Synthesis and Spectroscopic Characterisation of Lanthanide Polyoxometalate Complexes

Synthesis and spectroscopic characterisation of Lanthanide polyoxometalate complexes

1. Abstract

The syntheses and the crystal structures K₁₁[Ln³⁺(PMo₁₁O₃₉)₂]·16H₂O (Ln³⁺ = Pr and Gd) are reported in which an Ln³⁺ cation is sandwiched between two ‘lacunary’ [PMo₁₁O₃₉]⁷⁻ anions to give a complex with eight oxygen atoms coordinated to the lanthanide centre in a twisted square antiprismatic geometry. Phosphomolybdate solution speciation was controlled as a function of pH characterising the complexes synthesised. A pH of 4.3 was determined as the optimum pH for stabilisation of the defect lacunary phosphomolybdate anion, [Mo₁₁O₃₉]^7-. Both compounds were synthesised in high yield and characterised by spectroscopic techniques. UV/Vis/nIR spectroscopy gave a clear indication that the [Ln³⁺(PMo₁₁O₃₉)₂]¹¹⁻anion is also stable in solution. As Ln³⁺ cations contract across the lanthanide series the Ln–O bond distances decrease and the splitting of the ν_P–O vibrational mode within the [PMo₁₁O₃₉]⁷⁻ complex increases, following the lanthanide contraction phenomenon. The electronic structures of Pr and Gd complexes, with general formula [Ln(PMo₁₁O₃₉)₂]^11- where Ln is the corresponding lanthanide, were explored with a focus on the splitting of the f-orbitals, which provided information about the strengths of the metal–ligand interactions. Due to lanthanide contraction, it could be concluded that the smaller lacunary complex, Gd, displayed less splitting between the P-O stretch than the Pr complex. Characteristic absorption bands of lanathanides in the UV/Vis region corresponding to the 4f transitions were used for the quantitive determination of the elements. Upon analysis of the IR spectra it can be deemed that complexation of the Ln(III) cation to the vacancy in the lacunary anion increases the splitting between the P-O stretching modes.

 

2. Introduction

The lanthanide series of elements are made up of 15 metallic elements ranging from atomic number 57 to 71. The 15 lanthanide elements, from lanthanum to lutetium, including the chemically similar scandium and yttrium elements are typically deemed the rare earth elements. In general, the term ‘Ln’ is used to represent any of the lanthanide chemicals ^[1]. The lanthanides are f-block elements due to the series filling the 4f electron shell. The effect of the 4f orbitals on the chemistry of the lanthanides is the fundamental difference between them and the transition metals. The 7 4f orbitals are: f_z³, f_xz², f_yz², f_xyz, f_z(x²_−y²₎, f_x(x²_−3y²₎ and f_y(3x²_−y²₎  ^[2]. The 4f orbitals of lanthanide elements are “isolated within the [Xe] core” and hence do not participate in bonding; because of this crystal field effects observed are small and also why lanthanide elements do not form pi bonds ^[3].  As there are 7 4f orbitals there can be up to 7 unpaired electrons in the lanthanide elements resulting in large magnetic moments measured for lanthanide compounds ^[4]. Finding the magnetic moment of lanthanide compounds is one method used to determine the 4f electron configurations and hence is a valuable tool in the investigation of chemical bonding in the compounds ^[3].

The electronic configuration for the structure of the lanthanide elements have the general formula, [Xe]6s²4fⁿ, where n is any integer from 1 to 14. The chemistry observed in the lanthanides is derived from their +3 oxidation states. All lanthanide elements form trivalent cations, Ln³⁺; their “chemistry is established by the ionic radius” which, decreases across the period from lanthanum to lutetium ^[6]. In the Ln(III) compounds the 6s and 1 4f electron are removed leaving the general formula of: [Xe]4f^m. Lanthanides exist primarily in their +3 oxidation state when in coordination complex form because the probability of finding 4f orbitals is very high at the nucleus, meaning the 4^th ionisation energy is very high and unfavourable. Ce³⁺ is able to “lose its single f electron”, forming Ce⁴⁺ cation with the stable electronic configuration of xenon. The Eu³⁺ cation can “gain an electron to form Eu²⁺” with an f⁷ configuration that has the enhanced stability of a half-filled shell. Other than these exceptions, none of the lanthanides are “stable in oxidation states” other than +3 in aqueous solution ^[6]. These observed states are very strongly electropositve making the lanthanide ions hard Lewis acids. The trivalent lanthanide ions, excluding lanthanum and lutetium, have unpaired f electrons-they are paramagnetic.

In “multi-electron atoms”, the decrease in ionic radius results from an increase in nuclear charge being offset by increasing electrostatic repulsion among electrons. As more electrons are added to the outer shells, the electrons in the inner shells “shield the outer electrons from nuclear charge”, making these outer electrons “experience a lower effective charge on the nucleus”. The “shielding effect exerted by the inner electrons decreases in the order s > p > d > f ”. Generally, as a certain subshell is filled in a period, atomic radius decreases. This effect is observed strongly in the case of lanthanides, as the 4f subshell which is filled across these elements has a poor effect at shielding the outer shell electrons. Hence, the shielding effect is “less able to counter the decrease in radius” caused by increasing nuclear charge. This leads to the phenomenon observed in lanthanides of lanthanide contraction. ^[3]
The lanthanide contraction is a trend in which the reduction in size of the Ln³⁺ cation can be observed from La³⁺ to Lu³⁺ can be explained by the inadequate shielding of the electrons in the 5s and 5p orbitals by those in the 4f orbitals, as a direct result from poor shielding of nuclear charge by the 4f electrons; causing the 6s electrons to be pulled towards the nucleus, thus “resulting in a smaller atomic radius” ^[3]. A size of the La³⁺ is measured as 103 pm reduces to a size of 86.1 pm for the Lu³⁺ cation across the period ^[3]. The term was coined by the Norwegian geochemist Victor Goldschmidt in his series “Geochemische Verteilungsgesetze der Elemente” ^[7]. Without the lanthanide contraction, the ability to chemically separate lanthanides would be quite difficult due to their similar chemical properties.

There is a low probability of the 4f electrons existing at the outer region of the atom or ion resulting in reduced effective overlap between the orbitals of a lanthanide ion and any ligand binding to it ^[8]. Therefore, lanthanide complexes generally have little or no “covalent character” and are not typically affected by orbital geometries. This lack of orbital interaction suggests that altering the metal usually has no effect on the complex, apart from its size . Lanthanide complexes are held together by “weak electrostatic forces” which are “omni-directional” and so, it is the ligands that dictate the symmetry and coordination of the complexes ^[9]. Steric factors have a large effect on the coordination of complexes, with “coordinative saturation of the metal balanced against inter-ligand repulsion”. This creates a wide range of coordination geometries and results in the “highly fluxional nature” of lanthanide complexes ^[10]. Furthermore, quick intramolecular and intermolecular “ligand exchange” will take place as there is no real energetic cause for a coordination complex to be “locked into a single geometry”, resulting in complexes interchanging between a large number of configurations ^[11].

Crystal field splitting in lanthanide ions is quite small and so less important than spin-orbit coupling in terms of energy levels. The magnetic moments of the trivalent lanthanide ions “deviate from spin-only values” due to strong spin-orbit coupling ^[12]. Russel-Saunders coupling can be applied to the lanthanide ions in which the interaction of the quantum numbers L and S is observed; the coupling of L and S results in the angular momentum, J.

Some electronic transitions are not permitted; however, this does not mean that such a transition will never occur, but that it is less likely to and the intensity, and hence, the molar absorption coefficient, of such an absorption band is very low. Whether transitions are allowed or forbidden, and to what degree they may be forbidden depends on the selection rules. Transitions of electrons between the f orbitals are forbidden by the Laporte rule, f-f transitions are said to be Laporte forbidden. The Laporte Selection Rule states: “In a molecule having center of symmetry, transitions between states of the same parity (symmetry with respect to a center of inversion) are forbidden” ^[15]. Laporte-allowed transitions involve Δl = ±1.
‘s ↔ p’, ‘p ↔ d’, ‘d ↔ f’ etc allowed (Δl = ±1)
‘s ↔ d’, ‘p ↔ f’ etc forbidden (Δl = ±2)
‘s ↔ s’, ‘p ↔ p’ , ‘d ↔ d’, ‘f ↔ f’ etc forbidden (Δl = 0)

The Spin Selection Rule states: The overall spin, S, of a complex must not change during an electronic transition, hence, ΔS = 0. To be spin allowed, a transition must involve no change in the spin state. This is because electromagnetic radiation usually cannot change the “relative orientation of an electron spin” ^[13]. For example, [Mn(H₂O)₆]²⁺ has a d⁵ configuration and is a high-spin coordination complex; therefore electronic transitions are not only Laporte-forbidden, but also spin-forbidden. This results in dilute solutions of Mn²⁺ complexes being colourless. This explains why the Ln³⁺ cations are pale in colour. The lanthanide cations have unique colours due to their different valence 4f electronic configurations giving rise to specific electronic absorption spectra.

Polyoxometalates are soluble anions that form a class of chemical species ranging between discrete anions such as [MoO₄]^2- and infinite lattice solid oxides such as Mo₃. They operate as successful ligands for many metal ions using an octahederal {MoO₆}^6- unit as a basic building block ^[14].

The “lack of orbital interactions” and lanthanide contraction results in a change in lanthanide size across the series but a similar chemistry observed between elements. Lanthanide metal oxides can be used to operate as heterogeneous catalysts in a number of industrial processes. In Gd³⁺ all the electrons have parallel spin; a property which benefits the use of gadolinium complexes as “contrast reagent in MRI scans”. As f-f transitions are Laporte-forbidden, when an electron has been excited by absorbing a photon of energy, decay to the ground state will be slow. This makes lanthanide ions useful in laser applications as it facilitates the achievement of population inversion. The Nd:YAG laser is commonly used in oncology and laser eye surgery; the acronym stands for: neodymium-doped yttrium aluminium garnet; Nd:Y₃Al₅O₁₂. Nd(III) “typically replaces a small amount” of the yttrium ions, due to their similar ionic sizes, which are present in the crystalline structure of the YAG. The neodymium ion “provides the lasing activity in the crystal”. The Nd:YAG laser was first presented by “J. E. Geusic et al. in 1964. Lanthanide coordination complexes operate as effective homogeneous catalysts as hard Lewis acids are able to “polarise bonds upon coordination” affecting the electrophilicity of the compounds. This can be observed in the Luche reduction, in which lanthanidechlorides, mainly CeCl₃, and sodium borohydride are used to reduce α,β-unsaturated ketones to allylic alcohols in ethanol. ^[15]

There are a number of key properties of metal cations that dictate their coordination chemistry have certain consequences on their interactions with the lacunary phosphomolybdate anion, [PMo₁₁O₃₉]^7-. Firstly, since the f-orbitals are held deep in the core of the electron cloud of the lanthanides they are not readily available for bonding so ionic/electrostatic interactions dominate. Therefore, the Ln³⁺ cations are very hard metal centres and readily bond to oxygen donor ligands, including the lacunary phosphomolybdate anions ^[5].

The basicity of the lanthanide (III) cations, a measure of the ability of an atom to lose electrons, is a property that follows the attraction between the cations and anions. The basicity series is as followed: series is the following:

La³⁺ > Ce³⁺ > Pr³⁺ > Nd³⁺ > Pm³⁺ > Sm³⁺ > Eu³⁺ > Gd³⁺ > Tb³⁺ > Dy³⁺ > Ho³⁺ > Er³⁺ > Tm³⁺ > Yb³⁺ > Lu³⁺

In essence, the basicity decreases as the atomic number increased. These trends can be followed by observing the solubility of the lanthanide (III) salts and the creation of complex species ^[18].

The first monolacunary Keggin complexes with lanthanide cations were prepared by Peacock and Weakley in 1971. Polyoxometalate derivatives of lanthanides are of great importance in the world of chemistry currently due to their “unique structural, chemical and electronic properties”.

The vacant sites provided by the polyoxometalates and the varying coordination gemoetries of the lanthanide cations allows “large oxometalate clusters” to be formed which demonstrate “photoluminescence activity” ^[16].
Lanthanide contraction results in an increased charge:radius ratio across the series. The increased effective charge on the Ln³⁺ ion therefore results in stronger interaction with ligands. The decrease in ionic radius across the lanthanide series results in shorter Ln-O bond lengths in the lacunary complexes formed across the series. Due to the shielded nature of the 4f orbitals in lanthanides, “coupling with molecular vibrations” is very weak. This results in absorption spectra of lanthanide ions being weak with short, narrow absorption bands ^[9].
The main aim of this experiment was to gain experience of the preparation and chemistry of lanthanide polyoxometalate complexes by means of control of speciation as a function of pH; to apply several complementary spectroscopic techniques including UV/Vis/nIR and IR to probe the metal-ligand interactions in both solid state and in solution. Furthermore, to gain an appreciation of the key properties of the trivalent lanthanide metal cations and the trends observed in their coordination chemistry.

 

3. Reaction Scheme

 

H3PMo12O40 (s) + H2O (l)  [PMo12O40 ]3- (aq) + 3H+  (aq) + H2O (aq)

Scheme 1: H3PMo12O40 dissolved in water

[PMo12O40 ]3-(aq)            Na2CO3           [PMo11O39 ]7- (aq) + [MoO4]4+  (aq)

Scheme 2: Preparation of the lacunary anion at pH 4.3

[PMo11O39 ]7- (aq)  +12H2O (l)          Na2CO3          [MoO4]2- (aq) + [PO4]3-  (aq) + Na2Mo10O31.12H2O (aq)

Scheme 3: pH raised to 5.5

2H3PMo12O40 (s)  + PrCl3.16H2O (s) + 11KCl (s)                              K11[Pr(PMo11O39)2].16H2O] (aq) + [MoO]4- (aq) + 2H+ (aq) + 4HCl (aq)

Scheme 4: Preparation of Pr lacunary phosphomolybdate complex

2H3PMo12O40 (s)  + GdCl3.16H2O (s) + 11KCl (s)                              K11[Gd(PMo11O39)2].16H2O] (aq) + [MoO]4- (aq) + 2H+ (aq) + 4HCl (aq)

Scheme 5: Preparation of Gd lacunary phosphomolybdate complex

 

4. Experimental Procedure

 

4.1 Table of amounts for reagents used

The masses of lanthanide chloride salts used for the stock solutions made for the analysis of electronic absorption and IR spectroscopy and hence the calculated concentrations of the corresponding solutions in 30 mL are given in Table 1, along with the calculated concentrations of phosphomolybdic acid hydrate, potassium chloride and sodium carbonate.

 

Table 1. Quantities of reagents

Reagent	Mass /g	Molecular mass / g mol-1	Moles / mol	Concentration / mol dm-3
PrCl3	0.124	247.27	0.0005	0.02
GdCl3	0.132	263.61	0.0005	0.02
H3PMo12O40	1.17	1825.3	0.0005	0.02
H3PMo12O40	4.68	1825.3	0.002	0.03
KCl	1.10	74.550	0.01	0.3
Na2CO3	0.11	105.96	0.001	0.1

Table 1: Table of results for the calculation of concentrations for the reagents used in the experiment

Example calculation of concentration for PrCl₃ salt:

moles mol=mass (g)molecular mass (g mol-1)

Equation 1: Equation for number of moles (mol) from mass (g) and molecular mass (g mol^-1)

 

moles =

0.124g247.27 g mol-1=0.0005 mol (1 SF)

 

Concentraion mol dm-3=moles (mol)volume (dm3)

Equation 2: Equation for calculation of concentration (mol dm^-3) from number of moles (mol) and volume (dm³)

 

(where 1 dm³ = 1 L)

concentration of 0.124 g of PrCl₃ salt in 30 ml (0.03 dm³)of solution

=

0.0005 mol0.03 dm3= 0.02 mol dm^-3 (2 SF).

There were three main aspects of this experiment:

1. The determination of the pH at which the lacunary phosphomolybdate anion, [PMo₁₁O₃₉]^7- was present in the highest concentration, relative to other phosophorous containing species in aqueous solution by solution IR spectroscopy.
2. The preparation of two lanthanide lacunary phosphomolybdate complexes of general formula, [Ln(PMo₁₁O₃₉)₂]^11-, in aqueous solution. Both [Pr(PMo₁₁O₃₉)₂]^11- and [Gd(PMo₁₁O₃₉)₂]^11- complexes were observed and their solution IR and UV/Vis/nIR spectra of the Pr and Gd complexes and subsequent crystallisation.
3. The collection of the rare earth lacunary phosphomolybdate crystalline solids of general formula, K₁₁[Ln(PMo₁₁O₃₉)₂].16H₂O where Ln = Pr and Gd. IR spectra of the solid products were recorded and analysed.

Part 1: Phosphomolybdate solution speciation- a solution IR study

In order to determine the optimum pH for the formation of [PMo₁₁O₃₉]^7- in solution IR spectroscopy was used. From analysis of the spectra, the P-O stretch that occurs between 1000-1100 cm^-1 in the tetrahedral [PMo₁₂O₄₀]^3- is split into 2 bands for the of [PMo₁₁O₃₉]^7- complex with point symmetry of C_s (pseudo C_3v).
Prior to recording any pH measurements the pH meter had to be calibrated using sufficient volumes of pH 4 and pH 7 buffer solutions ensuring the probes were immersed in the solutions to allow accurate measurements. Once calibrated, H₃PMo₁₂O₄₀, (1.17 g, 0.0005 mol, 0.02 mol dm^-3) was dissolved in 25 cm³ distilled H₂O. A clear yellow solution was observed with a measured pH of 1.01. A few drops of the solution were placed in a labelled vial and put aside in preparation for solution IR spectrum measurements. The solution was made more basic, increasing the pH to 4.3 by adding small volumes of 0.1 mol dm^-3 Na₂CO₃ dropwise using a plastic pipette carefully observing the pH changes as the solution was stirred using a magnetic stirrer bar. The 0.1 mol dm^-3 Na₂CO₃ solution was made up in portions of 10 mL using 0.11 g (0.001 mol) in 10 mL distilled H₂O. At pH 4.3 the solution remained a clear yellow colour but slightly paler, forming the [PMo₁₁O₃₉]^7- monovacant lacunary anion. Again, a few drops of the solution were placed in a labelled vial and put aside in preparation for solution IR spectrum measurements. Lastly, the solution pH was raised further to 5.5 using additional drops of 0.1 mol dm^-3 Na₂CO₃ added slowly over time whilst stirred, carefully monitoring the pH change to avoid overshooting the pH.

The labelled vials of [PMo₁₂O₄₀]^3- with pH of 1.01, [PMo₁₁O₃₉]^7- with pH 4.3 and [PO₄]^3- with pH 5.5 were analysed using solution IR spectroscopy, in which the P-O stretches were observed between 1000-1100 cm^-1, acknowledging bridging peaks from [PMo₁₂O₄₀]^3- to [PMo₁₁O₃₉]^7-  . Results observed are given in part 4.2 of this report.

Part 2: Synthesis of lanthanide (III) lacunary phosphomolybdate complexes K₁₁[Pr(PMo₁₁O₃₉)₂].16H₂O and K₁₁[Gd(PMo₁₁O₃₉)₂].16H₂O

Two lanthanides, Pr and Gd, were used to prepare and spectroscopically characterise two different lanthanide complexes. PrCl₃.16H₂O and GdCl₃.16H₂O lanthanide (III) chloride salts were used to prepare the lacunary phosphomolybdate complexes. First, a stock solution was prepared using 4.68 g (0.002 mol) of H₃PMo₁₂O₄₀ dissolved in 60 mL distilled H₂O and the pH of the solution was raised to the optimum pH for [PMo₁₁O₃₉]^7- formation, pH 4.3, using solid Na₂CO_3. The solid Na₂CO₃ was added slowly to the solution carefully observing pH raise from 1.01 to 4.3. Once the pH was raised to 4.3 the solution was divided into 2 aliquots of 30 mL for each lanthanide. 0.124 g (0.0005 mol) of Pr lanthanide salt and 0.132g (0.0005 mol) of Gd lanthanide salt were weighed out were added to the separate 30 mL solutions, and the reaction vessels were labelled. To the Pr lanthanide solution solid Na₂CO₃ was added to raise the pH from 3.6 to 4.3 and similarly for the Gd lanthanide solution to raise the pH from 3.73 to 4.3. 1.1 g (0.01 mol) KCl was accurately weighed out and added to each reaction vessel containing the lanthanide solutions and was left to stir for 45 minutes. After this time, 5 mL of each lanthanide solution was placed in a labelled vial and set aside for spectroscopic analysis. To the remaining 25 cm³ not used for spectroscopic analysis for each lanthanide solutions 12 cm³ acetone was added drop wise to each reaction vessel with stirring. Each solution was then filtered by vacuum filtration to remove any residual solid material. The reaction vessel was then covered with parafilm and stored in the freezer over night to enable crystalisation. By the following day a crystalline product, the lacunary complex, precipitated from each solution. The products were filtered by vacuum filtration and washed with acetone and left to air dry. The dried lacunary products were weighed and their yields recorded in table 2 in section 4.2 of this report.

 

 

Part 3: Spectroscopic characterisation of lanthanide (III) lacunary phosphomolybdate complexes K₁₁[Pr(PMo₁₁O₃₉)₂].16H₂O and K₁₁[Gd(PMo₁₁O₃₉)₂].16H₂O

Solution IR spectra were recorded for the lacunary phosphomolybdate complexes: K₁₁[Pr(PMo₁₁O₃₉)₂].16H₂O and K₁₁[Gd(PMo₁₁O₃₉)₂].16H₂O recording the splitting difference between the P-O stretches between 1000-1100 cm^-1  for each complex. The IR spectra of the solid products crystallised were also recorded observing a loss of strong, broad O-H peak from removal of water, making the P-O stretches appear more intense on both spectra.

UV/Vis analysis was carried out in the spectrum region: 200-900 nm recorded in 1 cm³ polystyrene cells. The main peaks were labelled with wavelength and absorbance values. Wavenumbers and extinction coefficients were calculated from equation 3 and equation 4 respectively. Results of any main peaks observed recorded in table 3 in section 4.3 of this report.

 

 

 

 

 

 

 

Calculations:

υ (cm-1)=1λ (cm)

Equation 3: Equation to convert wavelength in cm to wavenumber in cm^-1

 

Conversion of wavelength in nm to cm required first to then convert into wavenumber.

1nm = 10^-7 cm
E.g.: Pr complex:

Wavelength, _max = 593.6.0 nm, in cm = 593.6 nm x 10^-7nm cm = 0.00005936 cm

 (cm^-1) =

10.00005936 cm

 cm^-1 = 16770 cm^-1 (5 SF)

ε (dm3mol-1cm-1)=Ac(mol dm-3)l(cm)

 

Equation 4: Beer Lambert Law

Calculation of the extinction coefficient for the Pr complex using its measured absorption relies upon the Beer Lambert Law.

Example calculation: Pr complex;

Absorption = 0.0812, using the calculated concentration value from Table 1 of 0.02 mol dm^-3

ε dm3mol-1cm-1=0.08120.02mol dm-3 x 1(cm)

= 4.06

dm3mol-1cm-1

 

4.2 Yields 

 

 

Table 2. Calculated yields

Complex	Theoretical yield / g	Actual yield / g	Percentage yield / %
Pr	2.14	1.43	67
Gd	2.15	1.49	69

 

Calculations example Pr complex:

Theoretical yield = Mr (K₁₁[Pr(PMo₁₁O₃₉)₂].16H₂O) (g mol^-1)x moles of PrCl₃ (mol)

= 4279.81 g mol^-1 x 0.0005 mol

= 2.14 g

Percentage yield =

actual yield (g)theoretical yield (g) x 100%

=

1.43 g2.14 g x 100%

= 67%

 

 

4.3 UV/Vis/nIR Spectroscopy

 

Table 3. UV/Vis/nIR data

Complex	Absorption	Wavelength max /nm	Wavenumber max / cm-1	Extinction coefficient   /dm3 mol-1 cm-1
Pr complex	0.0812	593.6	16770	4.06
Gd complex	–	–	–	–

 

 

4.4 Infra Red Spectroscopy v_max (method/cm⁻¹)

 

Table 4. IR data of [PMo₁₂O₄₀]^3- at pH 1.1.

Frequency (cm−1)	Br/s/m/w	Functional group assignment	Expected Range / cm−1	Wavenumber / cm−1	Splitting difference / cm−1
	Br	O-H	3000-3300	3256.49	–
	S,w	P-O	1000-1100	1062.22	–

 

 

Table 5. IR data of [PMo₁₁O₃₉]^7- at pH 4.3

Frequency (cm−1)	Br/s/m/w	Functional group assignment	Expected Range / cm−1	Wavenumber / cm−1	Splitting difference / cm−1
	Br	O-H	3000-3300	3291.02	–
	S,w	P-O	1000-1100	1057.60	29.85
	S,w	P-O	1000-1100	1027.75

 

 

Table 6. IR data of [PO₄]^3- at pH 5.5

Frequency (cm−1)	Br/s/m/w	Functional group assignment	Expected Range / cm−1	Wavenumber / cm−1	Splitting difference / cm−1
	Br	O-H	3000-3300	3205.49	–
	S,w	P-O	1000-1100	1043.57	–

 

 

Table 7. IR data of K₁₁[Pr(PMo₁₁O₃₉)₂].16H₂O

Frequency (cm−1)	Br/s/m/w	Functional group assignment	Expected Range / cm−1	Wavenumber / cm−1	Splitting difference / cm−1
	Br	O-H	3000-3300	3219.47	–
	S,w	P-O	1000-1100	1069.90	36.99
	S,w	P-O	1000-1100	1032.91

 

Table 8. IR data of K₁₁[Pr(PMo₁₁O₃₉)₂]

Frequency (cm−1)	Br/s/m/w	Functional group assignment	Expected Range / cm−1	Wavenumber / cm−1	Splitting difference / cm−1
	w	P-O	1000-1100	1072.47	43.60
	w	P-O	1000-1100	1028.87

 

Table 9. IR data of K₁₁[Gd(PMo₁₁O₃₉)₂].16H₂O

Frequency (cm−1)	Br/s/m/w	Functional group assignment	Expected Range / cm−1	Wavenumber / cm−1	Splitting difference / cm−1
	Br	O-H	3000-3300	3398.47	–
	w	P-O	1000-1100	1061.41	31.53
	w	P-O	1000-1100	1035.88

 

 

 

 

Table 10. IR data of K₁₁[Gd(PMo₁₁O₃₉)₂]

Frequency (cm−1)	Br/s/m/w	Functional group assignment	Expected Range / cm−1	Wavenumber / cm−1	Splitting difference / cm−1
	w	P-O	1000-1100	1070.14	38.58
	w	P-O	1000-1100	1031.56

 

 

 

 

5. Discussion of Results

Andrew J. Gaunt et al synthesised lanthanide complexes with the general formula: (NH₄)₁₁[Ln(PMo₁₁O₃₉)₂].16H₂O in which the Ln³⁺ cation was sandwiched in between two lacunary anions with eight oxygen atoms coordinated to lanthanide centre ^[17]. It is known that as Ln³⁺ cations contract across the lanthanide series, the Ln-O bond distances decrease and the P-O splitting of the vibrational mode in the lacunary anion increases.

Lanthanoid contraction explains the decrease in the relative size of ions in the lanthanide series corresponding to increasing atomic number from lanthanum, 57, to lutetium,71. The number of electrons present in the 4f orbitals surrounding the nucleus increases across the series. The 4f electrons are poorly shielded by each other from the increased positive charge of the nucleus, so that the “effective nuclear charge” attracting each electron steadily increases along the series, resulting in successive reductions of the ionic radii. The Praseodymium ion, Pr³⁺, has an ionic radius of 101.3 pm and the Gadolinium ion, Gd³⁺, has an ionic radius of 93.8 pm ^[9].

The “typical oscillator strengths of f-f transitions” are as small as 10^-6 because they are Laporte forbidden ^[18]. Nevertheless, f-f transitions have useful optic properties because the absorption and emission spectra have peaks in visible, near infrared and near ultraviolet regions. These peaks observed are sharp even in crystal fields because “4f electrons are not very affected by surrounding environment due to shielding effect of 5s and 5p electrons” ^[8]. Investigation into the “osicillator strengths of f-f transitions” has derived the “semi-empirical theory”: the Judd-Oflet theory ^[19]. The Judd-Oflet theory has been employed in a large number of experiments to deduce the intensities of lanthanide systems. The theory itself is based on crystal field theory and angular momentum coupling, it can be considered that the essential properties of f-f transition usually come from the crystal field generated by the surrounding environment but as previously mentioned in this report, the 4f electrons aren’t greatly affected by their surrounding environment due to shielding from the “closed shell” 5s and 5p electrons. Resulting in the crystal-field splittings being smaller than the spin-orbit splittings.

Gruen et al. reported that the “oscillator strengths of hypersensitive transitions” in gaseous lanthanide trihalide molecules,(LnX₃),( PrCl₃ andGdCl₃ in this experiment) were much larger than those of Ln³⁺ in solutions or crystals, however their wavelengths remained similar. Several models were developed in order to explain hypersensitivity; including the dynamic-coupling  model which could qualitatively explain the oscillator strengths of hypersensitive transitions in LnX₃ molecules ^[20].

The electronic structures of lanthanide complexes were studied by Amberger and co-workers using UV/Vis/nIR spectroscopy with crystal field modeling in order to determine the 4f orbital splitting as a result of ligand interaction ^[13].

The f-f transitions give rise to sharp, narrow bands of weak intensities as they are laporte forbidden. With Xenon having an electron configuration of: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s² 4p⁶ 4d¹⁰ 5s² 5p⁶; Pr has the following electron configuration: (Xe)4f²6s² and Gd (Xe)4f⁷ 5d¹ 6s². This explains the observed UV/Vis/nIR spectra for the lanthanide complexes. No major peaks were observed in the Gd complex, this is because with a 4f⁷ configuration it has 7 unpaired electrons in the 7 f orbitals meaning f-f transitions are not only laporte forbidden but they are also spin forbidden. The Pr complex with 4f² configuration also has laporte forbidden f-f transitions but does have spin allowed f-f transitions. The observed wavelength on 593.6 nm for the Pr complex is consistent with f-f transitions typically around 600 nm. A low extinction coefficient calculated of 4.06 dm³ mol^-1 cm^-1 for the Pr complex signifies the reduced likelihood of f-f transitions occurring in the complex.

The speciation of phosphomolybdate complexes refers to the varying concentration of the types of ion as the pH of the phosphomolybdate solution changes. This follows the Henderson-Hasselbalch ^[21] equation which states: the ration of acid: conjugate base concentration changes as the difference between pH and pKa of the solution changes. In the first part of the experiment the solution speciation of phosphomolybdates was observed. When H₃PMo₁₂O₄₀ was dissolved in water a colour change from colourless to a clear yellow solution was observed with a pH of 1.1. The equation for the reaction:

H₃PMo₁₂O_{40 (s)} + H₂O _(l)  [PMo₁₂O₄₀ ]^3- _(aq) + 3H⁺  _(aq) + H₂O _(aq)

The solution IR spectrum recorded for the [PMo₁₂O₄₀ ]^3-  anion, revealed 2 distinct peaks: a broad O-H peak at 3256.49 cm^-1 and a weak, sharp P-O peak at 1062.22 cm^-1 . The one P-O stretch observed in the [PMo₁₂O₄₀ ]^3-  anion spectrum was split into two bands in the lacunary anion, [PMo₁₁O₃₉ ]^7-  when the pH was raised to 4.3 and a slightly paler clear yellow colour formed; results shown in table 5 in section 4.3 of this report. Peaks at 1057.60 cm^-1 and 1027.75 cm^-1 were measured for the two P-O stretches observed in the pseudo c_3v point group symmetry, with a splitting of 29.85 cm^-1. When the pH was raised to 5.5, the solution appeared an even paler shade of yellow and the two bands converged to one at 1043.57 cm^-1. Only one peak was observed at pH 5.5 because the point group of the molecule determines the number of IR active modes. At pH 5.5 the [MoO₄]^2- and [PO₄]^3- are hydrolysed from H₃PMo₁₂O₄₀ reverting to tetrahederal, Td point symmetry, which shows one IR active P-O mode, the two bands present in [PMo₁₁O₃₉ ]^7-  at pH 4.3 converge to one.

Dalton Transactions, 2005, claimed that complexation of a Ln (III) cation into the vacancy in lacunary anion increases the splitting between the P-O stretching modes. This is demonstrated by the lower P-O splitting observed in the phosphomolybdate lacunary anion of 29.85 cm^-1 shown in table 5, than the lanthanide complexes which displayed splitting of 43.60 cm^-1 shown in table 8 for the Pr complex and 38.58 cm^-1 shown in table 10 for the Gd complex. Both the [Pr(PMo₁₁O₃₉)₂]^11- and the [Gd(PMo₁₁O₃₉)₂]^11- anions can be deemed as stable in the solid state and in solution as the splitting for the individual lanthanide complexes are similar as shown in table 7-10 in section 4.4 of this report. In tables 7 and 9 we can observe the solution phases of the lanthanide complexes; K₁₁[Pr(PMo₁₁O₃₉)₂].16H₂O and K₁₁[Gd(PMo₁₁O₃₉)₂].16H₂O giving P-O stretch splitting values of 36.99 cm^-1 and 31.53 cm^-1 respectively. Tables 8 and 10 show the solid phases of the complexes K₁₁[Pr(PMo₁₁O₃₉)₂] and K₁₁[Gd(PMo₁₁O₃₉)₂] respectively. Upon the removal of water a slight increase in the splitting values for each lanthanide complex can be observed; 43.60 cm^-1 for Pr and 38.58 cm^-1 for Gd.

A primary aim of this experiment was to prepare two lanthanide lacunary phosphomolybdate complexes of general formula K₁₁[Ln(PMo₁₁O₃₉)₂]^11- . Table 2 in section 4.2 of this report states the achieved yields of crystalline solid products and the calculated percentage yields. A calculated percentage yield of 67% for the K₁₁[Pr(PMo₁₁O₃₉)₂].16H₂O complex and 69% for the K₁₁[Gd(PMo₁₁O₃₉)₂].16H₂O complex concludes successful synthesis was carried out and hence the methodology was not only reliable but carried out accurately.

 

 

6. Conclusions

Overall, the experiment succeeded in determining the pH at which [PMo₁₁O₃₉]^7- was present at its highest concentration relative to other phosphorous containing species in aqueous solution by solution IR spectroscopy observing speciation in the phosphomolybdate solution. Two lanthanide lacunary phosphomolybdate complexes, Pr and Gd were successfully synthesised in aqueous solution and crystalised to high percentage yields of 67% and 69% respectively. The subsequent spectroscopic analysis of the complexes synthesised in the experiment allowed the quantification of splitting observed in the phosphomolybdate lacunary anion and hence an insight into metal-ligand interactions and relationship between point group symmetry and IR active bands. Furthermore, an insight into the phenomenon of lanthanoid contraction was gained and understanding of how complexation of the Ln (III) cation to the vacancy in the lacunary anion increased the P-O splitting. Lanthanoid contraction could be used to explain the trends observed in the P-O splitting between the Pr and Gd complexes, relating the smaller Gd³⁺ cation to less splitting. These confirmations of literature claims confirm that not only were the complexes synthesised successfully but the analysis of the UV/Vis/nIR absorption spectra and IR spectra were carried out accurately and successfully, useful skills for future experiments. Research into phenomenon such as the lanthanide contraction enabled further understanding of the trends observed in the lanthanide complexes. The experiment, as a whole, was particularly new to the researcher and hence provided great insight into a new period of elements- the lanthanides.

7. Future Work

 

The outcomes of this study have concluded successful preparation of lanthanide complexes and their subsequent analysis based on the P-O splitting differences due to lanthanide contraction. For further development of this experiment, ligand to metal charge transfer transitions could be defined using molecular orbital theory; this would provide deeper understanding to the behavior of the lanthanide cations and the orbital overlaps resulting in the varying P-O and O-Mo stretches.

There are a number of excellent reports offering different aspects of rare earth magnetism in the  Handbook on the Physics and Chemistry of Rare Earths, edited by K.A. Gschneidner, Jr. and L. Eyring (North-Holland 1978), providing a detailed investigation into the magnetic properties of lanthanides and the consequences of the magnetic interactions. The employment of these sources could be used to further enhance magnetic understanding of the lanthanides and aid in development of future experiments. A factor that could be altered to improve the methodology in this project is the addition of solid Na₂CO₃ to H₃PMo₁₂O₄₀ in part 2 of the experiment. The addition of Na₂CO₃ dropwise in aqueous form enabled more control and observation of speciaition, whereas solid addition could result in an overshoot of pH.

 

8. References

 

[1] G.Theodore The Elements: A Visual Exploration of Every Known Atom in the Universe. New York: Black Dog & Leventhal Publishers. p. 240. 2009

[2] Anon, (2018). [online] Available at: https://chemistry.tutorvista.com/inorganic-chemistry/f-block-elements.html [Accessed 9 Mar. 2018].

[3] C. Simon (2006). Lanthanide and Actinide Chemistry. John Wiley & Sons Ltd

[4] B. D. Cullity, and C. D. Graham, (2011) Introduction to Magnetic Materials, John Wiley & Sons

[5] Bochkarev, Mikhail N.; Fedushkin, Igor L.; Fagin, Anatoly A.; Petrovskaya, Tatyana V.; Ziller, Joseph W.; Broomhall-Dillard, Randy N. R.; Evans, William J. (1997). “Synthesis and Structure of the First Molecular Thulium(II) Complex: [TmI2(MeOCH2CH2OMe)3]”. Angewandte Chemie International Edition in English. 36 (12): 133–135.

[6] MacDonald, Matthew R.; Bates, Jefferson E.; Ziller, Joseph W.; Furche, Filipp; Evans, William J. (3 July 2013). “Completing the Series of +2 Ions for the Lanthanide Elements: Synthesis of Molecular Complexes of Pr, Gd, Tb, and Lu”. Journal of the American Chemical Society. 135 (26): 9857–9868.

[7] Hakala, Reino W. (1952). “Letters”. Journal of Chemical Education. 29 (11): 581.

[8] Lukens, W., Speldrich, M., Yang, P., Duignan, T., Autschbach, J. and Kögerler, P. (2016). The roles of 4f- and 5f-orbitals in bonding: a magnetochemical, crystal field, density functional theory, and multi-reference wavefunction study. Dalton Transactions, 45(28), pp.11508-11521.

[9] Adamo, C. and Maldivi, P. (1997). Ionic versus covalent character in lanthanide complexes. A hybrid density functional study. Chemical Physics Letters, 268(1-2), pp.61-68.

[10] Burgess, J. (1978). Metal ions in solution. New York: Ellis Horwood

[11] Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. pp. 1230–1242.

[12] Robert J. Lancashire (2006). “Selection Rules for Electronic Spectra of Transition Metal Complexes”. University of the West Indies, Mona.

[13] Gouzerh, P.; Che, M. (2006). “From Scheele and Berzelius to Müller: polyoxometalates (POMs) revisited and the “missing link” between the bottom up and top down approaches”. L’Actualité Chimique. 298: 9.

[14] Levine, Albert K.; Palilla, Frank C. (1964). “A new, highly efficient red-emitting cathodoluminescent phosphor (YVO4:Eu) for color television”. Applied Physics Letters. 5 (6): 118.

[15] Mason, S., Peacock, R. and Stewart, B. (1975). Ligand-polarization contributions to the intensity of hypersensitive trivalent lanthanide transitions. Molecular Physics, 30(6), pp.1829-1841.

[16] Keggin, J. F. (1934). “The Structure and Formula of 12-Phosphotungstic Acid”. Proc. Roy. Soc. A. 144 (851): 75–100.

[17] Gaunt A. J.; May I.; Sarsfield M. J.; Collison D.; Helliwell M.; Denniss I. S. Dalton Trans. (2003)

[18] James W. Robinson (1996). Atomic Spectroscopy. MARCEL DEKKER Incorporated. pp. 26

[19] Walsh, Brian M. “Chapter 21: Judd-Ofelt theory: Principles and practices”. In Di Bartolo, B.; Forte, O. Advances in Spectroscopy for Lasers and Sensing. Springer Netherlands. pp. 403–433. Retrieved 18 November 2015.

[20] Mason, S., Peacock, R. and Stewart, B. (1975). Ligand-polarization contributions to the intensity of hypersensitive trivalent lanthanide transitions. Molecular Physics, 30(6), pp.1829-1841.

[21] Po, Henry N.; Senozan, N. M. (2001). “Henderson–Hasselbalch Equation: Its History and Limitations”. J. Chem. Educ. 78 (11): 1499–1503