negative deviations could also be attributed to the existence of a m i xed-ligand species FeSCNOH+, which could absorb light at wavelengths over the same range as ferric thiocyanate Previous investigations of the ferric thiocyanate system (Yalman, 1961) provided evidence of absorbing mixed ligand halide species (FeSCNF + and FeSCNBr+). To the best of my knowledge, no studies have been carried out investigating the absorbance characteristics of the mixed ligand species FeSCNOH + Future work in this area will be required to calibrate this method for iron hydrolysis determinations. Table 8 shows previously published logp 1"' results at 25C and 0. 7 molal ionic strength The results are shown on the molal concentration scale In addition to possible problems caused by formation of colloids both the negative deviations from linearity (Figures 12a and 12b) and the smaller-than expected results for logp1"' may be attributable to absorbing mixed ligand species such as FeSCN(OHt that are not taken into account in Equation (27). The importance o f this possibility can be evaluated by conducting experiments over a range of thiocyanate concentrations 38

PAGE 48

Source method Milburn and Vosburgh (1955) -2.73 Spectrophotometry Byrne and Kester (1976) -2.77 Potentiometry Byrne and Kester (1978) -2.72 Spectrophotometry Soli and Byrne (1996) -2 .75 Potentiometry Byrne and Thompson (1997) -2 .75 Potentiometry Baes and Mesmer (1976) -2 68 Assessment of literature data Millero et al. (1995) -2.62 Reanalysis of literature data Byrne et al. (2000) 2.74 Potentiometry Table 8. Previous determinations of logl31 The fair agreement between previous results and the 1* results of this work indicates that the procedure developed in this work for iron hydrolysis investigations is promising. These procedures a l so have applications in determinations of weak Fe3+ ion pairing constants. Figure 13 shows Fe111 absorbance spectra obtained in perchlorate, nitrate and chloride solutions Each spectrum was obtained using 6x 1o -5M Fe111, 0.02M HC104 and 0 00549M NaSCN in each background e l ec trol yte (0.68M NaCI, 0.68M NaN03 and 0 68M NaC104 ) One experiment performed with no thiocyanate, was used to demonstrate that ferric chloride complexes are not significantly absorbing species at 460nm. 39

PAGE 49

650 700 Figure 13. FeSCN absorbance i n NaCI04 NaN03 and NaCI The P1 results obtained in these simple experiments demonstrate that NO; complexes Fern appreciably compared to CI04and that Cr complexes Fern substantially under these conditions These observations also indicate that these procedures might be particularly useful in investigations of Fern association with weakly complexing electrolytes. 40

PAGE 50

REFERENCES Baes C. F. and Mesmer R. E (1976) The Hydrolysis of Cations. John Wiley & Sons, Inc. Baht a A ., et al. (1997) Critical survey of stability constants of comple xes of thiocyanat e ion (Technical Report) In Critical Evaluation of Stability Constants of M e tal Complexes in Solution Vol. 69, pp 1489-1548 Pergamon Press. Balistrieri L., et al. (1981) Scavenging residence times oftrace metals and surface chemistry of s inking particles in the deep ocean Deep-Sea Research 28A, 1 01121. Barber R T. and Chavez F P. (1991) Regulation of primary productivity rate in equatorial Pacific Limnology and Oceanography 36(8), 1803-1815 Beck M. T. (1977) Introductory chapter to Series A of the critical evaluations of equilibrium constants in solution. In Critical Evaluation of Stability Constants of Metal Complexes in Solution Vol. 49, pp 129-135 Pergamon Press. Berger W H and Wefe r G. (1991) Productivity of the glacial ocean: Discussion of the iron hypothesis. Limnology and Oceanograph y 36(8), 1899-1918 Breland J A. and Byrne R H. (1993) Spectrophotometric procedures for detennination of sea-water alkalinity using bromocresol green. Deep -Sea R esea rch I 40(3), 629641. Byrne R H (1987) Standardization of Standard Buffers by Visible Spectrometry Analytical Chemistry 59, 14 79-1481. 41

PAGE 51

Byrne R. H., et al. (1988) The influence of temperature and pH on trace metal speciation in seawater. Marine Chemist1y 25, 163-181. Byrne R. H. and Lee J. H. (1993) Comparative yttrium and rare earth element chemistries in seawater. Marine Chemistry 44, 121-130 Byrne R. H., et al. (2000a) Iron hydrolysis and solubility revisited: observations and comments on iron hydrolysis characterizations. Marine Chemist1y 70, 23-35. Byrne R. H., et al. (2000b) Construction of a compact spectrofluorometer/spectrophotometer system using a flexible liquid core waveguide. Talanta 50, 1307-1312. Clayton T. D and Byrne R H (1993) Spectrophotometric seawater pH measurements : Total hydrogen ion concentration scale calibration of m-cresol purple and at-sea results. Deep-Sea Reseasrch 140(10), 2115-2129 Dzombak D. A and Morel F. M. M. (1990) Surface Complexation Modeling. Wiley. Frolova U.K., et al. (1966) Gidroliz Ionov Redkozemelnykh Elementov i Ittriya v Vodnykh Rastvorakh. Izvestiya Vysshikh Uchebnykh Zavedenii SSSR 9(2), 176-179. Hancock R. D. and Martell A. E. (1989) Ligand design for selective complexation of metal ions in aqueous solution. Chemical Reviews 89, 1875-1914 Hamed H. S. and Owen B. B. (1958) The Physical Chemistry of Electrolytic Solutions. Reinhold. Hedstrom B. 0. A. (1952) Studies on the hydrolysis of metal ions Vll. The hydrolysis of the iron (IID ion, fe3+. Arkiv For Kemi. 6(1), 1-16. 42

PAGE 52

Huang C. and Stumm W (1973) Specific adsorption of cations on hydrous y-Al203. Journal of Colloid and Interface Science 43, 409-420. Johannessen K. H., et al. (1997) Rare earth elements as geochemical tracers of regional groundwater mixing Geochimica et Cosmochimica Acta 61(17), 3605-3618. Kumok V. N. and Serebrennikov V V (1965) Stability of rare-earth complexes Russian Journal of Inorganic Chemistry 1 0(9), 1095-1098 Lee J. H and Byrne R. H (1992) Examination of comparative rare earth element complexation b ehavior using linear free-energy relationships Geochimica et Cosmochimica Acta 56, 112 7-1137. LiB. and Byrne R. H. (1997) Ionic strenth dependence of rare earth-NTA stability constants at 25 C. Aquatic Geochemistry 3, 99-115. Liu X. and Byrne R. H. (1998) Comprehensive investigation of yttrium and rare earth element complexation by carbonate ions using ICP-mass spectrometry. Journal of Solution Chemistry 27(9), 803-815. Luo Y. and Byrne R. H (2000) The ionic strength dependence of rare earth and yttrium fluoride complexation at 25C Journal of Solution ChemistJ)' 29(11) Martin J. H., et al. (1994) Testing the iron hypothesis in ecosystems of the equatorial Pacific Ocean. Nature 371, 123-129. Martin J. H. and Fitzwater S E. (1988) Iron deficiency lim its phytoplankton growth in the northeast Pacific subarctic Nature 331, 341-343 Martin J. H., et al. ( 1991) Iron Limitation? Limnology and Oceanography 36(8), 17931802. 43

PAGE 53

Milburn R.N. and Vosburgh W. C. (1955) A spectrophotometric study of the hydrolysis of iron (III) ion. II. Polynuclear species. Journal of the American Chemical Society 77, 1352-1357. Robert-Baldo G. L., et al. (1985) Spectrophotometric Determination of Seawater pH Using Phenol Red. Analytical Chemistry 57, 2564-2567. Schijf J and Byrne R. H. (1999) Determination of stability constants for the monoand difluoro-complexes ofY and the REE, using a cation-exchange resin and ICP MS. Polyhedron 18, 2839-2844. Shannon R. D. ( 197 6) Revised effective ionic radii and systematic studies of interatomic distances in halides and chalcogenides. Acta Crystallographica A 32, 751-767. Stumm W. (1987) Aquatic Surface Chemistry. Wiley. van den Berg C. M .G. (1995) Evidence for organic complexation of iron in seawater. Marine Chemistry 50, 139-157. Wells M. L., et al. (1995) Iron chemistry in seawater and its relationship to phytoplankton: A workshop report. Marine Chemistry 48, 157-182. Wu J. and Luther G. W. (1995) Complexation of Fe (III) by natural organic ligands in the Northwest Atlantic Ocean by a competitive ligand equilibration method and kinetic approach. Marine Chemistry 50(1-4), 159-177. Yalman R. G (1961) Stability ofthe mixed complex FeSCNF+. Journal of the American Chemical Society 83, 4142-4146. 44

PAGE 54

APPENDICES 45

PAGE 55

0\ Appendix A: Previous logf31 data (from Figure 1) Reference : I y La Ce Pr Frolova e t al. 1966 0.3 -7.87 -8. 59 -8.08 Guillaumont e t al. 1971 0 1 7 03 K.ragten and Decnop-Weever 1978-1987 1.0 -7 .99 -7 49 Amaya e t al. 1973 3.0 -9 36 Burkov et al. 1975 0.3 Burkov et al. 198 2 3 0 -9 64 -8 8 4 Biedermann and Ciavatta 1961 3.0 -9 38 C ac ec i and Choppin 1983 0. 7 Din Ngo and Burkov 1974 3 0 Davydo v and Voronik 1983 0 5 -7 .16 Ivano v-Emin e t al. 1970 Lopez-Gonzale z et a l. 1997 -7 00 Moeller 1946 0 0 -8 64 9 94 -8.94 -8 94 Nair et al. 1982 1.0 Schmidt et al. 197 8 0 0 Tobia s and Garrett 1958 3 0 -7.78 Usherenko and Skorik 1972 0 1 -7 74 -8 76 Wheelwright e t al. 1953 0 1 -8 0 3 Baes and Mesmer 1976 0 0 -7 70 8 50 8 30 8 .10 Le e a nd Byrne 199 1 0.0 -8 66 8.4 1 8 27 Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu -7 .96 -7.87 -7 84 -7 88 7.69 -7 63 -7.57 -7.52 -7.48 -7.45 -7.43 -6 .63 -4.03 -6 73 -5.13 -3.93 -7 .49 -6.89 -6 69 -5.69 -7 09 7 48 -8.51 8 68 -7 88 -6.73 8.48 -8.40 -8.94 -8.84 -8.84 -8. 34 -7 .51 -8 60 -7 .78 -7.73 -7 97 -7.71 -7 36 -8 00 -7 90 7 80 -8 00 -7 90 8 00 -8 00 -7 90 -7 70 7 70 -7 60 -8 .16 -8.06 -7 96 7.90 7.93 7 86 -7 8 1 7 78 -7 73 7 6 8 7 60 -7.59

PAGE 56

Appendix B : Potentiometric data at 55C 40C and 25 C REE 55C (1) 55C (2) 40C (1) 40C (2) 25C (1) 25C (2) y -7.479 7.306 -7.722 -7.660 -8.106 -8.149 La -8 .59 4 -8 .427 -8.749 -8.570 -9.179 -9 .211 Ce -8 165 -6.969 -7.941 -7.927 -8.405 8.549 Pr -8.004 -7.595 -8.143 -7.982 -8.658 -8.679 Nd -7.767 -7.635 -7.915 -7.737 -8.314 8.399 Pm Sm -7.515 -7.389 -7.716 -7.615 -8.140 -8.124 E u 7.415 -7.260 -7 588 -7.543 -8.039 -8.073 Gd -7.382 -7.233 -7.612 -7.509 8.169 -8.131 Tb -7 .279 -7. 105 -7.495 -7.459 -8.021 -7.961 Dy -7 .22 1 -7.10 5 -7 .525 -7.461 -7.945 -7.895 Ho -7.173 7.079 -7.517 -7.452 -7.942 -7.873 Er -7.118 -7.105 -7.455 -7 .399 -7 .866 -7 .82 4 Tm -6.971 -6.919 -7.307 -7.296 -7.754 -7.707 Yb -6.890 -6.886 -7.195 -7.151 -7.588 -7.527 Lu -6.965 -6.98 4 7.234 -7.211 -7 668 -7.599 Appendix C: Spectrophotometric data at 25C and I = 0.7 YREE 1 2 3 4 5 y -7.943 -8.03 5 -8.152 -8.177 8.194 La -8.787 -8.931 -9.073 -9.429 -9 195 Ce -8.328 -8.567 -8.769 -9.039 -8.886 Pr 8.355 -8 .497 -8.586 -8 .813 -8.775 Nd -8.261 -8 .424 -8.564 -8.793 8.679 Pm Sm 8.034 -7.741 -8.28 4 -8.385 -8.33 4 E u -7.818 -7.858 -8 .191 8.287 -8.178 Gd -7.981 -7.92 1 8.243 8.306 -8.199 Tb -7 784 -7.759 -8.035 -8.059 -8.003 Dy 7.730 -7.713 -7 .983 -7 .998 7.992 Ho -7.678 7.658 -7.953 -8.006 -7.9 47 Er -7.663 -7.621 -7 914 -7.922 -7.955 Tm -7.550 -7 .489 -7.779 -7.774 -7.820 Yb -7.395 -7 379 7.687 -7.668 -7.593 Lu -7.408 -7 333 -7.666 -7.666 -7.660 47

PAGE 57

Appendix D: Absorbance and [NaSCN] data used to calculate and at constant pH Exp.1 Exp. 2 Exp.3 Exp. 4 Exp.5 [Fe111h [NaSCN] Abs Abs Abs Abs Abs 0.00002 0 0.000 0.000 0 000 0 000 0 000 2E-05 0.004 0.326 0.328 0 323 0.325 0.329 2E-05 0.0045 0.354 0.355 0.351 0.356 0 361 2E-05 0 00499 0.380 0.383 0 378 0 384 0.389 2E-05 0.00599 0.432 0.435 0.428 0.435 0.436 2E-05 0 00699 0.476 0.480 0.474 0.481 0.485 2E-05 0 00799 0.516 0.520 0.513 0 523 0 528 2E-05 0.01199 0.640 0.649 0 643 0 651 0 659 2E-05 0.01599 0.736 0.743 0.737 0.746 0 753 1.9E-05 0.01999 0.809 0.815 0.807 0.818 0 822 1.9E-05 0.024 0.866 0.871 0.863 0.875 0 880 1.9E-05 0 028 0.912 0.916 0.909 0.922 0.928 1.9E-05 0.032 0.949 0 956 0.949 0.959 0.966 1.9E-05 0.036 0.983 0.988 0 .981 0.991 1 000 1.9E-05 0 03999 1.008 1 008 1.009 1.019 1.027 1.9E-05 0.04399 1.033 1.040 1 033 1.041 1.051 1.9E-05 0.04798 1.054 1.060 1.053 1.060 1.071 48

PAGE 58

A pp e ndi x E: Thioc yana t e abs orb anc e data at 460 nm f ro m Ex p er ime nt 1 1 1 1 -2 1 3 1-4 1-5 1-6 A b s p H Ab s pH A b s p H Abs pH A bs p H A b s p H 0.074 2.70 0.075 2 .7 0 0 075 2 70 0.075 2.70 0 075 2.70 0 074 2.70 0.079 2 59 0.08 1 2 60 0.08 2.59 0.081 2.59 0 08 2.59 0.08 2.60 0.083 2.5 1 0.084 2.51 0 084 2.5 1 0 08 5 2 .51 0 084 2 .51 0 084 2 5 1 0 086 2.43 0.087 2.44 0.087 2.4 4 0.087 2.43 0 087 2.44 0 08 7 2.44 0 088 2.37 0 089 2.38 0 089 2.38 0 089 2.37 0.089 2.38 0 .088 2 38 0.09 2.27 0 092 2 28 0 09 1 2 28 0.093 2.27 0 092 2 28 0.092 2 28 0 093 2.19 0 093 2.19 0.094 2 20 0 095 2.19 0 093 2.19 0.094 2.20 0 094 2 .13 0 096 2.13 0 096 2.13 0 096 2. 1 2 0 095 2.13 0 095 2.13 0 097 1.99 0 098 1.99 0.097 2.00 0 098 1.98 0 098 1.99 0 097 2.00 A ppendi x F : Thioc y an a t e ab s orb anc e dat a at 4 6 0 nm f rom E x p e rim e nt 2 2-1 2-2 2 3 2 -4 2-5 A b s pH A bs pH A b s p H A bs p H A b s p H 0.079 2.70 0.08 2.70 0.078 2.70 0.078 2.70 0 079 2.70 0 078 2 .71 0.078 2 .75 0.073 2 79 0.073 2 79 0.074 2 .79 0.075 2 76 0.074 2.80 0.067 2 89 0.066 2 89 0.067 2.90 0.057 3 04 0 066 2 .93 0.057 3.04 0.058 3.03 0 058 3 04 0 045 3 23 0 061 3 .01 0 05 3.15 0 .051 3 14 0.051 3.15 0 .031 3 29 0.047 3.22 0 044 3.24 0 045 3 23 0 045 3.24 0 015 3.94 0 043 3.29 0 .0 29 3.52 0.035 3.41 0 034 3.43 0.014 4 .05 0 034 3.44 0 019 3.77 0.026 3 60 0 025 3 62 0 .023 3.68 0.008 4 29 0.0 1 2 4 .0 4 0.0 1 4 1 6 0 013 4 .01 4 9

PAGE 59

Appendix G: Thiocyanate absorbance data at 460 nm from Experiment 3 3-1 3-2 3-3 Abs pH Abs pH Abs pH 0.077 2.70 0.078 2.70 0 .077 2.70 0.072 2.79 0.072 2.79 0.072 2.79 0.066 2.89 0.066 2.89 0.066 2.89 0.057 3.03 0.057 3.02 0.057 3 .0 3 0.050 3.14 0.051 3.13 0.051 3.13 0.044 3.23 0.046 3.22 0 .045 3 22 0.034 3.41 0.036 3.39 0.035 3.40 0.025 3.60 0.028 3.55 0.026 3.59 0.011 4 .06 0 .0 14 3.97 0 .0 12 4.01 Appendix H: Thiocyanate absorbance data at 460 nm from Experiment 4 4-1 4-2 4-3 4-4 4-5 Abs pH Abs pH Abs pH Abs pH Abs pH 0.072 2.70 0.081 2.70 0.080 2.70 0.081 2.70 0.081 2.70 0 .061 2 89 0.069 2.88 0.069 2.89 0.070 2.87 0.069 2.88 0 042 3.20 0.060 3 .0 1 0.060 3.01 0.062 2.99 0 .0 60 3 02 0.025 3.54 0.047 3.21 0.047 3.20 0.050 3.17 0.048 3.21 0.015 3.81 0.027 3.56 0.027 3 .55 0 .027 3.57 0.027 3 .56 0.010 4 .09 0.017 3 .83 0.017 3.81 0.017 3.85 0.016 3.86 0 009 4.17 0.011 4 .07 0.011 4.05 0.010 4 .15 0.010 4.13 0.008 4 .26 0 007 4.30 0 .008 4.20 0.007 4.37 0.007 4.29 0 006 4 37 0 .006 4.38 0.007 4.30 0.006 4.48 0.006 4.39 0.005 4 54 0.005 4.48 0.005 4.52 0 005 4.53 0.005 4.58 50


printinsert_linkshareget_appmore_horiz

Download Options

close
No images are available for this item.
Cite this item close

APA

Cras ut cursus ante, a fringilla nunc. Mauris lorem nunc, cursus sit amet enim ac, vehicula vestibulum mi. Mauris viverra nisl vel enim faucibus porta. Praesent sit amet ornare diam, non finibus nulla.

MLA

Cras efficitur magna et sapien varius, luctus ullamcorper dolor convallis. Orci varius natoque penatibus et magnis dis parturient montes, nascetur ridiculus mus. Fusce sit amet justo ut erat laoreet congue sed a ante.

CHICAGO

Phasellus ornare in augue eu imperdiet. Donec malesuada sapien ante, at vehicula orci tempor molestie. Proin vitae urna elit. Pellentesque vitae nisi et diam euismod malesuada aliquet non erat.

WIKIPEDIA

Nunc fringilla dolor ut dictum placerat. Proin ac neque rutrum, consectetur ligula id, laoreet ligula. Nulla lorem massa, consectetur vitae consequat in, lobortis at dolor. Nunc sed leo odio.