Lanthanide Ternary Complexes Relevant to the Nuclear Fuel Cycle

Transcription

1 Lanthanide Ternary Complexes Relevant to the Nuclear Fuel Cycle A Dissertation submitted to The University of Manchester for the degree of Master of Science by Research in the Faculty of Engineering and Physical Sciences 2012 Georgina Roughley School of Chemistry 1

2 Contents List of Figures 4 List of Tables 5 List of Equations 6 Abstract 7 Declaration 8 Copyright Statement 9 Acknowledgements 10 List of Abbreviations 11 Chapter 1: Introduction 1.1 The Nuclear Fuel Cycle The TALSPEAK Process The f elements The Lanthanides Lanthanide Hydrolysis Lanthanide Carbonate Complexes Lanthanum Europium Lutetium EDTA Potentiometry Speciation Diagrams Ultraviolet-Visible Spectroscopy Aims 29 Chapter 2: Experimental 2.1 Potentiometric Titration Solution Preparation Lanthanide(III) Nitrate Solutions Ligand Solutions Lanthanide(III) : EDTA Binary Complex Solutions Lanthanide(III) : EDTA : Na 2 CO 3 Ternary Complex Solutions Carbonate free NaOH 31 2

3 2.2 Instrumentation Potentiometry UV-Vis Spectroscopy 32 Chapter 3: Results/Discussion Binary Complexes La(III) : EDTA (1:1) Eu(III) : EDTA (1:1) Lu(III) : EDTA (1:1) Ternary Complexes La(III) : EDTA : Carbonate (1:1:1) Eu(III) : EDTA : Carbonate (1:1:1) Lu(III) : EDTA : Carbonate (1:1:1) UV-Vis Spectroscopy Titrations for Eu:EDTA:Carbonate (1:1:1) Potentiometric Results Summary 44 Chapter 4: Conclusion 45 Chapter 5: Future work 46 Chapter 6: References 47 Word Count: 8,450 3

4 List of Figures Figure 1: Representation of the nuclear fuel cycle. 1 Figure 2: Diagrammatic representation of the dry method of conversion for yellowcake to uranium hexafluoride. Figure 3: Diethylenetriaminepentaacetate (DTPA 5- ). Figure 4: A basic diagrammatical representation of the TALSPEAK process. Figure 5: A basic diagrammatical representation of the reverse TALSPEAK process. Figure 6: Diagrammatic representations of f orbitals. 16 Figure 7: Ethylenediaminetetraacetate (EDTA 4- ) Figure 8: Metrohm751 GPD Titrino apparatus. 28 Figure 9: Example of a speciation model from Hyperquad. Figure 10: To show transitions between two energy levels. Figure 11: Speciation diagram of the deprotonation of EDTA complexes. Figure 12: Potentiometric titration of La 3+ /EDTA complexation at 25 o C. I = 0.1 M NaNO 3. V 0 = 10 ml. C La = C EDTA = 10 mm. Titrant = 0.1 M NaOH. Figure 13: Potentiometric titration of Eu 3+ / EDTA complexation at 25 o C. I = 0.1 M NaNO 3. V 0 = 10 ml. C La = C EDTA = 10 mm. Titrant = 0.1 M NaOH. Figure 14: Potentiometric titration of Lu 3+ /EDTA complexation at 25 o C. I = 0.1 M NaNO 3. V 0 = 10 ml. C La = C EDTA = 10 mm. Titrant = 0.1 M NaOH. Figure 15: Speciation diagram for carbonate. Figure 16: Potentiometric titration of La 3+ /EDTA/Carbonate complexation at 25 o C. I = 0.1 M NaNO 3. V 0 = 10 ml. C La = C EDTA = C carbonate = 10 mm. Titrant = 0.1 M NaOH. Figure 17: Potentiometric titration of Eu 3+ /EDTA/Carbonate complexation at 25 o C. I = 0.1 M NaNO 3. V 0 = 10 ml. C La = C EDTA = C carbonate = 10 mm. Titrant = 0.1 M NaOH. Figure 18: Potentiometric titration of Lu 3+ /EDTA/Carbonate complexation at 25 o C. I = 0.1 M NaNO 3. V 0 = 10 ml. C La = C EDTA = C carbonate = 10 mm. Titrant = 0.1 M NaOH. 4

5 List of Tables Table 1: Ionic radii for Ln 3+ ions. 4 Table 2: Hydrolysis constants for selected lanthanides. 18 Table 3: Stability constants of Ln 3+ complexation with carbonate; I = 0.7 M NaClO Table 4: Stability constants for the deprotonation of Na 2 EDTAH 2 in solution. 18 Table 5: Stability constants for lanthanides La 3+, Eu 3+ and Lu 3+ ions with EDTA 4- ligand. 18 Table 6: Table to show the concentrations and volumes used for each titration. Table 7: Stability constants of La(III) complexation with EDTA and hydroxide; I = 0.1 M NaNO 3. Table 8: Stability constants of Eu(III) complexation with EDTA and hydroxide; I = 0.1 M NaNO 3. Table 9: Stability constants of Lu(III) complexation with EDTA and hydroxide; I = 0.1 M NaNO 3. Table 10: Stability constants of La(III) complexation with EDTA and carbonate; I = 0.1 M NaNO 3. Table 11: Stability constants of Eu(III) complexation with EDTA and carbonate; I = 0.1 M NaNO 3. Table 12: Stability constants of Lu(III) complexation with EDTA and carbonate; I = 0.1 M NaNO 3. Table 13: Summary of Log K values found when conducting potentiometry experiments 5

6 List of Equations Equation 1: Equation to calculate the distribution factors of solutes Equation 2: Equation to calculate separation factors Equation 3: The parameter ph is related the concentration of protons Equation 4: The Beer- Lambert law. 32 6

7 Abstract This project has been focused on modelling actinide behaviour using lanthanide analogues. Ternary complexes of Lanthanides-EDTA-Carbonate have been studied with techniques such as potentiometry and ultraviolet-visible absorption spectroscopy. It was found that ternary complexes of Ln-EDTA-Carbonate and Ln-EDTA-OH do exist for La, Eu and Lu. Log K values were calculated from potentiometry titrations which involve the formation of these ternary species. For lanthanum, the log K values for the complexes [La(EDTA)] -, [La(EDTA)(OH)] 2- and [La(EDTA)(CO 3 )] 3- were 14.20, and 4.70, respectively. For europium, the log K for the complexes [Eu(EDTA)] -, [Eu(EDTA)(OH)] 2- and [Eu(EDTA)(CO 3 )] 3- were 13.77, and 4.55, respectively. For lutetium, the results for the complexes [Lu(EDTA)] -, [Lu(EDTA)(OH)] 2-, [Lu(EDTA)(OH) 2 ] 2- and [Lu(EDTA)(CO 3 )] 3- were 17.38, , and 3.30, respectively. The results are useful for helping to understanding solution chemistry of lanthanides with EDTA in aqueous environments, which is applicable to actinide-lanthanide separations and actinide/lanthanide behaviour in storage pond conditions within the area of nuclear waste management. 7

8 Declaration No portion of this work has been submitted in support of an application for another degree or qualification of this or any other university or other institute of learning. 8

9 Copyright Statement The author of this dissertation (including any appendices and/or schedules to this dissertation) owns any copyright in it (the Copyright ) and she has given The University of Manchester the right to use such copyright for any administrative, promotional, educational and/or teaching purposes. Copies of this dissertation, either in full or in extracts, may be made only in accordance with the regulations of the John Rylands University Library of Manchester. Details of these regulations may be obtained from the librarian. This page must form part of any such copies made. The ownership of any patents, designs, trade marks and any and all other intellectual property rights except from the Copyright (the Intellectual Property Rights ) and any reproductions of copyright works, for example graphs and tables ( Reproductions ), which may be described in this dissertation, may not be owned by the author and may be owned by third parties. Such Intellectual Property Rights and Reproductions cannot and must not be made available for use without the prior written permission of the owner of the relevant Intellectual Property Rights and/or Reproductions. Further information on the conditions under which disclosure, publication and exploitation of this dissertation, the Copyright and any Intellectual Property Rights and/or Reproductions described in it may take place is available from the Head of School of Chemistry. 9

10 Acknowledgements I would like to thank my supervisor Dr. Clint Sharrad for all his brilliant guidance and help! Lucy Jones, Dan Whittaker, Nick Bryan and the rest of The University of Manchester Radiochemistry Department for all their help and support throughout my masters. And a massive, huge, gigantic thank you to Tamara Griffiths for being amazingly patient, supporting and for just generally everything she has done over this past year. 10

11 List of Abbreviations DTPA EDTA g HDEHP Ln An mm M ml mg nm NMR TALSPEAK Diethylene Triamine Pentaacetic Acid Ethylenediaminetraacetic Acid Gram Di(2-ethylhexyl) phosphoric acid Lanthanide Actinide Millimolar Molarity Millilitres Milligrams Nanometres Nuclear Magnetic Resonance Trivalent Actinide-Lanthanide Separation by Phosphorus Reagent Extraction from Aqueous Complexes UV-Vis Ultraviolet- Visible 11

12 Chapter 1: Introduction 1.1 Nuclear Fuel Cycle The nuclear fuel cycle is a series of processes that encompasses all aspects of nuclear energy generation, from mining uranium ore to nuclear reactors and the management of spent nuclear fuel. 1 There are various steps that make up the nuclear fuel cycle, these steps can be split up into 2 categories, the front end and the back end. The front end of the fuel cycle is the preparation stage, made up of mining and milling, conversion, enrichment, fuel fabrication and energy generation. The back end of the fuel cycle is the management of fuel after energy production. The steps involved could include temporary storage, reprocessing and recycling, and eventually disposal as waste. 1 Figure 1: Representation of the nuclear fuel cycle. 2 The nuclear fuel cycle begins with the mining of uranium ore. The uranium is found and removed by either underground or open pit methods. Mining which occurs in Canada, Australia and Kazakhstan contributes towards 62% of the world s production of uranium. 3 Next, the mined uranium is converted to uranium oxide U 3 O 8, which is also known as yellowcake, by crushing, separation and processing. The yellowcake is then converted by two-stage process into uranium hexafluoride gas. 1 The first stage of the process involves dissolving Yellowcake in nitric acid to produce UO 2 (NO 3 ) 2, which is then purified and denitrated to produce UO 3 powder. This is then reduced to UO 2 by hydrogenation and hydrofluorinated with HF, which results in UF 4. The next stage 12

13 involves the conversion of UF 4 to UF 6 by fluorination (F 2 ). This is the wet conversion method. There is another method of conversion that involves dry techniques. 1 The dry method is essentially the same, but unlike the wet method the first step is hydrogenation of yellowcake. The purification stage in the dry method is in the form of a fractional distillation of UF 6 at the end of the process. Figure 2 below shows step by step the dry process. Yellowcake H 2 UO 2 HF UF 4 F 2 UF 6 Fractional distillation Pure UF 6 Figure 2: Diagrammatic representation of the dry method of conversion for yellowcake to uranium hexafluoride. From here, the uranium needs to be enriched. Enrichment is a process to increase the concentration of 235 U atoms within the uranium material. It is an essential step, as for uranium to be useful in most nuclear reactors it must contain 2-3% of 235 U. 4 There are two common enrichment methods; gas diffusion and a gas centrifuge. 4 Both methods account for the isotopes of uranium in uranium hexafluoride having different mass presenting the basis by which isotopic separation can occur. After enrichment, the UF 6 is converted back to UO 2, which is then fabricated into pellet form. These fuel pellets are placed into tubes referred to as fuel rods. These fuel rods are then used to make up the nuclear fuel core of a power reactor. The nuclear fuel cycle can either be described as an open or a closed cycle. The closed cycle involves the reprocessing of spent nuclear fuel. Spent fuel is the product 13

14 after the irradiation of the fuel rods. Spent nuclear fuel can either be reprocessed (closed cycle) or disposed in a geological repository (open cycle). In many countries, especially the USA, reprocessing spent nuclear fuel is not considered to be ecologically viable, and so spent fuel rods are temporarily stored at the site until a long-term location for the waste is found. 5 Numerous chemical techniques have been developed for the reprocessing and treatment of spent fuel where particular elements are separated from the remaining fuel mixture. The PUREX process is the prime reprocessing method for spent nuclear fuels. It is a process whereby uranium and plutonium are separated from spent nuclear fuel so they can be used again within the nuclear fuel cycle (Plutonium-URanium EXtraction). 6 After the PUREX process has been used, the remaining waste is called PUREX raffinate. This waste contains minor actinides, which causes major problems in terms of safe management and storage, since they have long-term radioactivity. A possible way of decreasing the half-life of the actinides within the waste would be to employ neutron bombardment on the minor actinides so they undergo nuclear transmutation. 7 The problem with this technique is that lanthanides also exist within the PUREX raffinate. The separation of the trivalent actinides from lanthanides is necessary for the process of nuclear transmutation to be efficient. Nuclear transmutation is a method whereby long-lived isotopes are irradiated with fast neutrons in a nuclear reactor. This converts the long-lived isotopes to fission products with a significantly shorter halflife, thereby reducing the lifetime required to contain this material within a disposal facility. Without the separation process, the lanthanides would compete with the actinide ions for neutrons, used in the transmutation process, as the lanthanides have very large neutron capture cross sections. For this reason, it is important that actinides and lanthanides are separated. This process is called partitioning and transmutation. 7 The chemical behaviour of lanthanides and actinides are very similar, so separation of these species is difficult. There are a number of different processes for the partitioning of actinides from lanthanides; DIAMEX (DIAMide EXtraction), SANEX (Selective ActiNide EXtraction), TRUEX (TRansUranium EXtraction) and TALSPEAK (Trivalent Actinide-Lanthanide Separation by Phosphorus Reagent Extraction from Aqueous Complexes). 1,8,9,10 14

15 1.2 TALSPEAK The TALSPEAK process (Trivalent Actinide-Lanthanide Separation by Phosphorus Reagent Extraction from Aqueous Complexes) is a separating method for removing short-lived lanthanides from long-lived actinides in spent nuclear fuel. 10 This separation means that minor actinides with a long half life can be disposed of efficiently, and also with the possibility of allowing transmutation. It also means that it will permit a greater storage capacity for fuel containing short-lived isotopes. 11 The TALSPEAK process is a solvent extraction technique that uses the ligand diethylenetriamine pentaacetate (DTPA, Figure 3) in an aqueous phase, buffered with lactic acid, and di(2-ethylhexyl) phosphoric acid (HDEHP) in the organic phase. 12 It is believed separation of the lanthanides and actinides by this method is achieved on the basis of the metal ion hardness. The actinides are retained in the DTPA containing aqueous phase, while the lanthanides are extracted into the HDEHP containing organic phase. This is shown visually in Figure 4. The actinides are considered softer metals than lanthanides, so the actinides prefer binding to DTPA as it contains soft nitrogen donor atoms over HDEHP. The lanthanides prefer to coordinate to HDEHP, which only contains hard oxygen donor atoms. 13 Figure 3: Diethylenetriaminepentaacetate (DTPA 5- ) 15

16 Organic Phase HDEHP Ln 3+ DTPA/ Lactic acid An 3+ Aqueous Phase Figure 4: A basic diagrammatical representation of the TALSPEAK process, upon complete separation. Reverse TALSPEAK is a modified version of the original TALSPEAK process. In this method of separation the actinides and lanthanides present are extracted together in an organic phase containing HDEHP and TBP. TBP is used as a phase modifier, as it reduces the viscosity of the organic phase. 14 Then, the actinides are selectively back extracted to an aqueous phase, leaving the lanthanides in the organic phase (see figure 5). The aqueous phase used in the back extraction contains DTPA and an organic acid, such as lactic or citric acid, which gives it an acidity of around ph Organic Phase Organic Phase HDEHP / TBP An 3+ Ln 3+ HDEHP / TBP Ln 3+ HNO 3 DTPA/ Citric Acid An 3+ Aqueous Phase Aqueous Phase Step 1 Step 2 Figure 5: A basic diagrammatical representation of the reverse TALSPEAK process. The chemistry of the TALSPEAK separation process is still not fully understood and is still a work in progress. 15,16,17 16

17 Separations factors are used as a way of quantifying the performance of the TALSPEAK process. They are the measurement of how efficiently two solutes have been separated between the two liquid phases. Each solute will have a distribution ratio (Equation 1). This distribution ratio is then used to calculate the separation factor (Equation 2). D Solute 1 = [Solute in organic phase]/ [Solute in aqueous phase] Equation 1: Equation to calculate the distribution factors of solutes SF x/y = x/y where: D Solute 1 = x D Solute 2 = y Equation 2: Equation to calculate separation factors In 1964, separation factors in relation to the TALSPEAK process were determined by Boyd Weaver et al. 18 Although this review was extremely beneficial it is very dated, and so recent investigations into determining these separation factors have been undertaken

18 1.3 The f- elements The f elements are made up of the lanthanide and actinide series. The shorthand used for each series is Ln and Ac, for the lanthanides and actinides, respectively. Lanthanides and actinides are defined as the f elements as it is the 4f and 5f orbitals that progressively fill moving from left to right across the period. The lanthanides consist of the elements lanthanum (La, Atomic number=57) to lutetium (Lu, Atomic Number =71). The actinides consist of the elements from actinium (Ac, atomic number = 89) to lawrencium (Lr, atomic number = 103). Figure 6 below shows representations of the f orbitals. 5 Figure 6: Diagrammatic representations of f orbitals

19 1.4 The Lanthanides The most common oxidation state of the lanthanides is +3. This is due the 4 th ionisation energy being greater than the sum of the first three ionisation energies, for most of the lanthanides. Traversing across the lanthanide series, there is a decrease in ionic radii (Table 1). There is therefore an increase in charge density across the 4f series. The contracted nature of the 4f orbitals causes the lanthanides to participate in ionic bonding. They are hard Lewis acids and so prefer binding with hard Lewis bases such as hydroxide, halogens and carbonate. The greater the charge density of the ions, the stronger the ionic bonds. This explanation for the decrease in ionic radii across the series is referred to as the lanthanide contraction. 21 Element Ln 4+ Electronic Ln 4+ Ionic Radius / Configuration pm La [Xe] Ce [Xe]4f Pr [Xe]4f Nd [Xe]4f Pm [Xe]4f Sm [Xe]4f Eu [Xe]4f Gd [Xe]4f Tb [Xe]4f Dy [Xe]4f Ho [Xe]4f Er [Xe]4f Tm [Xe]4f Yb [Xe]4f Lu [Xe]4f Table 1: Ionic radii for Ln 3+ ions. 5 The lanthanides can coordinate to up to twelve ligands as they are particularly large ions. Although, coordination number does decrease with decreasing ionic radius. 19

20 1.4.1 Lanthanide Hydrolysis Hydrolysis of metal ions dominates in aqueous solutions, with the absence of organic ligands and with low carbonate concentrations. 22 Knowing the hydrolysis constants of lanthanide ions is important for understanding their chemical behaviour in natural environments. The susceptibility of hydrolysis of the lanthanides increases across the series and is shown in table 2 below. La 3+ + OH - La(OH) 2+ log β La OH - + La(OH) 2 log β La(OH) 3(s) La OH - log K sp La OH - 3+ La 2 (OH) 3 log β La OH - 6+ La 5 (OH) 9 log β La OH - 8+ La 6 (OH) 10 log β Eu 3+ + OH - Eu(OH) 2+ log β Eu(OH) 3(s) Eu OH - log K sp Eu OH - - Eu(OH) 4 log β Eu OH - 3+ Eu 2 (OH) 3 log β Lu 3+ + OH - Lu(OH) 2+ log β Lu(OH) 3(s) Lu OH - log K sp Table 2: Hydrolysis constants for selected lanthanides Lanthanide Carbonate Complexes Lanthanides form soluble, highly charged anionic species in concentrated solutions of alkali metal carbonates. 24 The high basicity of the CO 2-3 anions and the large effective charge of the Ln 3+ cations give these complexes their substantial stability. Table 3 below shows the stability constants for lanthanide carbonate complexes. Reaction log K (literature values) 2 La CO 3 [La(CO 3 )] Eu CO 3 [Eu(CO 3 )] Lu CO 3 [Lu(CO 3 )] Table 3: Stability constants of Ln 3+ complexation with carbonate; I = 0.7 M NaClO

21 1.4.3 Lanthanum Lanthanum (Lu) has an atomic number of 57 and is a silvery, white, malleable metallic element, but is more commonly found in the oxidation state of +3. Carl Gustav Mosander discovered lanthanum in 1839 when it was discovered in a sample of cerium salt that he was working on at the time. The name lanthanum came from the Greek word lanthano which translates to to be hidden. It was isolated to pure form in General applications of lanthanum (III) include swimming pool additives, as lanthanum binds phosphates which otherwise would act as a fertilizer and encourage algae growth. Lanthanum binding to phosphates is also used for medicinal purposes. Phosphate in blood can cause bone problems and can be found in large quantities in patients with kidney disease. Lanthanum works by preventing absorption of phosphate from food when administered in the stomach. 26 Lanthanum (III) has an empty outer electronic shell, so 4f-4f transitions do not occur. For this reason lanthanum complexes generally do not exhibit any metal-based absorptions in the UV-Visible region of the spectrum Europium Europium (Eu) is a moderately hard, silvery metal with an atomic number of 63. In metallic form, it readily oxidises in water and air. Europium was discovered in 1892, but was not isolated into its pure form until 1901, by the French chemist Eugene- Anatole Demarcay. He named europium after Europe. 27 It has a common oxidation state of +3 but also can exist in the +2 oxidation state. Eu 3+ has an incomplete outer electron shell, so 4f-4f transitions are able to take place, giving Eu 3+ absorption properties in the visible region of the electromagnetic spectrum. The absorptions are weak as the 4f-4f transitions are Laporte forbidden. One application for europium compounds is that europium oxide, along with yttrium and gadolinium, is used as the red phosphor in colour television tubes

22 1.4.5 Lutetium Lutetium (Lu) is a silvery, white metal and has an atomic number of 71. It can be considered as the last element in the lanthanide series and so has the smallest ionic radius within the series due to lanthanide contraction. As with the other lanthanides, lutetium always assumes a +3 oxidation state. 21 Georges Urbain and Carl Auer Von Welsbach discovered lutetium in It was named after Lutetia, which is the Latin name for Paris. This stemmed from the fact that Georges Urbain was a French chemist. Lutetium is rare and therefore quite an expensive element, so its use commercially is somewhat limited. 21 Lutetium has a full outer 4f electronic shell, so 4f-4f transitions do not occur. For this reason lutetium complexes typically do not have any absorptions in the visible region of the spectra

23 1.5 EDTA Figure 7: Ethylenediaminetetraacetate (EDTA 4- ) Ethylenediaminetraacetate (EDTA 4-, Figure 7) is a polyaminocarboxylate. It mostly forms 1:1 complexes with metal ions and binds through the four carboxy O-atoms, and the two N-donors. EDTA typically binds to metal ions in a hexadentate manner but, depending on the size of the co-ordinated metal ion, it can bind pentadentate or tetradentate. EDTA was chosen as the multidentate, organic ligand to probe ternary complex formation with CO 2-3 and OH - due to the propensity for EDTA to form highly soluble metal complexes in water. The EDTA ligand can effectively cap the majority of the Ln coordination sites, leaving the remaining few sites free for coordination to other ligands that may be present. Na 2 EDTAH 2 forms acidic solutions due to the deprotonation of the acetate groups, stability constants for the deprotonations are found in table 4. Stability constant EDTAH 3- (aq) EDTA 4- (aq) + H + (aq) log β 11 = EDTAH 2 (aq) EDTAH 3- (aq) + H + (aq) log β 11 = EDTAH 3 (aq) EDTAH 2 (aq) + H + (aq) log β 11 = EDTAH 4(aq) EDTAH 3 (aq) + H + (aq) log β 11 = Table 4: Stability constants for the deprotonation of Na 2 EDTAH 2 in solution 21 Early studies of the stability of lanthanides and EDTA were undertaken by F.H. Spedding et al. In this experiment a potentiometric method was used whereby the complex would be formed or decomposed simultaneously with the production or consumption of an equivalent amount of hydrogen ions. The concentration of these ions gives information on the reactions progress and also allows calculation of the 23

24 stability constant of the reaction. The results of these experiments indicated an increase in stability constant of the Ln-EDTA complexes with increasing atomic number. 29 Spedding then developed this analogy and applied it to be able to use EDTA in the separation of lanthanide metals by ion exchange chromatography. 30 EDTA is of high interest in nuclear waste management operations 31 and has been reported to be used as a decontaminating agent. 32 Aminopolycarboxylic chelating agents such as EDTA and also DTPA are used to decontaminate reactor walls and fuel cladding that has become active due to neutrons and gamma rays being used during nuclear fuel generation. 33 The chelating ligands form strong complexes with the radionuclides and increases the solubility, making it easier to clean contaminated surfaces. This is a positive in terms of decontamination, but EDTA forming strong bonds with actinides also has its negatives. If EDTA is still present in nuclear waste disposal, then radionuclides will bind to it. Radionuclides would usually bind to humic acids and become immobilised in the soil. EDTA stops this from happening as it has a higher affinity for radionuclides than humic acid. 34 The stability of lanthanide- EDTA complexes is due to the large and positive change in entropy during complex formation. Table 5 shows log betas for the lanthanide series binding to the EDTA 4- ligand, proving the thermodynamic stabilities of these complexes are relatively high. These stability constants have been obtained from potentiometric titration experiments, which were reported by Martell et al. 23 Metal Ion Ligand Stability Constant La 3+ EDTA 4- [La(EDTA)] - log β Ce 3+ EDTA 4- [Ce(EDTA)] - log β Pr 3+ EDTA 4- [Pr(EDTA)] - log β Nd 3+ EDTA 4- [Nd(EDTA)] - log β Sm 3+ EDTA 4- [Sm(EDTA)] - log β Eu 3+ EDTA 4- [Eu(EDTA)] - log β Tb 3+ EDTA 4- [Tb(EDTA)] - log β Ho 3+ EDTA 4- [Ho(EDTA)] - log β Er 3+ EDTA 4- [Er(EDTA)] - log β Yb 3+ EDTA 4- [Yb(EDTA)] - log β Lu 3+ EDTA 4- [Lu(EDTA)] - log β Table 5: Stability constants for lanthanides ions with EDTA 4- ligand

25 1.6 Potentiometric titrations Figure 8: - Metrohm751 GPD Titrino apparatus 35 Potentiometry is an electrochemical analysis method, which measures the potential of an electrode system. Potentiometry is used as a way of finding the location of the equivalence point of a titration. The type of potentiometric titration used in this experiment is the standard addition method, whereby known measures of base are added to an acidic solution, and the change in ph is recorded 36. ph is related the concentration of protons, as shown in equation 3 below. ph= -log 10 [H + ]. Equation 3: The parameter ph is related the concentration of protons In a potentiometric cell containing, for example, lanthanide and EDTA ions in aqueous solution, there is competition of lanthanide and proton binding to EDTA. As the ph is increased, the concentration of protons decreases which favours the lanthanide binding to EDTA. At high ph, the concentration of hydroxide increases which will also compete for lanthanide binding. 25

26 1.7 Speciation Models Speciation models are used to indicate the proportion of species that are present in a system at a specific ph. Changes in concentration can be modelled to fit the equilibria that are present in the solution. Figure 9 below shows an example of a speciation diagram that has been created by using the computer software Hyperquad. The titration data used is that from potentiometric experiments, which gives values for ph vs. amount of acid/base added. The data was modelled to a series of equilibria that could be present within the solution. The stability constants for the equilibria that have been previously established in the literature are fixed in the model, while the remaining unknown stability constants are refined in an iterative process. Equilibria may be removed or added and the refinement process continued until appropriate fit of the modelled titration curve to the experimental data is obtained. [La(EDTA)] - (aq) ph(exp.) --- ph(cal.) Free La 3+ [La(EDTA)(OH)] 2- (aq) mols of added Titrant Figure 9: Example of a titration curve fitting using Hyperquad 26

27 1.8 Ultraviolet visible spectroscopy Ultraviolet visible spectroscopy (UV-Vis) is a form of electronic spectroscopy, which deals with the transitions of electrons between molecular orbitals. 37 The orbitals that the electron is being excited from and to each have a specific energy. If an appropriate amount of energy is absorbed by a molecule, then an electron may be excited from an occupied orbital to a partially or unoccupied orbital. These transitions occur when light of energy ΔE is absorbed by a molecule, as shown in Figure 10. Energy E 2 ΔE E 1 Figure 10: Diagrammatical representation to show the transition between two energy levels Absorption of most metal complexes typically occurs in the ultraviolet and visible regions of the electromagnetic spectrum between 200 and 800 nm. An absorption spectrum is used to identify the molecule in question, which shows a series of absorption bands. Different types of molecules will absorb the UV-Vis light at different wavelengths, which give rise to different absorption bands. The concentration or absorbance of a compound can be determined by UV-Vis spectroscopy. The Beer-Lambert law relates the absorbance to the concentration, cell path length and molar extinction coefficient of a solution sample. 38 The Beer-Lambert law is stated in equation 4 below. A = ε x c x l A= absorbance, ε= extinction coefficient, l = path length, c = concentration Equation 4: The Beer- Lambert law. 38 In Ln 3+, the 4f electrons are heavily shielded from ligands by the 5s and 5p orbitals. This is due to the large radial contraction of the 4f orbitals, which means that f-f 27

28 transitions are unlikely. The f-f transitions that are observed do not shift substantially when the coordination environment changes, compared to those for d transition metals. UV-Vis analysis of Ln 3+ species typically gives sharp peaks that are due to the 4f-4f transitions. These transitions are forbidden by selection rules, and so have low extinction coefficients, relative to the analogues d-d transitions. 38 Within this study UV-Vis is used as an extra method of analysis for ternary lanthanide complexes, by observing the effect of ph on the electronic absorption spectrum of [Eu(EDTA)] - (aq) in the presence and absence of carbonate. Nicole Graeppi et al completed a similar study, with the difference in their experiment being that they analysed the effect of temperature and pressure on their europium complexes

29 1.9 Aims The aims of this project are to investigate ternary complexes of lanthanides/edta/carbonate with the techniques potentiometry and ultravioletvisible spectroscopy. The reason for this is to gain more information of the behaviour of lanthanides, EDTA and carbonate so it can be applied to reprocessing, decontamination or decommissioning techniques relevant to the nuclear fuel cycle (see section 1.1). Potentiometry is a method where the competition process between lanthanide metal ions, ligands (e.g. EDTA, OH -, CO 2-3 ) and protons can be analysed. Changing the ph of an aqueous solution containing these components can be used to determine the thermodynamic parameters of the equilibria that occur in a given system. The potentiometric titration results will be analysed and modelled using the computer program Hyperquad. Stability constants for Ln-OH and Ln-EDTA species (Ln= La,Eu,Lu) will first be modelled and compared to literature values. The results from this will then be used in modelling for Ln-carbonate species. La, Eu and Lu were selected as a broad cover of the lanthanide series. One from the beginning, one from the middle and one from the end of the series. This will allow the possibility to understand how ternary complex formation is influenced by the lanthanide contraction. Unlike La 3+ and Lu 3+, Eu 3+ has absorption properties and so UV-Vis spectroscopy will be used as an extra method of analysis to see if ph affects the spectra of Eu- EDTA-CO 3 2- system. This project will be applied to the nuclear fuel cycle in the sense that it will provide information on some species that exist within nuclear reprocessing. The findings can be applied to TALSPEAK separation in that greater understanding of the lanthanides and their complexes within solution will be investigated. EDTA is used as a ligand as theoretically it is likely that ternary complexes will be formed. EDTA is generally very soluble and easy to manage, compared to DTPA. 29

30 Chapter 2: Experimental This section documents the chemicals, preparation and instrumentation used in this research. 0.1 M NaNO 3 / 6 mm nitric acid solution was used as a background electrolyte for the solutions. 2.1 Potentiometric Titrations Solution Preparation [Ln 3+ ] / [EDTA]/ [(CO 3 ) 2- ]/ Volume of each mm mm mm used/ ml Ln(III) (aq) Ln (III):EDTA, 1: Ln(III):EDTA:Carbonate, 1:1: Table 6: Table to show the concentrations and volumes used for each titration Lanthanide (III) Nitrate Solutions Solutions were prepared by dissolving a known mass of La(NO 3 ) 3, Eu(NO 3 ) 3 and Lu(NO 3 ) 3 (all %, Sigma Aldrich) in 0.1 M NaNO 3 / 6 mm nitric acid solution. A typical experimental solution was 3.33 ml to 10.0 ml, with a lanthanide ion concentration ranging from 10 mm to 30 mm Ligand Solutions Na 2 EDTA (BDH Chemical), Na 2 CO 3 (Fisher Chemical) were used without further purification. Solutions of Na 2 EDTA were prepared by dissolving of a known mass of ligand in 0.1 M NaNO 3 / 6 mm nitric acid solution. A typical experiment solution was 3.33 ml to 5 ml with a known concentration ranging from 20.0 mm to 30 mm. Solutions of Na 2 CO 3 were prepared by dissolving of a known mass of ligand in 0.1 M NaNO 3 solution. A typical experiment solution was 3.33 ml with a known concentration of 30 mm. 30

31 2.1.3 Lanthanide(III)-EDTA Binary Complex Solutions Typical experimental solutions of both the lanthanide and EDTA were 5 ml of each, with known lanthanide and ligand concentrations of 20 mm Lanthanide(III)-EDTA /Na 2 CO 3 Ternary Complex Solutions Within potentiometric titrations, where Na 2 CO 3 was used as the secondary ligand, the ph of the solution is increased to ph 7 using NaOH solution before the addition of carbonate anions. Typical experimental solutions of the lanthanide, EDTA and the carbonate were 3.33 ml of each, with a known lanthanide, EDTA and the carbonate concentration of 30 mm Carbonate free NaOH A 25 M NaOH solution was prepared under nitrogen and left to stand for 24 hours. The solution was filtered to remove any carbonate precipitate. From this, a 0.1 M solution was made by dilution of the 25 M NaOH solution with boiled de-ionised water (after cooling). This was then standardised against 0.1 M HCl (Fisher Scientific) by titration. 31

32 2.2 Instrumentation Potentiometry All titrations were carried out on a 751 GPD Titrino apparatus using a Metrohm aquatrode electrode filled with KCl electrolyte, which is within a glass cell. The cell temperature was controlled by a pumping water bath and maintained at 298 K. The apparatus was controlled remotely using a Dell Inspirion 5100 computer loaded with Tiamo 1.2 software. The data was analysed using Hyperquad 2006 software. All titrations were prepared and run within a nitrogen environment, which was used to eliminate CO 2, which could interfere with the titrations. This consisted of an airtight Perspex glove box with a nitrogen feed. The nitrogen feed was treated with concentrated NaOH to remove any inherent carbonate from the system. The solutions for the calibration were ph 4.00 (±0.01), 7.00 (± 0.01) and (± 0.01) standard buffers (Fisher Scientific). An aqueous alkali solution of Carbonate free NaOH was always used as the titrant. The alkali solution used was standardised using analytical grade aqueous HCl. The maximum volume and molarity of the alkali solution was 1000 ml and 0.1 M, respectively. Ionic strength was kept constant throughout the experimentation using an aqueous electrolyte solution of 0.1 M NaNO 3 / 6 mm nitric acid. The maximum volume and molarity of the electrolyte solution was 1000 ml and 0.1 M, respectively UV-Vis Spectroscopy The UV-Vis spectra were recorded on a Cary Varian 500 scan UV-vis-NIR spectrophotometer, with a typical scan rate of 600 nm min -1. The background spectra of water were obtained and subtracted for each of the experimental spectra required. UV-Vis spectroscopy was used to observe the effect of ph on the absorption profile of [Eu(EDTA)] - (aq) in the presence and absence of carbonate. 32

33 Chapter 3: Results and Discussion Previous work has shown that the species [Ln(EDTA)(CO 3 )(H 2 O)] 3- and [Ln(EDTA)(OH)(H 2 O)] 2- are likely to form for all the Ln series. 6 Therefore, these two species have been included within the speciation model. The second hydrolysis of [Ln(EDTA)] - was also considered within the speciation model to assess if inclusion of this species improved the fits for the potentiometric data. The [Ln(EDTA)(CO 3 ) 2 ] 3- complex was not used for the model as its unlikely to form due to steric hindrance around the Ln 3+ ion. There is also increased electronic repulsion between the [Ln(EDTA)(CO 3 )] 3- complex and another carbonate anion. All experimental speciation diagrams (Figures 11-17) show discrepancy between the experimental results and calculated model. This suggests that all the possible species may not have been correctly identified. The models presented are those that gave the best fit for species of [Ln(EDTA)] -, [La(EDTA)(OH)] 2- and [Ln(EDTA)(CO 3 )] 3--. All potentiometry experiments were run under nitrogen and titrated with carbonate free NaOH. A number of experiments were first performed in different conditions, such as without the nitrogen environment and carbonate free NaOH. Not much difference was seen in the results but the nitrogen environment and the carbonate free NaOH were kept for consistency. All presented potentiometric titrations were completed in triplicate and the averages taken and used for the titration values. 33

34 3.1 Binary Complexes H 2EDTA 2 - HEDTA 3- EDTA 4- % Composition H 3 EDTA - H 5 EDTA + ph Figure 11: Speciation diagram of the deprotonation of EDTA complexes. 40 Figure 11 above shows a speciation diagram of the different EDTA species that can occur at a certain ph. As ph is increased, so is deprotonation of the EDTA ligand. 34

35 3.1.1 La:EDTA (1:1) [La(EDTA)] - (aq) ph(exp.) --- ph(cal.) Free La 3+ [La(EDTA)(OH)] 2- (aq) mols of added NaOH Figure 12: Potentiometric titration of La 3+ /EDTA complexation at 25 o C. I = 0.1 M NaNO 3. V 0 = 10 ml. C La = C EDTA = 10 mm. Titrant = 0.1 M NaOH Potentiometric titrations of 1:1 acidic mixtures of La(NO 3 ) 3 and Na 2 EDTA in 0.1 M NaNO 3 against NaOH were performed (Figure 12). The stability constant of the [La(EDTA)] - (aq) complex was determined by potentiometry (Table 7). The speciation diagram in Figure 11 shows that the dominant species in solution up to around ph 11.5 is the [La(EDTA)] - (aq) complex. After this the [La(EDTA)(OH)] 2- (aq) complex then becomes dominant. The speciation diagram also indicates that around ph 10, hydrolysis of the [La(EDTA)] - (aq) complex begins. Reaction log K La 3+ + EDTA 4- [La(EDTA)] ± 0.01 [La 3+ ( EDTA) 4- ] - + OH - [La(EDTA)(OH)] ± 0.01 Table 7: Stability constants of La 3+ complexation with EDTA and hydroxide; I = 0.1 M NaNO 3 35

36 3.1.2 Eu: EDTA (1:1) [Eu(EDTA)(OH)] 2- (aq) [Eu(EDTA)] - (aq) ph(exp.) --- ph(cal.) Free Eu 3+ mols of added NaOH Figure 13: Potentiometric titration of Eu 3+ /EDTA complexation at 25 o C. I = 0.1 M NaNO 3. V 0 = 10 ml. C Eu = C EDTA = 10 mm. Titrant = 0.1 M NaOH Potentiometric titrations of 1:1 acidic mixtures of Eu(NO 3 ) 3 and Na 2 EDTA in 0.1 M NaNO 3 against NaOH were performed (Figure 13). The stability constants of the [Eu(EDTA)] - (aq) and [Eu(EDTA)(OH)] 2- (aq) complexes have been determined by potentiometry (Table 8). The speciation diagram in Figure 12 shows that up to ph 11 the [Eu(EDTA)] - (aq) complex is the dominant species in solution. After this, the [Eu(EDTA)(OH) 2- ] (aq) dominates. Stability constant of the [Eu(EDTA) - ] complex has increased compared to that of [La(EDTA) - ]. Reaction log K Eu 3+ + EDTA 4- [Eu(EDTA)] ± 0.17 [Eu 3+ ( EDTA 4- )] - + OH ± 0.22 [Eu(EDTA)(OH)] 2- Table 8: Stability constants of Eu 3+ complexation with EDTA and hydroxide; I = 0.1 M NaNO 3 36

37 3.1.3 Lu:EDTA (1:1) [Lu(EDTA)] - (aq) ph(exp.) --- ph(cal.) [Lu(EDTA)(OH)] 2- (aq) Free Lu 3+ mols of added NaOH Figure 14: Potentiometric titration of Lu 3+ /EDTA complexation at 25 o C. I = 0.1 M NaNO 3. V 0 = 10 ml. C Lu = C EDTA = 10 mm. Titrant = 0.1 M NaOH Potentiometric titrations of 1:1 acidic mixtures of Lu(NO 3 ) 3 and Na 2 EDTA in 0.1 M NaNO 3 against NaOH were performed (Figure 14). The stability constants of the [Lu(EDTA)] - (aq) and [Lu(EDTA)(OH)] 2- (aq) complexes have been determined by potentiometry (Table 9). The speciation diagram in Figure 13 shows that up to around ph 11 the [Lu(EDTA)] - (aq) complex is the dominant species in solution. After this, the [Lu(EDTA)(OH) 2- ] (aq) species dominates. Reaction log K Lu 3+ + EDTA 4- [Lu(EDTA)] ± 2.25 [Lu 3+ (EDTA 4- )] - + OH - [Lu(EDTA)(OH)] ± 2.25 Table 9: Stability constants of Lu 3+ complexation with EDTA and hydroxide; I = 0.1 M NaNO 3 37

38 3.2 Ternary complexes H 2 CO 3 HCO 3 - CO 3 2- % Composition ph Figure 15: Speciation Diagram for Carbonate. 40 Figure 15 shows a speciation diagram for carbonate equilibria in solution. There are three species and each are dependant on ph. 38

39 3.2.1 La: EDTA: Carbonate (1:1:1) ph(exp.) --- ph(cal.) [La(EDTA)(OH)] 2- (aq) [La(EDTA)(CO 3 )] 3- (aq) [La(EDTA)] - (aq) mols of added NaOH Figure 16: Potentiometric titration of La 3+ /EDTA/Carbonate complexation at 25 o C. I = 0.1 M NaNO 3. V 0 = 10 ml. C La = C EDTA = C carbonate = 10 mm. Titrant = 0.1 M NaOH Potentiometric titrations of 1:1:1 acidic mixtures of La(NO 3 ) 3, Na 2 EDTA and Na 2 CO 3 in 0.1 M NaNO 3 against NaOH were performed (Figure 16). The stability constant of the [La(EDTA)(CO 3 )] 3- (aq) complex was determined by potentiometry (Table 10). The speciation diagram in Figure 15 shows that the dominant species in solution up to around ph 10.5 is the [La(EDTA)(CO 3 )] 3- (aq) complex. After this the [La(EDTA)(OH)] 2- (aq) complex then becomes dominant. Reaction log K [La(EDTA)] CO 3 [La(EDTA)(CO 3 )] ± 0.32 Table 10: Stability constant for La 3+ complexation with EDTA and carbonate; I = 0.1 M NaNO 3 39

40 3.2.2 Eu:EDTA:Carbonate (1:1:1) ph(exp.) --- ph(cal.) [Eu(EDTA)(CO 3 )] 3- (aq) [Eu(EDTA)(OH)] 2- (aq) [Eu(EDTA)] - (aq) [Eu(OH) 3 ] 2- mols of added NaOH Figure 17: Potentiometric titration of Eu 3+ /EDTA/Carbonate complexation at 25 o C. I = 0.1 M NaNO 3. V 0 = 10 ml. C Eu = C EDTA = C carbonate = 10 mm. Titrant = 0.1 M NaOH Potentiometric titrations of 1:1:1 acidic mixtures of Eu(NO 3 ) 3, Na 2 EDTA and Na 2 CO 3 in 0.1 M NaNO 3 against NaOH were performed (Figure 17). The stability constant of the [Eu(EDTA)(CO 3 )] 3- (aq) complex was determined by potentiometry (Table 11). The speciation diagram in Figure 16 shows that the dominant species in solution at ph 11.2 is the ternary [Eu(EDTA)(CO 3 )] 3- (aq) complex. Reaction log k [Eu(EDTA)] CO 3 [Eu(EDTA)(CO 3 )] ± 0.50 Table 11: Stability constant for Eu 3+ complexation with EDTA and carbonate; I = 0.1 M NaNO 3 40

41 3.2.3 Lu: EDTA: Carbonate (1:1:1) ph(exp.) --- ph(cal.) [Lu(EDTA)(CO 3 )] 3- (aq) [Lu(EDTA)] - (aq) [Lu(EDTA)(OH) 2 ] 3- (aq) [Lu(EDTA)(OH)] 2- (aq) mols of added NaOH Figure 18: Potentiometric titration of Lu 3+ /EDTA/Carbonate complexation at 25 o C. I = 0.1 M NaNO 3. V 0 = 10 ml. C Lu = C EDTA = C carbonate = 10 mm. Titrant = 0.1 M NaOH Potentiometric titrations of 1:1:1 acidic mixtures of Lu(NO 3 ) 3, Na 2 EDTA and Na 2 CO 3 in 0.1 M NaNO 3 against NaOH were performed (Figure 18). The stability constant of the [Lu(EDTA)(CO 3 )] 3- (aq) complex was determined by potentiometry (Table 12). The speciation diagram in Figure 16 shows that the dominant species in solution at ph 10.5 is the ternary [Lu(EDTA)(CO 3 )] 3- (aq) complex. Reaction log K [Lu(EDTA)] CO 3 [Lu(EDTA)(CO 3 )] ± 0.09 Table 12: Stability constant for Lu 3+ complexation with EDTA and carbonate; I = 0.1 M NaNO 3 41

42 3.3 Uv-Vis spectroscopy titrations for Eu:EDTA:Carbonate (1:1:1) The effect of ph on the electronic absorption spectrum of [Eu(EDTA)] - (aq) in the presence and absence of carbonate was studied (Figure 19 and Figure 20). In the electronic absorption spectra of [Eu(EDTA)] - (aq) (Figure 19), as ph is increased the extinction coefficient for each of the absorption maxima decreases slightly. This is probably due to hydrolysis of [Eu(EDTA)] - to form the [Eu(EDTA)(OH)] 2- species, which is included in the relevant potentiometric experiment (see section 3.1.2). The main transition at 397 nm corresponds to the 7 F 0 to 5 D 0 transition. For the electronic absorption spectra of [Eu(EDTA(CO 3 )] 3- (Figure 20), as ph is increased the extinction coefficient is increasing. This indicates the binding of carbonate to the [Eu(EDTA)] - complex. There is an increase in the absorbance of the bands at 362 nm, 375 nm and 382 nm. This is not observed in the [Eu(EDTA)] - (aq) binary system. As there is a difference between the spectra when carbonate is present and absent, it suggests that the ternary [Eu(EDTA)(CO 3 )] 3- (aq) complex is forming in solution. UV-Vis experiments confirmed the presence of the interaction between carbonate and EDTA, therefore the [Ln(EDTA)(CO 3 )] 3- species was included in the model for fitting the potentiometric data (see section 3.2) 42

43 1.4 EuEDTA ph 6 EuEDTA ph 9.65 EuEDTA ph Absorption [Eu]/ dm- 3 mol- 1 cm Wavelength/nm Figure 19: Electronic absorption spectra of 1:1 Eu:EDTA as ph is increased; [Eu] i =[EDTA] i = 30 mm 2.0 EuEDTACarbonate ph 8.7 EuEDTACarbonate ph EuEDTACarbonate ph Absorption [Eu]/ dm3 mol- 1 cm Wavenumber/nm Figure 20: Electronic absorption spectra of 1:1:1 Eu:EDTA: Carbonate as ph is increased [Eu] i =[EDTA] i =[CO 3 2- ] i = 30 mm 43

44 3.4 Potentiometric Results Summary [Ln(EDTA)] - [Ln(EDTA)(OH)] 2- [Ln(EDTA)(CO 3 )] 3- La ± ± ± 0.32 Eu ± ± ± 0.50 Lu ± ± ± 0.09 Table 13: Summary of Log K values found when conducting potentiometry experiments. 44

45 Chapter 4: Conclusions The results summarised in Table 13 show an increase in the stability constants for the [Ln(EDTA)] - complexes as the series is traversed. Lutetium has a higher charge density than europium, which has a higher charge density than lanthanum. The higher the charge density, the stronger ionic interaction with EDTA. Hence, the stability constants increase across the lanthanide series Table 13 also shows stability constants for the formation of [Ln(EDTA)(OH)] 2- for lanthanum, europium and lutetium. The trend shows the hydrolysis constants becoming more negative as the series is crossed. This suggests that the higher the charge density, the more susceptible the lanthanide ion is to hydrolysis, as would be expected The log K values for the complexation of carbonate to [Ln(EDTA)] - decrease across the lanthanide series. This contradicts with the theory that the harder the metal ion, the stronger the ionic interaction with ligands. It is unclear why this trend occurs, but further work will be performed to understand this trend (see chapter 5). Formation of the ternary [Ln(EDTA)(CO 3 )] 3- complex was found to occur at the ph region of around 9 to 11. At ph 11 hydrolysis competes to form the [Ln(EDTA)(OH)] 2- complex, with the carbonate ligand replaced by hydroxide. 2- Evidence has been found for ternary complexes with CO 3 and OH -. This has applications to reprocessing as part of the nuclear fuel cycle. Separation of actinides from lanthanides is important in reprocessing as it allows the actinides to be extracted from the waste and then potentially reused or stored correctly. Understanding the chemistry of the lanthanides and how it forms ternary complexes can be applied to these separation techniques, especially TALSPEAK. 45