f Block Elements

INTRODUCTION

These 14 elements are called the lanthanons, or acanthanides

They are characterized by the filling up of the antepenultimate 4f energy levels.

They are extremely similar to each other in properties, and until 1907 they were thought to be a single element.In the past they were called the rare earths.

ELECTRONIC STRUCTURE

Lanthanum (the d-block element preceding this series) has the electronic structure: xenon core 5d1 6s2.

It might be expected that the 14 elements from cerium to lutetium would be formed by adding 1, 2, 3, ... 14 electrons into the 4f level.

However, it is energetically favourable to move the single 5d electron into the 4f level in most of the elements, but not in the cases of Ce, Gd and Lu.

The reason why Gd has a 5d1 arrangement is that this leaves a half filled 4f level, which gives increased stability.

Lu has a 5d1 arrangement because the f shell is already full.

The lanthanides are characterized by the uniform (+III) oxidation state shown by all the metals.

They typically form compounds which are ionic and trivalent. The electronic structures of the ions are Ce3+ f1, Pr3+ f2, Nd3+ f3,... Lu3+ f14.

Consequently the 4f electrons do not take part in bonding.

They are neither removed to produce ions nor do they take any significant part in crystal field stabilization of complexes.

OXIDATION STATES

Thus the oxidation state (+III) is ionic and Ln3+ dominates the chemistry of these elements. The Ln2+ and Ln4+ ions that do occur are always less stable than Ln3+. (In this chapter the symbol Ln is used to denote any of the lanthanides.)

In just the same way as for other elements, the higher oxidation states occur in the fluorite and oxides, and the lower oxidation states occur in the other halides particularly bromides and iodides.

Oxidation numbers (+II) and (+IV) do occur particularly when they lead to:

1. a noble gas configuration. e.g. Ce4+ (f0)

2. a half filled f shell. e.g. Eu2+ and Tb4+ (f7)

3. a completely filled f level. e.g. Yb2+ (f14).

In addition (+II) and (+IV) states exist for elements that are close to these states.

Thus Sm2+ and Tm2+ occur with f6 and f13 arrangements and Pr2+ and Nd4+ have f1 and f2 arrangements.

The (+III) state is always the most common and the most stable.

The only (+IV) and (+II) states which have any aqueous chemistry are Ce4+, Sm2+, Eu2+ and –Yb2+.

The lanthanide elements resemble each other much more closely than a horizontal row of the transition elements.

This is because the lanthanides effectively have only one stable oxidation state, (+III).

SEPARATION OF THE LANTHANIDE ELEMENTS

The properties of metal ions are determined by their size and charge.

The lanthanides are all typically trivalent and are almost identical in size, and so their chemical properties are almost identical.

The separation of one lanthanide from another is an exceedingly difficult task, almost as difficult as the separation of isotopes of one element.

The classical methods of separation exploit slight differences in basic properties, stability or solubility. These are outlined below.

However, in recent years the only methods used are ion exchange and valency change.

1. Precipitation

With a limited amount of precipitating agent the substance with the lowest solubility is precipitated most rapidly and most completely.

Suppose hydroxyl ions are added to solution conyaining a mixture of Ln(NO3)3.

The weakest base Lu(OH)3 precipitated first, and the strongest base La(OH)3 is precipitate last.

The precipitate contains more of the elements at the right of the series.

Thus the solution contains more of the elements at the left of the series.

The precipitate can be filtered off. Only partial separation is effected, but the precipitate can be redissolved in HNO3 and the process repeated to obtain greater purity.

2. Complex formation

A mixture of lanthanide ions is treated with a complexing agent such as EDTA (ethylenediaminetetraacetic acid).

All the ions form complexes.

The ions at the right hand side of the lanthanide series such as Lu3+ form the strongest complexes as they have the smallest ions.

Oxalates of the lanthanides are insoluble.

However addition of oxalate ions to this solution does not give a precipitate since the Ln3+ ions are all complexed with EDTA.

If some acid is added to the solution, the least stable EDT A complexes are dissociated.

This releases ions at the left hand side of the series Ce3+, pr3+, Nd3+ which are immediately precipitated as the oxalates.

These are filtered off. Separation is not complete, so the oxalates are redissolved and the process repeated many times.

3. Ion exchange

This is the most important, the most rapid and most effective general method for the separation and purification of the lanthanides.

A solution of lanthanide ions is run down a column of synthetic ion-exchange resin such as Dowex-50.

This is a sulphonated polystyrene and contains the functional groups -SO3H.

The Ln3+ ions are absorbed onto the resin and replace the hydrogen atom on -SO3H.

The H+ ions produced are washed through the column.

Then the metal ions are eluted, that is are washed off the column in a selective manner.

The eluting agent is a complexing agent, for example a buffered solution of citric acid/ammonium citrate or a dilute solution of (NH4)3(H • EDTA) at pH 8.

Consider the citrate case.An equilibrium is set up:

As the citrate solution flows down the column Ln3+ ions are removed from the resin and form the citrate complex.

A little lower down the column the Ln3+ ions go back onto the resin.

As the citrate solution runs down the column the metal ions form complexes alternately with the resin and the citrate solution many times.

The metal ion gradually travels down the column, and eventually passes out of the bottom of the column as the citrate complex.

The smaller lanthanide ions such as Lu3+ form stronger complexes with the citrate ions than do the larger ions like La3+, Thus the smaller and heavier ions spend more time in solution, and less time on the colournn, and are thus eluted from the column first.

The different metal ions present separated into bands which pass down the column.

The progress of the bands may be followed spectroscopically by atomic fluorescence.

The solution leaving the column is collected by means of an automatic fraction collector in separate containers.

By this means the individual elements can be separated.

The metals may be precipitated as insoluble oxalates, and then heated to give the oxides.

The chromatographic process is analogous to carrying out many separations or many crystallizations, but the separation is carried out on a single column. By using a long ion-exchange column the elements may be obtained 99.9% pure with one pass.

CHEMICAL PROPERTIES OF (+III) COMPOUNDS

The heavier metals are less reactive than the lighter ones because they form a layer of oxide on the surface.

The chemical properties of the group are essentially the properties of trivalent ionic compounds.

The sum of the first three ionization energies varies with minima at La3+, Gd3+ and Lu3+ which are associated with attaining an empty, half full or full f shell. Maxima occur at Eu3+ and Yb3+ associated with breaking a half full or full shell.

The hydroxides Ln(OH)3 are precipitated as gelatinous precipitates by the addition of NH4OH to aqueous solutions.

These hydroxides are ionic and basic.

They are less basic than Ca(OH)2 but more basic than Al(OH)3 which is amphoteric.

The metals, oxides and hydroxides all dissolve in dilute acids, forming salts.

Ln(OH)3 are sufficiently basic to absorb CO2 from the air and form carbonates.

The basicity decreases as the ionic radius decreases from Ce to Lu.

Thus Ce(OH)3 is the most basic, arid Lu(OH)3, which is the least basic, is intermediate between scandium and yttrium in basic strength.

The metals react with H2, but often require heating up to 300-400°C to start the reaction.

The products are solids of formula LnH2.

Eu and Yb both have a tendency to form divalent compounds and EuH2 and YbH2 are salt-like hydrides and contain M2+ and two H-.

The others all-form hydrides LnH2 which are black metallic and conduct electricity.

These are better formulated as Ln3+. 2H- and an electron which occupies a conduction band. In addition Yb forms a nonstoichiometric compound approximating to YbH2.5.

OXIDATION STATE (+IV)

The only (+IV) lanthanide which exists in solution and has any aqueous chemistry is Ce4+.

It is rare to find 4+ ions in solution.

The high charge on the ion leads to it being heavily hydrated, and except in strongly acidic solutions the hydrated Ce4+ is hydrolysed, giving polymeric species and H+.

Ce(+IV) solutions are widely used as an oxidizing agent in volumetric analysis instead of KMnO4 and K2Cr2O7.

Aqueous cerium(IV) solutions can be prepared by oxidizing a Ce3+ solution with a very strong oxidizing agent such as ammonium peroxodisulphate(NH4)2S2O8.

Ce(+ IV) is also used in organic reactions, for example the oxidation of alcohols, aldehydes and ketones at the a-carbon atom.

It has a three-dimensional crystal structure with the metal at the centre of square antiprism.

A number of complexes are stable, for example ammonium cerium(IV) nitrate (NH4)2(Ce(NO3)6].

The crystal structure is unusual and contains bidentate NO3- groups.

The Ce atom has a coordination number of 12 and the shape is an icosahedron.

This structure is stable even in solution.

Two of the NO3- ions may be replaced by phosphine ligands Ph3PO, giving a neutral 10-coordinate complex (CeIV(NO3)4(Ph3PO)2].

The other (+IV) compounds are not stable in water and are known only as oxides, fluorides and a few fluoro complexes.

Thus PrO2, PrF4, Na2[PrF6], TbO2, TbF4, TbO2, DyF4 and Cs3(DyF7] are all known.

The elements Pr, Nd Tb, and Dy also form(+IV) states.

These are generally unstable, occur only as solids, and are found as fluorides or oxides which may be nonstoichiometric.

OXIDATION STATE (+II)

The only (+II) states which have any aqueous chemistry are Sm2+, Eu2+ and Yb2+.

The most stable divalent lanthanide is Eu2+.

This is stable in water, but the solution is strongly reducing.

EuIISO4 can be prepared by electrolysing Eu2III solutions, when the divalent sulphate is precipitated.

EuIICl2 can be made as a solid by reducing EuIIICl3 with H2.

2EuCl3 + H2 → 2EuCl2 + 2HCl

Aqueous Eu3+ solutions can be reduced by Mg, Zn, zinc amalgam or electrolytically to give Eu2+.

EuH2 is ionic and similar to CaH2. Eu(II) resembles Ca in several ways:

The insolubility of the sulphate and carbonate in water.

The insolubility of the dichloride in strong HCI.

The solubility of the metals in liquid NH3.

Yb2+ and Sm2+ can be prepared by electrolytic reduction of their trivalent ions in aqueous solution.

However, the Ln2+ ions are readily oxidized by air.

These two elements form hydroxides, carbonates, halides, sulphates and phosphates.

The states Nd(+II), Pm(+II), Sm(+II) and Gd(+II) are only found in solid dihalides LnCl2 and LnI2.

These dihalides can be made by reducing the trihalide with hydrogen, with the metal, or with sodium amalgam,

The dihalides such as LaI2 and NdI2 tend to be nonstoichiometric.

They show metallic conduction, and are better represented as La3+ + 2I- + electron,

A detailed study of the third ionization energy shows the stability of a half filled and completely filled shell.

The ionization energies also suggest that there may also be extra stability associated with a three quarters filled shell.

COLOUR AND SPECTRA

Many trivalent lanthanide ions are strikingly coloured both in the solid state and in aqueous solution, The colour seems to depend cm the number, of unpaired f electrons.

Elements with (n) f electrons often have a similar colour to those with (14 - n) f electrons. (See Table) However, the elements in other valency states do not all have colours similar to their isoelectronic 3+ counterparts (Table).

Colour arises because light of a particular wavelength is absorbed in the visible region.

The wavelength absorbed corresponds to the energy required to promote an electron to a higher energy level.

In the lanthanides spin orbit coupling is more important than crystal field splitting.

In the spectra of transition metals, crystal field splitting is of major importance.

All but one of the lanthanide ions show absorptions in the visible or near-

Table-Colours of Ln4+, Ln2+ and their isoelectronic Ln3+ counterparts

UV region of the spectrum.

The exception is Lu3+ which has a full f shell.

These colours arise from - transitions.

Strictly these transitions are Laporte forbidden (since the change in the subsidiary quantum number is zero).

Thus the colours are pale because they depend on relaxation of the rule.

The f orbitals are deep inside the atom.

Absorption spectra of lanthanide ions are useful both for the qualitative detection and the quantitative estimation of lanthanides.

Lanthanide elements are sometimes used as biological tracers for drugs in humans and animals.

This is because lanthanide elements can quite easily be followed in the body by spectroscopy, because their peaks are narrow and very characteristic.

Ce3+ and Yb3+ are colourless because they do not absorb in the visible region However, they show exceptionally strong absorption in the UV region, because of transitions from 4f to 5d. Absorption is very strong for two reasons. Since ∆l = l This is an allowed transition and so gives stronger absorption than forbidden f-f transitions.

Furthermore, promotion of electrons in these ions is easier than for other ions.

The electronic configuration of Ce3+ is f1 and Yb3+ is f 8 Loss of one electron gives the extra stability of an empty or half full shell. f-d peaks are broad, in contrast to the narrow f-f peaks.

Charge transfer spectra are possible due to the transfer of an electron from the ligand to the metal.

This is more probable if the metal is in a high oxidation state or the ligand has reducing properties.

Charge transfer usually produces intense colours.

The strong yellow colour of Ce4+ solutions arises from charge transfer rather than f-f spectra.

The blood red colour of Sm2+ is also due to charge transfer.

MAGNETIC PROPERTIES

La3+ and Ce3+ have an f° configuration, and Lu3+ has an f 14 configuration.

These have no unpaired electrons, and are diamagnetic.

All other f states contain unpaired electrons and are therefore paramagnetic.

The magnetic moment of transition elements may be calculated from the equation:

μ(S+L) is the magnetic moment in Bohr magnetons calculated using both the spin and orbital momentum contributions.

S is the resultant spin quantum number and L is the resultant orbital momentum quantum number.

For the first row transition elements, the orbital contribution is usually quenched out by interaction with the electric fields of the ligands in its environment.

Thus as a first approximation the magnetic moment can be calculated using the simple spin only formula. (μs is the spin only magnetic moment in Bohr magnetons. s S is the resultant spin quantum number and n is the number of unpaired electrons.)

This simple relationship works with La3+ (f0), and two of the lanthanides Gd3+ (f7) and Lu3+ (f14).

La3+ and Lu3+ have no unpaired electrons, n = 0 and μs = √(0(0 +2)) = 0

Gd3+ has seven unpaired electrons. n = 7 and μs = √(7(7 +2)) = √63 = 7.9BM

In the lanthanides the spin contribution S and orbital contribution L couple together to give a new quantum number J.

J = L-S When the shell is less than half full and J = L+S when the shell is more than half full

The magnetic moment μ is calculated in Bohr magnetons (BM) by:

μ = g√(J(J +1))

Where, g = 1(1/2) + (S(S+1)-L(L+1))/(2J(J+1))

Figure shows the calculated magnetic moments for the lanthanides using both the simple spin only formula and the coupled spin plus orbital momentum formula.

For most of the elements there is excellent agreement between the calculated values using the coupled spin + orbital momentum formula and experimental values measured at 300 K.

The range of experimental values are shown as bars.

The agreement for Eu3+ is poor, and that for Sm3+ is not very good.

The reason is that with Eu3+ the spin orbit coupling constant is only about 300 cm-1.

This means that the difference in energy between the ground state and the next state is small.

Thus the energy of thermal motion is sufficient to promote some electrons and partially populate the higher state.

Because of this the magnetic properties are not solely determined by the ground state configuration.

Measuring the magnetic moment at a low temperature prevents the population of higher energy levels.

The magnetic moment of Eu3+ at low temperature is close to zero as expected.

The unusual shape of the spin plus orbital motion curve arises because of Hund's third rule.

When the f level is less than half full the spin and orbital momenta contributions work in opposition (J= L - S).

When the f shell is more than half full they work together (J = L + S).

LANTHANIDE CONTRACTION

Covalent and ionic radii normally increase on descending a group in the periodic table due to the presence of extra filled shells of electrons.

On moving from left to right across a period, the covalent and ionic radii decrease.

This is because the extra orbital electrons incompletely shield the extra nuclear charge.

Thus all the electrons are pulled in closer.

The shielding effect of electrons decreases in the order s > p > d > f.

The contraction in size from one element to another is fairly small.

However, the additive effect over the 14 lanthanide elements from Ce to Lu is about 0.2 Å, and this is known as the lanthanide contraction.

The hardness, melting points and boiling points of the elements all increase from Ce to Lu.

This is because the attraction between the atoms increases as the size decreases.

The properties of an ion depend on its size and its charge.

The Ln3+ lanthanide ions change by only a small amount from one element to the next (Table 29.9), and their charge is the same, and so their chemical properties are very similar.

Since Lu3+ is the smallest ion it is the most heavily hydrated.

Though the lanthanides do not form complexes very extensively; since Lu3+ is the smallest ion the complexes formed by Lu3+ are the are the strongest.

La3+ and Ce3+ are the largest ions so La(OH)3 and Ce(OH)3 are the strongest bases.

The lanthanide contraction reduces the radii of the last four elements in the series below that for Y in the preceding transition series.

Since the size of the heavier lanthanide ions, particularly Dy3+ and Ho3+, are similar to that of Y3+ it follows that their chemical properties also very similar.

As a result the separation of these elements is very difficult.

Ionic radii depend on how many electrons are removed.

For simplicity, the radii of ions with the same charge are compared in Table.

A similar change in size across the series is observed if covalent radii are compared (Table).

Because of this contraction in size across the lanthanide series, the elements which follow in the third transition series are considerably smaller than would otherwise be expected.

The normal size increase Sc → Y→ La uisappears after the lanthanides.

Thus pairs of elements such as Zr/Hf, Nb/Ta and Mo/W are almost identical in size.

The close similarity of properties in such a pair makes chemical separation very difficult.

The sizes of the third row of transition elements are very similar to those of the second row of transition elements (see Table 29.10).

Thus the second and third rows of transition elements resemble each other more closely than do the first and second rows.

COMPLEXES

The lanthanide ions Ln3+ have a high charge, which favours the formation of complexes. However, the ions are rather large (1.03 - 0.86 Å) compared with the transition elements (Cr3+ = 0.615 Å, Fe3+ = 0.55 Å (low spin)) and consequently they do not form complexes very readily.

Complexes with amines are not formed in aqueous solution because water is a stronger ligand than the amine.

However, amine complexes can be made in nonaqueous solvents.

Very few stable complexes are formed with CO, CN- and organometallic groups.

This is in contrast to the transition metals.

The difference arises because the 4f orbitals are well shielded and are 'inside the atom'.

Thus they are not take part in π back bonding, whereas in the transition elements the d orbitals are involved in π bonding.

The most common and stable complexes are those with chelating oxygen ligands such as citric acid, oxalic acid, EDTA4- and acetylacetone.

These complexes frequently have high and variable coordination numbers, and water or solvent molecules are often attached to the central metal.

β-Diketone complexes of Eu3+ and Pr3+ dissolved in organic solvents are used as lanthanide shift reagents in nmr spectroscopy.

Coordination numbers below 6 are uncommon, and occur only with bulky ligands such as (2,6-dimethylphenyl)- and [N(SiMe3)2]-.

In contrast to the transition elements, the coordination number 6 is not common.

The most common coordination numbers are 7, 8 and 9 and these give a variety of stereochemistries.

Coordination numbers 10 and 12 occur with the larger (lighter) lanthanides and small chelating ligands NO3- and SO42-) (Table ).

Complexes with monodentate oxygen ligands are much less stable than the chelates, and tend to dissociate in aqueous solution.

There are hardly any complexes with nitrogen donor ligands except ethylenediamine and NCS-, and these are decomposed by water.

Fluoride complexes LnF2+ are formed particularly by the smaller ions, but chloride complexes are not formed in aqueous media or concentrated HCl.

This is an important distinction between the lanthanide and actinide groups.

Ce4+ is smaller and more highly charged, and [Ce(NO3)6]2- is formed in the non-aqueous solvent N2O4, and is 12-coordinate. Each NO3- uses two oxygen atoms to coordinate to the metal.

The lanthanides form no complex with π bonding ligands, and the lack of π bonding is attributed to the unavailability of the f orbitals for bonding.

It is difficult to explain the bonding in complexes with high coordination numbers.

If one s orbital, three p orbitals and all six d orbitals in the valency shell are used for bonding, this accounts for a maximum coordination number of 9.

The higher coordination numbers of 10 and 12 present a problem.

They imply either participation of f orbitals in bonding, or bond orders of less than one.

There are few organic compounds of the lanthanides. Alkyls and aryls can be made with lithium reagents in ether solution:

LnCl3 + 3LiR → LnR + 3LiCl

LnR3 + LiR → Li[LnR4] and [LnMe6]3-

Cyclopentadienyl compounds [Ln(C5H5)3], [Ln(C5H5)2CI) and [Ln(C5H5)Cl2] are known but are sensitive to water and air.