UNIT 11 CHEMISTRY OF d- ANDf- ELEMENTS

Structure 11.1 Introduction Objectives 11.2 Transition and Inner Transition Elements - An Introduction 11.3 IUPAC Nomenclature of 6d Transition Series Elements 11.4 .Electronic Configuration of d-Block and f-Block Elements Electronic Configurations of Transition Elements and Ions Electronic Configurations of and Elements 11.5 in Properties Atomic Radii and Ioaic Rad~i Melting and Boiling Points Enthalpies of Oxidation States Colour of the Complexes Magnetic Properties Catalytic Properties Formation of Complexes Formation of Interstitial Compounds (Interstitial Solid Solutions) and Alloys (Substitutional Solid Solutions) 11.6 Summary 11.7 Terminal Questions

11.1 INTRODUCTION

In last unit we have studies about the periodicity and representative elements. In this unit we will study the chemistry of d and f block elements. First we will study the IUPAC nomenclature of these elements then we will discuss the electronic configuration, periodicity, variation of size, melting and boiling points. We shall also study the , , electrode potential, oxidation sate of these elements in detail.

Objectives

After studying this unit, you should be able to: explain the IUPAC nomenclature of d and f block elements, describe the electronic configuration of d and f block elements, outline the general properties of these elements, and discuss the colour, magnetic complex formation catalytic properties.

1 1.2 TRANSITION AND INNER TRANSITION ELEMENTS - AN INTRODUCTION We already know that in the the elements are classified into four blocks; namely, s-block, p-block, d-block andfiblock, based on the name of that accepts the or differentiating . In the elements of d-block orfiblock the valence enters the penultimate (inner to the outermost) (n-l)d orbitals or ante- penultimate (third to the outermost) (n-2lforbitals respectively.

The d-block elements occupy position in between s-block andp-block elements and possess properties that are intermediate (or transitional) between elements of s- andp- block and are, therefore, called transition elem'ents. Thef-block elements are often Chemistry of Elements called as inner-transition elements because in these elements the differentiating electron enters thef -orbitals of an inner shell. The term transition is used here because they exhlbit transition element behaviour by exhibiting variable oxidation states, forming coloured ions and exhibiting paramagnetism. The (also called lanthanoids or lanthanons) are placed along with and (also known as actinoids or actinons) along with in the periodic table. Customarily, they are listed separately in two series at the bottom of the periodic table in order to avoid it being excessively wide.

Presently most chemists consider as transition elements, only those elements that, either as free (or neutral atoms) or in any of their common oxidation states, have partly filled d-orbitals. According to this definition, the elements , and are excluded from the list of transition elements, as they possess completely filled d- orbitals in their neutral atoms and in their common oxidation states (~'3.

These large numbers of transition elements are further classified into four series of elements- first transition series, second transition series, third transition series and fourth transition series according to the filling of 3d, 4d, 5d and 6d atomic orbitals. The elements from to constitute the first transition series while a of nine elements from to constitute the second transition series. The third transition series begins with and goes up to . The fourth transition series (also called super heavy elements) starts with element having 104 and goes up to 1 12. These elements have been synthesized by artificial nuclear reactions and are very unstable with respect to a-decay or spontaneous fission with very short half- lives. Thef-block elements comprise two series of elements- the lanthanide series and the actinide series according to the filling of 4f and Sf orbitals, respectively. A group of fourteen elements following lanthanum in the periodic table are called lanthanides. Si~nilarly,the fourteen elements following actinium in the periodic table are termed as actinide elements. The elements following have been produced artificially and are collectively called transuranium elements.

------11.3 IUPAC NOMENCLATURE OF 6d TRANSITION SERIES ELEMENTS

As we already know each element in the periodic table have been given definite names and their symbols are derived by either taking the first alphabet, or by taking first alphabet and one more alphabet from the name of the elements. But the 4d transition series elements are given special names and symbols according to TUPAC rules, which are given below: 1: The name of the element is directly derived from the atomic number of the element as per the following numerical roots: 0 1 2 3 4 5 6 7 8 9 nil un bi tri quad pent hex sept oct enn

2. These roots are written together in order of appearance of the digits in the atomic number and terminated by -iurn. If enn comes before nil, the last n of enn is elided. Similarly, the final i of bi and tri is omitted when it occurs before -ium.

3. Taking the first letter of the initial roots that make up the number derives the symbols of the elements.

~hus,the name of the elements with atomic numbi 104 will be unnilquadium and the symbol will be Unq. The name and symbol of elen :~tshaving atomic numbers 104-112 are given in Table 1 1.1. . Table 11.1: IUPAC names and symbols of fourth transition series elements. Chemistry of d- and .f-Block Elements Atomic Name of Symbol Atomic Name of Symbol Number Element Number Element

104 Unnilquadium Unq 109 Unnilennium Une 105 Unnilpentium Unp 1 10 Ununnilium Uun 106 Unnilhexium Unh 111 Unununium Uuu 107 Unnilseptium Uns 112 Ununbium Uub 108 Unniloctium Uno

It is pertinent to note that these IUPAC names are assigned to elements only temporarily till the claimls about their synthesis are confirmed after which they are assigned proper names and symbols. Elements with atomic numbers 109 have already been assigned proper names. Recently the element with atomic number 110 has been named as (Ds), on 16Ihaugust, 2003 at 42"d General Assembly of IUPAC in Ottawa, Canada.

SAQ 1 a) Why zinc, cadmium and mercury are not considered as transition elements?

b) Write the rCTPAC name and symbol of unknown elements with atomic number 118, 120and 150.

11.4 ELECTRONIC CONFIGURATION OF I-BLOCK ANDfiBLOCK ELEMENTS

In the preceding section we discussed the position of transition elements and the inner transition elements in the periodic table and the rules for naming of elements of fourth transition series. We shall now discuss the ground state electronic configurations of transition and inner transition elements and their ions.

11.4.1 Electronic Configurations of Transition Elements and Ions We already know that the electronic bon figuration of is id, 2s22p6,3~~3~'. In the atom of next element the differentiating electron enters the 4s level. The 4s level becomes filled at the element , which has the electronic configuration [A-]4s2.In the atoms of successive elements from scandium to zinc, differentiating electrons enter 3d level instead of 4p. The electronic configurations of the atoms of second and third series transition elements follow similar pattern. In atoms of 4d series of transition elements, after filling up of the 5s level at [ICr]5s2 the differentiating electrons enter 4d level instead of 5p. Similarly, in atoms of 5d transition series elements the 6s level is filled at and the filling of 5d level begins at lanthanum which intempted by fourteen elements in which the electron enter 4f level (i.e. lanthanides), resumes at hafnium. The resulting electronic configurations of the atoms of transition elements are given in Table 1 1.2. The electronic configuration of the Chemistry of Elements ions can be obtained by removing first the outer ns electrons and then (n-l)d electron of the atom until the number of removed electrons become equal to the charge on the ion. For example, co3+will have electronic configuration as [Ar]3d6 and ~i~'as [Ar]3d, etc.

Table 11.2: Electronic configurations of the free atoms of transition elements.

it transition series Second transition series Element Free atom Element Free atom F [Ar]3d14s' Y [ICr]4d15s2 [Ar]3d24s' Zr [Kr]4d25s2 [Ar]3d34s' Nb [Kr]4# 5s' [Ar]3ds 4s' Mo [Kr]4ds 5s'

[Ar]3ds 4s' b Tc [Kr]4d6 5s' [Ar]3d6 4s2 Ru [Kr]4d75s' [Ar]3d74s2 Rh [Kr]4d5s1 [~r]3d4s~ Pd [Kr14d'~5s' [Ar]3d1° 4s' Ag [Kr14d'~5s' [~r]3d"4s2 Cd [~r]4d'O52

Third transition series I ~ourthtransitionseries I Element Free atom Element Free atom La rxel5d1 6s2 Unq [Rn]5f4 6d7s2

Hf ,U~P [Rn]5f4 6d3 7s' Ta unh lRn15f 6# 7s' W Uns 687s2 Re Uno [Rn]5f4 6d6 7s' 0s Une [Rn]5f4 6d7 7s' Ir Unn [Rn]5f4 627s' Pt Uuu [Rn]5f 6d7s2 Au Uub lRn15f 6dI0 7s'

Now you may be wondering as to why the ns level is filled first before the (n-l)d or the (n-22flevels and then later why (n-l)d or (n-22flevels are filled prior to np level. It is the radial dependence of the d-orbitals that is responsible for this order of filling of electrons in these elements. The stability of an electron in various atomic orbitals can be evaluated by comparison of radial probability functions 4n22 T 2 . The plot of radial distribution functions for 3d and 4s electrons is shown in Fig. 1 1.1.

Fig. 11.1: Radial probability functions for 3d and 4s orbitals In atom 26 A comparison of the plots for 3d and 4s orbitals reveals that for 4s orbitals the electron Chemistry of d- and humps are closer to the origin (i.e. nucleus), and well inside the hump of 3d .f-Block Elements probability function. This suggests that 4s electrons penetrates close into the argon core and spends an appreciable portion of its time close to the nucleus and hence experiences a greater average nuclear charge than that felt by 3d electrons. Hence (n-l)d orbitals is higher in energy than the ns orbital. Consequently, in potassium and calcium the nineteenth and twentieth electrons enter the 4s level rather than 3d level. As these two electrons are added, the nuclear charge increases by two units. Since the 3d orbitals penetrate the 4s orbitals strongly, it experiences a marked increase in effective nuclear charge and its energy falls abruptly below the 4p level to about the level of the 4s orbital. The next electron therefore enters the 3d orbital before the 4p. With an increase in atomic number the 3d probability functions maximum progressively moves closer to the nucleus and there is a continuous drop in its energy below that of the 4p level. Thus the next nine electrons enter the 3d level from to zinc. The change in the relative energies of various orbitals with atomic number is given in Fig. 11.2.

Fig. 11.2: The change in relative energies of various atomic orbitals with increasing atomic number in free atoms. I0 2 Chemistry of Elements When 3d level is filled at zhc, with the configuration [Ar]3d 4s , the level next in energy is 4p orbital which gets successively filled up to . Thus at krypton we have the configuration [Ar]3d10 4s2 4p6. Since the 5s level experiences more effective nuclear charge than does the 4d level in the next two elements following krypton, the 4d levels are higher in energy than the 5s level. The next two successive electrons enter the 5s orbital. The addition of these two electrons increases the nuclear charge by two units. Since the 5s electrons shield the 4d orbitals poorly, it experiences an increased effective 1 nuclear charge and its energy falls sharply below that of 5p orbitals. Hence the next electron goes into the 4d level which is filled at Ag with configuration [~r]4d'O5s~.The filling of next lowest energy level 5p is completed at .

At xenon the next available orbitals are 4f, 5d, 6s and 6p orbitals. The 6s, 6p and 5d orbitals penetrate the Xe core more strongly than does the 4f orbitals. Hence the 6s, 6p and 5d levels all drop beIow the 4f level in energy and the next lower energy level is the 6s. The next two electrons therefore enter the 6s orbitaldat caesiurn and barium, respectively. Both the 4f and 5d are strongly affected by increase in effective nuclear charge due to scarce shielding of 4f and 5d orbitals by 6s electrons and suffer a steep drop in energy below the 6p level, and 4f and 5d levels are all of about the same energy. The next two electrons in the elements lanthanum and enter 5d level, but the next twelve electrons enter the 4f orbitals. When the 4f shell is filled at , with the configuration [xe]4f4 586s2, the next lowest level 5d is successively filled fiom through mercury, then the 6p levels are filled to give the configuration of , [~e]4f5d'~6~*6~~.

The Sf orbitals are relatively non-penetrating and do not significantly drop in energy than do the 7s and 7p orbitals. Hence the next two electrons enter the 7s level. Again both the 5f and 6d orbitals penetrate the 7s orbitals very considerably and thus suffer a steep drop in energy below the 7p orbital. The Sf and 6d levels are very close in energy and the next entering electron enters 6d and/or Sf orbitals. The Sf orbitals probably become more stable in heavier actinides and are therefore successively filled.

It may be noted that the configurations of and copper (belonging to first transition series), and silver (second transition series) and gold (third transition series) are written as [~r]3d64s',[Ar]3di04s', [~r]4&5s',[~r]4d"5s' and [Xe]4f5d1°6s', respectively. This is because these configurations have some additional stability compared to their predicted configurations. The addtional stability is associated with half-filled subshell and completely filled configurations.Half-filled and completely filled orbitals have exchange energies due to inter-electronic interactions that are greater than the exchange energies associated with other configurations. This exchange energy arises from the indistinguishability of electrons and their spins. This energy is responsible for the stability of half-filled and completely filled subshell and is the main driving force for these configurationsto take an electron out of turn to achieve or maintain these configurations. Also these configurations provide the most symmetrical or uniform distribution and minimum inter-electronic repulsion.

If the positions of two electrons with parallel spins are exchanged in a subshell the decrease in energy per electron set is called exchange energy. It should be noted that the parallel and anti-parallel electrons act as independent sets. The exchange energy (E,,) for any configuration is proportional to tlie total number of pairs of electrons with parallel spins (P) in the orbital i.e. Ee, = KxP,where K is a constant and P = the combination nC2where n is the number of electrons with parallel spins. Thus P can have the values: Let us compare the stabilization due to exchange energy for the actual ground state Chemistry of d- and configuration with that for the predicated configuration of chromium. +Block Elements

Since the electrons in the 4s orbitals of these two configurations do not constitute any pair with parallel spins (exchange pair) they do not contribute to the exchange energy. The five unpaired 3d electrons in the actual configuration can make ten pairs of electrons with parallel spins and thus contribute 10K towards exchange energy. The exchange energy of the second configuration ds2from four electrons with parallel spins is 6K. This gain of 4K in exchange energy imparts additional stability and favours the 3d54s1configuration for chromium. But it should be remembered that this configuration is attained by moving an electron fiom 4s orbital to 3d orbital and there is some loss of energy (equal to promotional energy) in the process. Since the difference in energy between 4s and 3d orbitals is small this loss in energy is less than the gain in exchange energy and thus the configuration 324s' is favoured.

Similarly, we can compare the exchange energies for the two possible configurations 3d104s'and 384s' for copper.

In the actual configuration we have two sets of five electrons with parallel spins. Each of these two sets of electrons constitutes ten pairs of electrons with parallel spins and contributes 20K (10K each) towards exchange energy. But the alternative configuration has two sets of electrons with parallel spins- a set of five electrons with upward spins and another set of four electrons with downward spins. These two sets of electrons constitute ten and six electron pairs and thus contribute a total of 16K (10K and 6K) towards exchange energy. Thus in the actual configuration 3dI04s' the net gain in exchange energy is 4K. Again there is some loss of energy in promotion of an electron from 4s to 3d orbital and also in the pairing of electrons. In case of copper the gain in exchange energy is more than loss of energy and therefore the configuration 3d1°4s' is favoured over 384s'.

The exchange energy is responsible for the stability of half-filled and filled configurations for molybdenum (4d5s1), (4dI05s@), silver (4d1°5s'), gold (5dI06s'), (4f) and lutetium (4f4). In the configurations of the atoms of some second and third transition series elements such as (4d5s1), (4d75s'), (4d5s1)and (584s') both the inter-electronic and nuclear- electronic forces play important part in determ.@ng these configurations.

It should be noted that though in transition elements the ns orbital is filled before the (n-l)d orbital, it can not be assumed that s orbitals are more stable than d orbitals. In fact, when elements of the first transition series form ions it is always the s electrons that are lost first. This is because removal of an electron from any 3d transition element increases the effective nuclear charge on the 3d electrons over that of any 4s electron. Consequently, the 3d orbitals are expected to drop significantly in energy well below the 4s orbitals. The stability of an electronic configuration is, therefore, largely determined by the net effect of many forces such as nuclear electronic attraction, shielding of one electron by several other electrons fiom the nuclear charge, inter- electronic repulsion and exchange energy.

When the Sf level is filled at in actinide series the gradual filling of 6d orbitals starts with element having atomic number 104 and continues up to atomic try of Elements number 1 12. The electronic configurations of the elements of 6d transition series are given in Table 1 1.2.

11.4.2 Electronic Configurations of Lanthanide and Actinide Elements , In the preceding section we have learnt about the order of filling of various available atomic orbitals in the atoms of transition elements as well as lanthanide and actinide elements.

The element lanthanuni;mmediately preceding the lanthanide elements has the electronic configuration [xe]5d1 6s2. There is a sudden decrease in energy of 4f orbital at the next element cerium and the next-entering electron enters the 4f orbital. In subsequent thirteen elements the decrease in energy of 4f orbitals is so much that no electron enters 5d orbital and the 4f level is successively filled and the filling of 5d level is resumed only at hahium. The electron enters the 5d orbitals in case of gadolinium and lutetium to impart thedstable half-filled and completely filled configurations. As the 5d and 4f orbitals remain very close in energy, electrons readily shuttle between / them. But whether the electron remains in d orbital or not is of little consequence since in the most common + the d electron (if present) is removed along with two s electrons and tfe electronic configuration of M~'cations varies regularly from [~e]4f(ince3') to [xe]4f4(in LU~~.The ic configurations of the free atoms of lanthanide and actinide elements are dete their complicated atomic spectra, and are given in Table 11.3.

Table 11.3: Electronic configurationyd lanthanide and actinide elements. /

nfiguration outside

It is difficult to assign a precise configuration to the actinide elements. This is because the difference in energy between Sf and 6d orbitals in the beginning of the series is less than that between the 4f and 5d orbitals in the lanthanidb. Thus even the small amount of energy released during bond formation is sufficient enough to shift an electron from Sf to 6d. After actinium the energy of Sf orbitals falls below that of 6d orbitals. But in early actinide elements (Th -Np) the Sf and 6d orbitals are very close in energy and electron can occupy any ~r~sometimesboth of these levels. In and subsequent elements the Sf levels drop well below the 6d level so the electrons preferably fill the former. Thus the configuration of the atoms of actinide elements actually fluctuates between 5f7s2 and'5f'6d'7s2. The ground state electronic configuration of actinium [h]6d'7s2is similar to that of lanthanum and indeed these two elements possess similar I "I, chemical properties. Since in both lanthanide and actinide elements the f orbitals are successively filled up, \ I Chemistry of d- and these elements bear close relationship in 'their properties. However, there are some .f-Block Elements I significant differences also. The lanthanides invariably form ionic compounds and 4f ! electrons are little involved in bonding, whereas in actinides Sf orbitals are involved in 1 bond formation and there is some covalent contribution to . This is mainly because Sf orbitals have greater spatial extension relative to the 7s and 7p orbitals. Since in actinide elements 5f, 6d, 7s and 7p levels are close in energy particularly for uranium- elements and Sf orbitals have greater spatial extension, actinide elements can involve any or all of them in bonding. It is because of this reason that actinide elements have greater complex formation tendency as compared to ltmthanide elements. Thus actinides, unlike lanthanides, form complexes even with n-bonding ligands. Because of small difference in energies of these orbitals (even less than the binding energies) the electronic configuration of actinide ions varies with compounds and nature of ligands. It is therefore impossible to indicate which orbitals are involved in bonding or whether the bonding is ionic or covalent.

SAQ 2 a) Which of the two orbitals, 5d and 4f has higher energy at (i) lanthanum, and (ii) cerium?

b) Write the outer electronic configurations of the following ions: (i) cr5' ...... (ii) w3+...... (iii) ~b~'...... (iv) LU" ......

(v). . ~m~'...... (vi) N~~+......

SAQ 3 Which of the two ions, Fe3+andFez+ is more stable?

SAQ 4. a) Why is the filling of 4f orbital is not regular in the lanthanide elements?

b) In what ways the electronic configuration of actinides is different fiom that of lanthanides. Chemistry of Elements 11.5 PERIODIC TRENDS IN PROPERTIES

In the preceding sections we have studied the position and the electronic configuration of both d-block andf-block elements. These elements having incomplete d orf sub- shells exhibit many properties that depend primarily on their electronic configurations. Thus the elements of the d-block andf-block, like main group elements, show periodicity in their properties based on electronic configuration. The transition elements in addition to the horizontal similarity (along a ) show a vertical similarity (along a group) in their physical and chemical properties. The horizontal similarity is more pronaunced in the elements of first transition series, whereas in second and third transition series the vertical similarity predominates. Thus the chemistry of the elements of 3d series differs considerably from those of 4d and 5d series elements, which are quite similar to each other. The lanthanide elements show much more close horizontal similarity as compared to the d- block elements. There are many similarities between lanthanides and the later actinide elements. We shall now see how these properties vary from one group to the other and from one period to another.

11.5.1 Atomic and Ionic Radii Since it is not possible to measure the radius of an isolated atom or ion, the atomic radius or is deduced from the measured intemuclear distances in gaseous or in ionic . These calculated radii depend on bond order, bond'type (whether covalent, ionic or metallic), oxidation states of the bonded atoms or ions and and molecular structure. Thus we have several kinds of radii such as atomic (either covalent or metallic) and ionic.

The atomic radii of the elements of 3d, 4d and 5d transition series given in Table 11.4 indicate a general decrease from left to right across a period.

Table 11.4: Atomic radii of 3d, 4d and 5d transition elements in pm.

3d Atomic no. 2 1 22 23 24 25 26 27 28 29 30 series Element Sc Ti V Cr Mn Fe CO Ni Cu Zn Atomic Radius 144 132 122 118 117 117 116 115 117 125

4d Atomic no. 39 40 41 42 44 44 45 46 47 48 series Element Y Zr Nb Mo Ru Ru Rh Pd Ag Cd Atomic Radius 162 145 134 130 125 125 125 128 134 144

5d Atomicno. 57 58-71 72 73 74 76 76 77 78 79 80 series Element La Ce-Lu Hf Ta W 0s 0s Ir Pt AU Hg AtOmicRadius 169 165- 144 134 130 126 126 127 130 134 147 156

Across a row of transition elements the addition of electrons to (n-l)d orbitals is accompanied by a corresponding increase in nuclear charge. Due to imperfect screening of the outer ns electrons by each other and by the inner (n-l)d electrons the effective nuclear charge experienced by the ns electrons increases. The outer ns electrons become mote and more tightly bound and hence the atomic radius decreases. It is important to mention that shielding of the outer ns electron(s) by (n-l)d electron(s) is more efficient than the shielding of a ns electron by another ns electron (or that of a np electron by another np electron). The contraction in atomic radii from scandium through copper causes a reduction in atomic size of the succeedingp block elements, to krypton. As a consequence gallium and are almost the same size as and respectiveljr and the chemical properties of aluminium and gallium and that of germanium and silicon are almost similar. This is because of the cun~ulativeeffect of the poor shielding of (n-1)d electrons and increased effective Chemistry of d- and nuclear charge felt by the outer IIS$'-~ electrons. .f-Block Elements

Similarly, the that occurs from lanthanum to lutenium due to the filling of 4f orbitals has important consequences in the third transition series, hafnium through gold. Thus the atomic and ionic radii of the 5d transition elements are similar to the corresponding elements of the 4d transition series from to silver. In 3d transition elements, due to increase in effective nuclear charge in the beginning of the series, the atomic radii gradually decrease from scandium up to chromium but from chroinium to copper it is nearly the same. As the number of d electrons increases, the screening effect also increases. In the midway onwards of the series the effective nuclear charge and the screening effect become nearly equal and thus, there is no change in atomic radii from chromium to copper. However, at the end of the series the electron-electron repulsion among (n-l)d electrons become predominant and thus atonlic radius increases in zinc.

On descending a group, a small increase in atomic radius occurs between the first transition series and the second transition series. However, the elements of third transition series, except lanthanum, have nearly the same radii as elements of second series due to lanthanide contraction. The ionic radii follow the same trend as the atomic radii.

The metallic and ionic radii of lanthanide and actinide elements are given in Table 11.5.

Table 11.5: Metallic and ionic radii of lanthanide and actinide elements,

From the radii given in the table 11.5 it is seen that lanthanides are very close in size, with a small but regular decrease observed from lanthanum to lutetium. A similar gradation is found in the atomic radii and ionic radii of the actinide elements. The small decrease in atomic radius of these elements across a period is because of better screening effect of the additional inner ~lectronsfrom one element to the next on the size determining outer electrons from the nucleus. The rate of decrease in atomic radius in the lanthanide and actinide series is even less than that in the transition series elements. This steady decrease is not so regular and the first large decrease occurs with addition of first f electron and then after3 configuration between gadolinium and . These slow decreases in sizes from lantharium to lutetium and from actinium to are respectively called lanthanide contraction and actinide contraction. The decrease in size arises from the slow increase in effective nuclear charge with the

addition off electrons due to their imperfect . The jump in contraction , between consecutive elements in actinides is greater than in lanthanides. Chemistry of Elements The similarity in the ionic radii of both series to similarities in many of their chemical properties. The ionic radii of tripositive lanthanide ions are comparatively larger than those of the transition elements. The metallic radii of and ytterbium are exceptionally high because of difference in their . While each lanthanide contributes three electrons for metallic bonding Eu and Yb contribute only two electrons thus leaving half-filled and completely filled f orbitals. For cerium the smaller value is bkcause of the presence of cerium ions in oxidation state greater

thah 3+. As a result of lanthanide contraction, the normal increase in size from Sc -+ Y

-+ La disappears after the lanthanides, and pairs of elements such as Zr and Hf, Nb:and Ta, Mo and W, Ru and Os, etc, possess nearly similar sizes (Table 11.4). The properties of these elements are therefore very similar. The similarities in properties within these pairs make their separation extremely difficult. Thus, due to lanthanide contraction, the elements of 4d and 5d transition series resemble each other much more closely than do the elements of 3d and 4d transition elements. However, the effect of lanthanide contraction vanishes towards the right of the d- series.

11.5.2 Melting Points and Boiling Points The melting and boiling points of transition elements are generally very high. The melting points of the elements are correlated with the metallic bonding energy that is measured by enthalpies of atomization. with low enthalpies ofatomization are generally soft and have low melting points while those with high enthalpies of atomization are hard and have high melting points. The bonding energy of a increases with the number of unpaired electrons per atom available for bonding by delocalization. The enthalpy of atomization increases from left to right and becomes maximum with the transition metals where (n-l)d electrons are involved in bonding. Thus, there is a jump of about 700 K in the melting point between calcium (1 1 12 K) and scandium (1812 K). However, near the middle of the transition elements particularly 4d (Nb-Ru) and 5d (Hf-Ir) series the cohesive energies are maximum resulting in higher melting points. As we move towards the right of chromium, molybdenum and , - the d electrons become paired and the bonding is due only to s andp electrons resulting in a decrease in melting points. The metals of the second half of the transition senes with the beginning of electron pairing have fewer numbers of d electrons for bonding and thus the melting points decrease. However, the melting point of decreases sharply with respect to or although it has fiveunpaired electrons. Thus zinc and cadmium with no unpaired d electrons available for bonding are soft and have low melting points. The melting and boiling points of transition metals are given in Table 11.6.

Table 11.6: Variation of melting and boiling points of transition metals, lanthanides and actinides in K. It is evident that melting and boiling points of 3d metals are relativeIy lower than those Chemistry of d- and of 4d and 5d metals. The lanthanide and actinide metals are rather softer. However, no .f-Block Elements definite trend is observed in their melting and boiling points.

11.5.3 Enthalpies of Ionization The enthalpies of ionization of the d block elements are higher than s block elements but lower thanp block elements. In the transition series of elements, the enthalpy of ionization gradually increases across a period. This is due to an increase in the effective nuclear charge experien~edby ns elecion, causing a contraction in atomic size and thus making is difficult to ionize. But with increasing number of (n-l)d electrons the effect of increase in nuclear charge is partly balanced by increase in screening effect and consequently the increase is rather small. The first enthalpies of ionization of all three series transition elements are given in Table 11

Table 11.7: The first enthalpies of ionization of 3d, 4d and 5d transition elements in kJ mol-'.

- 3r It is important to note that the ionization enthalpies of 5d elements are comparatively higher than 4d and 3d elements. This is because of greater effective nuclear charge felt by outer ns electrons due to imperfect shielding by the intervening 4f electrons.

The second and third enthalpies of ionization follow the same trend except for second enthalpy of ionization of chromium and copper which are relatively high due to additions1 stability associated with half- filled and completely filled configuration of chromium and copper, respectively. A comparison of the periodic trends of ionization enthalpy reveals that for s andp block elements (Table 10.1) its variation is much greater than that for d or f block elements (Table 11.7).

Within a group the ionization enthalpy decreases from 3d to 4d transition series and then increases from 4d to Sd elements. This increase is because of smaller atomic size of 5d elements due to lanthanide contraction.

The ionization enthalpies of lanthanides are quite low. The first and second ionization enthalpies are comparable to those of alkaline earth metals.

The ionization enthalpies of actinides are not accurately known. It is, however, expected that for early actinides it must be lower than for the early lanthanides. This is because the Sf orbitals in early actinides will penetrate less into the inner core and thus the 5f electrons will be more effectively shielded from the nuclear charge than are the 4f electrons of the corresponding lanthanides. The ionization enthalpies are correlated with relative stabilities of various oxidation states of some ions. For example, the sum of first and second ionization enthalpies of Ni(lT) is less than that for Pt(lI), thus the compounds of Ni(lI) are thermodynamically more stable than Pt(II) compounds. Chemistry of Elements The standard electrode potentials for the half reaction ~~'(a~)+ 2e- - M(s) for 3d-transition metal ions, and for half reaction ~n~'(a~)+ 3e- - Ln(s) for lanthanides are given in Table 11.8

Table 11.8: Stapdard electrode potentials of 3d-transition elements (M~+/M)and of lanthanides (Ln3'/Ln) in volts. I Element I Electrode I Element 1 Electrode 1 Element 1 Electrode 1 Potential -1.2 -0.91 -1.18 -0.44 -0.28 -0.25 +0.34 -0.76

,The standard electrode potential is a me.asure of electropositive character and the ,reactivity of metals.

The standard electrode potential of a transition element is generally lower than that of standard hydrogen electrode with the exception of copper. Hence these elements are good reducing agents but not as good as s- andp-block elements. Copper that has a positive electrode potential is a poor reductant.

The electrode potentials of lanthanides and actinides are quite low. Therefore, they are highly electropositive and reactive metals. While in the transition metals there is no regular trend of variation of standard electrode potentials due to accumulative effects af ionization enthalpies, enthalpies of sublimation and enthalpies of hydration, the electrode potentials in lanthanides increase from cerium to lutetium consistent with . small decrease in ionic radli due to lanthanide contraction.

SAQ 5 a) The enthalpies of ionization of 3d elements increase slowly along a period. Explain.

b) Why is the first enthalpy of ionization of 5d elements greater than those of 3d

and 4d transition elements? .d

......

11.5.4 Oxidation States The transition elements exhibit a large number of oxidation states differing usually by units of one instead of the double units shown byp-block elements. The variable oxidation states of transition elements is due to the involvement of (n-l)d electrons, in addition to the ns electrons, in bonding, as they have comparable energy. The various oxidation states exhibited by transition elements are listed in Table 1 1.9. ------Table 11.9: Various oxidation states of the transition elements. Chemistry of d- and .f-Block Elements

(The most stable states are in bold type; less common and unstable states are enclosed in paranthesis)

It can be seen from the table that, in general, there are lesser number of oxidation state in the beginning and at each end of the series and a maximum number of oxidation states comes in the middle. The lowest oxidation state is equal to the number of ns electrons while the highest oxidation state is equal to the total of ns and (n-l)d electrons. The lesser number of oxidation state in the beginning of the series may be due to the . presence of too few electrons to share or lose, and at the end of the series it can be due to hepresence of too many electrons and thus fewer empty orbitals to share electrons. Also there is a reduced tendency of higher oxidation state towards the end of the series. This could be due to steady increase in the effective nuclear charge and thus increasing the binding energy of (n-l)d electrons relative to ns electrons along the transition series form left to right. The heavier elements of any d group are more stable in their higher oxidation states, contrary to the trend in the main group elements..Thus the Mn(I1) is most stable state while is stable in IV and VII states. Similarly, chromium(V1) . as in K2Ca4 is a strong oxidizing agent and is easily reduced to cIZ' but Mo(V1) and W(V1) as in K2Mo04and K2WO4,respectively are not easily reduced and ~0~'and w3+ are not easily formed. In the transition elements, the stability of higher oxidation states increases in the order 3d << 4d < 5d. Thus, although the most stable state for Ti, Zr and Hf is IV, Ti(II) and Ti@) can be formed from Ti(IV) with good reducing agent but lower oxidation state of Zr and Hf are extremely difficult to obtain.

The lowest common oxidation state is 2+, which arises due to the loss of ns electrons. For 4d and 5d series including Ti in the 3d series 2+ states is not very stable except for Pd and Pt. However, from manganese to copper 2+ state is quite stable. The highest oxidation state (VIII) is shown by Ru and 0swhich happens to be the highest oxidation Chemistry of Elements state in the periodic table along with that of Xe which also shows highest oxidation states are usually stabilized in the compounds VI VI v For example, MnO; ,c~o~-,vo; etc. This is due to the fact that 0 and F are strong oxidizing agents in these compounds. The relative stabilities of various oxidation states, for example ~i~'> ~i~',Fe3' > Fe2', ~n~'> Mn3' is associated with the extra stability of half filled configurations.

The sums of first three ionization enthalpies of lanthanides are comparatively lower, so the elements are highly electropositive. They readily form M3' ion. The known oxidation states of lanthanides and actinides are given in Table 11.10.

Table 11.10: Different oxidation of lanthanide and actinide elements.

(Bold number indicates the most stable state in aqueous solution)

For the first inner transition series (4j) elements, actinium and trans-americium (Am to Lr) elements the 3+ oxidation state is not only most common but it is the most stable state. It is the stabilizing effects on the 4f orbitals by increasing nuclear charge, which is . responsible for the prevalence of 3+ oxidation states in lanthanides. As successive electrons are removed from a neutral lanthanide atom the 4f orbitals, which penetrates through the inner core of electrons towards the nucleus more than 5d and 6s orbitals, becames more and more stabilized and by the time an ionic charge of 3+ is reached it becomes so stabilized that it is difficult to remove any further electron. It is believed that in farming tripositive lanthanide or actinide ions, the ns2 electron alongwith (n-l)di electrons are removed. In the elements in which no (n-l)d electron is availiue one of the electrons present in (n-2)forbital is lost.

Besides the 3+ oxidation state, some of the lanthanides and actinides exhibit other oxidation states also. In these cases, there is some evidence that the io-nswithf (La", ce4+,AC~', Th4', pa5', u6+),f7(EU~+, ~d~', Tb4+, cm3', ~k~')and f (yb2', Lu33 configurations show greater stability. However, pr4' (4f), ~d~+(4f), sm2' (4f), ~m~' (4f3) with non-f, nonf and non-f4 configurations are also known. This indicates that 3 there may be some other factors also such as enthalpies of ionization, enthalpies of sublimation of metals and lattice energies etc., which are responsible for the stability of these oxidation states. A few lanthanides, which exhibit oxidation state of +2 and +4, have a tendency to attain 3+ states. Thus, sm2+is a good reducing agent while ce4' is a good oxidizing agent.

It can be seen from the Table 11.10 that at least the first six actinide elements show a greater variation in the oxidation states. The 3+ oxidation state is not the most stable state until americium. The most stable oxidation state is the one, which involves all the valence electrons in bonding. This may be because the Sf electrons are held less firmly Chemistry of d- and by the nucleus due to more effective shielding from the nuclear charge than the 4f ,f-Block Elements

electrons of the corresponding lanthanides. Hence they are available for bonding in ,~ early actinides.

11.5.5 Colour of the Complexes We know that most of the compounds of transition elements are coloured in solid state or in solutions. This is unlike the compounds of main group elements, which are usually colourless unless the anion is coloured. A compound appears coloured because it absorbs certain coloured portion of the incident white and exhibits the cdmplimentary colour. For example, [Ti(H20)613+ is reddish violet because it absorbs the green portion of the white light. In the preceding section, it has been pointed out that C the transition elements or their free ions have incomplete (n-l)d orbitals, all of which are equal in energy or in a degenerate state. However, in complexes the degeneracy of the d-orbitals is lifted as these split into sets of orbitals having different energies. The F colours of the transition metal compounds are usually due to the electronic transition from a d-orbital of lower-energy to that of higher energy. Such electronic transition is called d-d transition. The energy required for this transition is fairly small and is available in the visible region. For example, [T~(H~o)~]'+is reddish violet because of single d electron transition from a d-orbital of lower energy (tz,orbitals) to that of higher energy (e, orbital). Whenever, the d-orbitals are empty or filled, the d-d transitions are not possible. In such cases, the ions will not show any colour. For example, the compounds of sc3+,~i~+, CU' and zn2+are white or colourless. The observed colours of some hydrated transition metal ions in different oxidation state are given in Table 11.1 1.

Table 11.11: Observed colours of hydrated transition metal ions in different oxidation states. Element + 2 +3 + 6 + 7 /' Colourless

Violet

Green

VioletJGreen OrangeNellow c~,o:- C~O- Red Green Purple Mn0;- MnO; Yellow/Brown

Blue

Blue

The absorption of light The observed deep colour of some transition metal compounds such as KMn04 (purple) sometimes causes the and K2O207(orange) having empty d orbitals is not due to d-d transition but due to a transfer of an electron differeqt type of electronic transition called charge transfer transition from a between molecular orbitals molecular orbital centred mainly on the ligand to the one centred mainly on the metal that are centred on different atoms. Such electronic atom (ligand-to-metal charge transfer). It is also possible that electron is transferred transitions are called charge- from a molecular orbital on metal to that on the ligand (metal-to-ligand charge transfer), transfer transitions. '4 9 Chemistry of Elements for example, in Cr(C0)6 end [N~(cN)~]~-.Generally, d-d transitions give pale co1,ours while charge transfer transitions give dark colours.

Like transition elements, most of the lanthanide and actinide ions are coloured in the solid state or in aqueous solutions. The observed colours of lanthanide ions and actinide ions in 3+ and 4+ oxidation states are given in Table 1 1.12.

Table 11.12: Observed colours of lanthanide ions and actinide ions.

La3'(4f) ce3'(4f) ~?+(4f) ~d~+(4f) pm3+(4f) sm3+(4P) Eu3+(4P) ~d~+(4f7)

Colourless Colourless Green Lilac , Pink Yellow Pale pink Colourless

T b3'(4f8) ~~~+(4f)~o~'(4f') E?+(4f1) ~m~'(4f') yb3+(4f3) Ld+(4f4)

Pale pink Yellow Pale Lilac Pale green Colourless Colourless yellow

Ac3+(5f) u3+(5f) Np3+(5f) pu3'(5.) Am3'(5f) cm3+(5f7) ~k~+(5f8) ~p(5f)

Colourless Red Blue Violet Pink Colourless Yellow Pale green green

~h~'(5f) pa4'(5f) u4+(5f) Np4+(5f) pu4+(5f) Am4'(5f) cm4'(5f)

Colourless Colourless Green Yellow Brown Rose pink Pale green yellow

It is seen that the colour sequence of lanthanide ions from La3+- Gd3+is accidentally repeated from Gd3+- Yb3+,and the colours of the lanthanide ions with 4f and 4f4-" configuration are almost same. For example, Nd3+and Er3+ions having 44 and 4f' configurations respectively are lilac in colour. However, several lanthanide ions in other oxidation states and their isoelectronic 3+ counterparts do not have similar colours. For example, Lanthanide ion Colour Isoelectronic Colour lanthanide ion La3+(4f') Colourless ce4' Orange red Eu3'(4f) Pale pink sm2' Blood red Gd3'(4J') Colourless Eu2' Pale greenish yellow Lu3+(4f4) Colourless Yb2' Yellow

As in case of transition metal ions, the lanthanide and actinide ions withf (La3+,~h~', pa5'),J' (Gd3+,cm3') and f (Lu3+)configurations are colourless. Other actinide ions having 5f to 5f configurations are coloured. For example, u3+(54) is red; Np3' (5f) is blue etc.

The colowrs of lanthanide ions and actinide ions are due to thef;f transition, similar to d-d transition in transition elements, and are virtually independent of the nature of anion present. A charge transfer phenomenon, as in transition elements, is also observed in certain lanthanide and actinide ions. For example, Eu3+,ce3+, cm3+ ions etc. The orange red colour of ce4+(4f configuration) is due to charge transfer transition from ligand to 4f orbitals of cerium. SAQ 6 Chemistry of d- and .f-Block Elements a) Which among the following ions are coloured and why?

b) Name any three lanthanide elements that show unusual oxidation states.

" 11.5.6 Magnetic Properties We are now aware that d-block andf-block elements have incomplete d or f sub shells k and they usually have unpaired electrons in their d or f orbitals in many of their compounds. An unpaired electron by virtue of its spin and its orbital motion produces a definite magnetic moment that exists in absence of any external magnetic field. When a sample of the substance is placed in a magnetic field, the applied field causes the magnetic moments to align themselves in the direction of the field. It is then attracted towards the magnetic field with a force proportional to the field strength times the field gradient. Such a substance is calledparamagnetic. However, a substance that has no unpaired electrons is repelled away by the magnetic field, and is called diamagnetic. Thus, most of d-block andf-block elements or their compounds are paramagnetic.

In some transition elements and their compounds such as , cobalt, and Cr02 and several lanthanides a spontaneous alignment of the moments in the same direction occurs due to coupling (or interaction) between individual magnetic moments of the unpaired electrons below a critical temperature called Curie temperature. This gives rise to a high permanent magnetic moment. These substances are calledferromagnetic. However, in certain compounds of ~e~',~n~' and ~d~' ions and MnO half of the moments align opposite to the other half, thus giving a zero net magnetic moment. These are called antiferromagnetic substance and the critical temperature is known as Nee1 temperature. In another type of substance calledferrimagnetic, e.g. magnetite Fe304some of the moments align opposite to the other but with unequal number or magnitude, which gives a finite net magnetic moment to the compound. It is important to mention that above the critical temperature, which is a characteristic of the substance, the ferromagnetic or antiferromagnetic substances become paramagnetic. The magnetic moment (p) is usually expressed in unit called Bohr Magneton (B.M). The general equation for the magnetic moment is given by

where S is the sum of spin quantum numbers and L is the sum of orbital angular momentum quantum numbers of all the electrons. However, in many compounds of the first transition series elements, the orbital contribution to the magnetic moment is quenched out by the electrical field of the surrounding atoms and as an approximation the observed magnetic moment is considered to arise only hmunpaired spins. Putting L= 0 in the above expression, a very simple expression is given for spin only magnetic moment b.Thus,

Since the total of spin quantum numbers is equal to half the number of unpaired electrons, n in any species, the above expression can be written as Chemistry of Elements

This equation is used to estimate the number of unpaired electrons.

The magnetic moment cannot be measured directly hence it is calculated from the measured magnetic susceptibility per mole or simply molar susceptibility (xM).When a substance is placed in a magnetic field of magnitude H, a magnetization I is induced. The ratio IIH is the magnetic susceptibility per unit volume or volume susceptibility K. The volume susceptibility relates to the molar susceptibility and is given by the following equation.

where d is the density and M is the molar mass of the substance. This molar susceptibility is corrected for the diamagnetism and also for some temperature independent paramagnetism, which gives a corrected molar susceptibility X? . The magnetic moment, p is calculated by using the following expression

In the above expression, NA is the Avogadro's constant, k is the Boltzmann constant and T is the absolute temperature. On rearranging and evaluating the constants, tEe magnetic moment can be expressed as

If uncorrected value of XM is used, the corresponding value of p is usually called the effective magnetic moment, h.

The magnetic susceptibility is commonly measured by weighing the sample in the presence and absence of magnetic field using a magnetic balance called Gouy balance. The paramagnetic substance shows an increase in mass in presence of magnetic field whereas diamagnetic materials show a decrease in mass. The susceptibility is calculated from the change in mass of the sample.

the basis of the values of magnetic susceptibility and its temperature dependence, it v'i possible to distinguish between different types of magnetic behaviour shown by a substance. The sign, magnitude and temperature dependence of magnetic susceptibility for a paramagnetic, diamagnetic, ferromagnetic and antiferromagnetic substance is given in Table 11.13.

Table 11.13: Magnetic susceptibility and its temperature dependence. Sign Magnitude Dependence on Magnetic Temperature Behaviour Positive 0 - Decreases Paramagnetism Negative 1 - 500 x None Diamagnetism Positive - lo6 Decreases Ferromagentism Positive 0 - Increases Antiferromagnetism I t It can baseen from Table 11.13 that for a paramagnetic substance magnetic Chemistry of d- and susceptibility is large and positive whereas for a diamagnetic substance it is small and .f-Block Elements slightly negative. In a ferromagnetic substance the susceptibility is highly positive and decreases with rise in temperature but in an antiferromagnetic material it is somewhat lower than that in a paramagnetic substance and increases with increase in temperature.

For most paramagnetic substances the magnetic susceptibility varies inversely with absolute temperature. Thus

where C is a constant called Curie constant characteristic for the particular substance. The above equation is called Curie law. A plot of 11 X: as a function of Tis a straight line (Fig. 11.3), which intersects the origin. In some paramagnetic substances this line does not intersect the origin as'shown in Fig. 1 1.3, because of the presence of some interionic and intermolecular interactions.

Fig. 11.3: Temperature dependence of magnetic susceptibility of a paramagnetic substance.

For such substances Curie-Weiss law is used which is

f'

where 8 is the temperature at which the line cuts the Taxis, hnd is known as Weiss constant and the magnetic moment is given by the following equation.

For ferromagnetic and antiferromagnetic substances, the magnetic susceptibility does not follow the Curie-Weiss law below the Curie temperature and Nee1 temperature, respectively. The temperature dependence of the susceptibility for a ferromagnetic and an antiferromagnetic substance is shown in Fig. 11.4. Chemistry of Elements

...... Ferromagnetic

Fig. 11.4: Temperature dependence of susceptibility of a ferromagnetic and antiferromagnetic substance.

From Fig. 1 1.4 it can be seen that a ferromagnetic substance has very large susceptibility values at low temperature. As the temperature increases its value decreases rapidly till the Curie temperature Tc is reached. Above the Tc, it follows the curie law or curie-weiss law and behaves like a paramu;,etic snbstance. For antiferromagnetic compounds the susceptibility actually inct :es with increasing temperature upto the Neel temperature, TN.Above TNthe co dund reverts back to paramagnetic behaviour. This peculiar behaviour of ferromagnetic and antiferromagnetic substance below the Curie temperature or Neel temperature is due to some interionic interactions of comparable magnitude to that of thermal energy at the Curie or Neel temperature. Above these temperatures thermal energy becomes greater and randomizes the perfectly ordered parallel or antiparallel alignment of the moments. Thus, the susceptibility for ferromagnetic and antiferromagnetic substances decreases with increasing temperature like that of paramagnetic substances.

I9 Table 1 1.14 the experimentally observed magnetic moment values along with those calculated by the spin only formula for some transition metal hydrated ions is given.

Table 11.14: Observed and calculated magnetic moments of some hydrated transition metal ions.

Ion Electronic Unpaired Magnetic Moment (BM) configuration electrons Calculated Experimental

[~i(H20)61~+ 3d' ? 1 1.73 1.75 [V(H~O)~I'' 3d2 ? ? 2 2.84 2.75 --3 3.87 [C~(H~O),]~' 3d4? 'r 'T 'r 4 4.90 [M~(H~~),I~' 3dSf'r ??? 5 5.92 4 4.90 ----[co(H~o)~~~+ 3d7 T.1 T.1 ? 'r ? 3 3.87 4.40 [~i(~20)~1~' 3d8 T.1 T.1 t.1 ? ? 2 2.84 2.90 [CU(H~O)~I~' 38 TL fL T1 TL I+ 1 1.73 1.80 It follows from Table 1 1.14 that for the ions of the first half of the 3d transition Chemistry of d- and elements series the magnetic moments increases with increase in the number of .f-BlockElements unpaircd electrons until the dSconfiguration is achieved. The magnetic moments then decrease because of decrease in number of unpaired electrons with pairing of d electrons.

All lanthanide and actinide ions, except those withf configuration (La3', Ce4', AC)', ~h~+,paS+,u6") and f configuration (yb3', Lu3'. Lr3') contain unpaired electrons and are therefore paramagnetic. The effective magnetic moments of lanthanides and - The quantum number J is actinide ions is given by: the vector sum of L and S quantum numbers and can have only positive values fiom [L+SI...... I L-s I or h whiic J is the resultant inner quantum number which is obtained by L-S coupling and be zero. g is Lande splitting factor or gyromagnetic ratio. The value of g is given by

The magnetic mment of the lanthanide ions with? (La3+),f (~d~')and f (Lu3') configuration, where there is no orbital contribution, agrees well with spin only value as observed in many 3d transition metal ions. In all other lanthanide ions the observed magnetic moments are higher than those calculated from spin only formula. When the nlonlents arc calculated taking both orbital and spin contributions into account the experimental and calculated values are in excellent agreement except in case of sm3' and EU~'.The experimental and calculated effective magnetic moments of the trivalent lanthanide ions is given in Table 11.15.

Table 11.15: Observed and calculated effective magnetic moment of trivalent lanthanide ions.

. Lanthanide Ion I Magnetic Moment Observed 1 Calculated

1 The quenching of orbital contributions as observed in first row transition elements is not possible in thef-block elements because of the fact that (n-2'forbitals are quite I deeply seated in the atom or ion and are well shielded by the outerlying 5s and 5p 1 electrons from the effects of electric field of the ligands. The magnetic-properties of I actinide ions are more complicated than those of lanthanide ions and are usually I lower than the calculated values. This in part arises form (a) the fact that 5f electrons Chemistry of Elements are nearer to the surface of the atom and are easily influenced by the chemical environment, (b) the less sharply defined distinctions between Sf and 6d electrons as compared Fo 4f and 5d electrons, and (c) the lesser applicability of the expression p, =g,/Gas compared to that in the lanthanide elements. pu3*and ~m"ions, like sm3' and EU~' ions, show anomalous magnetic moments.

SAQ 7

Find out the number of unpaired electrons in the ions 2r4', ~d~',D~~*, MO~', ~b~+, pa4', N~~'

11.5.7 Catalytic Properties Many transition metals and their con~poundshave good catalytic properties. The catalytic property of these metals is due to the fact that they can utilize both (n-l)d and ns electrons for the formations of bonds between reacting species and the surface atoms of the metals. This increases the concentration of the reactants at the surface of the metals in conformations favourable for reaction and weakens the bonds in the reactant molecules with the result that activation energy is lowered. The catalytic activity of transition metal and their compounds is attributed to the ability of these metals to form complexes with ease and to undergo oxidation-reduction reaction due to their variable oxidation states and thus act as a source of, or sink for electrons. There are many industrially important reactions, which are carried out with the help of transition metals or their compounds. Some common exanlples are: TiCl4 and trialkylaluminiun~(as Zeigler Natta catalyst) for polymerization of ethane. (i i) PtIRh in oxidation of ammonia to nitric oxide in manufacture of nitric acid. (iii) Fe/Mo in manufacture of ammonia by Haber process. (iv) V205in oxidation of SO2to SO3 in the manufacture of sulphuric acid by Contact process. (v) Ni in hydrogenation of oils. (vi) Ni (Raney nickel) in the reduction processes. (vii) Pd in many hydrogenation reactions. (e.g. phenol to cyclohexanone) (viii) FeS04and H202(Fenton's reagent) for oxidation of alcohols to aldehydes. (ix) RhC1(PPh3) (Wilkinson catalyst) for homogeneous reduction of alkenes and alkynes. (4 H2PtC16(Speier's catalyst) In sillcone technology. (xi) Mn02in decomposition of KC1O3 to give 02. (xii) Cu-ZnO in reduction of CO to methanol. (xiii) Cu in manufacture of (CH3)2SiC12used to make silicones.

Many transition elements are important catalysts in several biological processes. A number of metals including some transition metals present in very small quantities in plants and animals are essential for the enzymes to function. For exanlple, the enzyme 'nitrogenase' containing an iron protein and molybdenunl iron protein catalyses the reduction of molecular to ammonia. A zinc enzyme 'carbonic anhydrase' catalyses the hydration of C02to give carbonic acid (COz + OH -+ HCO; at pH 7) and a variety of superoxide dimutases which catalyse the dimutation (disproportion) of superoxide ion ( 0;) into hydrogen peroxide and O2(2 0; + 2~'-+ H202+0, contain Cu-Zn, Mn and Fe atoms. Vitamin B12coenzyme (formed by reaction of ATP ------with vitamin BI2)required in many biological processes contains a cobalt atom at its Chemistry of d- and centre. Iron atonms are involved in haemoglobin of blood and in ferrodoxins. ,f-Block Elements

11.5.8 Formation of Complexes The transition metals have a characteristic property to form complex ions with a large number of neutral molecules and ions. For example, (F~(cN),]~-,[Ni(NH3)412+, [P~(NH~)~]~+,[CO(NH,)~C~]~+ etc. This unique tendency to foml complex ions is because the d-block transition nmetal atoms form small, highly positively charged ions which have vacant d-orbitals of suitable energy to accept lone pairs of electrons donated by other groups or ligands. In case of transition metals in low oxidation states, the d electrons can also become involved in n bonding with ligands that have vacant n orbitals.

In complex ions, the traimsition metals are surrounded by a number of groups called ligands. The type of ligands and the number of ligands attached to central metal atom or ion and the arrangement of the ligands is also varied. In majority of conmplexes the transition metal ion is surrounded by six ligands arranged octahedrally or less frequently at the corners of a trigonal prism. In some complexes, the transition nmetal ion is bonded by four ligands arranged tetrahedrally or less frequently at the vertices of a square. There are few complexes in which five or seven ligands are attached to central nmetal ions giving trigonal bipyramidal or pentagonal pyramidal geometries. Complexes with coordination number greater than six are usually found in second and third transition series elements and also for lanthanide and actinide elements. The bonding between the ligand and the central metal ion can either be predominantly electrovalent or covalent or in many cases intermediate between the two extreme.

The lanthanide and actinide ions have a strong tendency to form complexes with a imunmber of and nitrogen donor ligands. They preferably form most stable coinplexes with oxygen containing chelating li@nds such as oxalic acid, nitric acid, ethylenediaminetetraacetic acid (EDTA) and P-diketones than with nitrogen donor ligands. The complexes of these metal ions with citric acid, tartaric acid and EDTA are watcr-solublc which facilitates separation of the metal ions by ion exchange chromatography. Probably, because of their comparatively higher charge to size ratio, thc actinide ions have a much higher tendency to form complexes than the lanthanide~.Also due to existence of more number of oxidation states than lanthanides the complexation behaviour of actinides is more varied. Because of larger size of lanthanide ions and also because 4f orbitals are embedded in the inert core the metal - ligand bonding is predonlinantly ionic. However, in actinides the Sf and 6d orbitals are of conmparable energy than the 4f and 5d orbitals and Sf orbitals also spatially projects outwards than 4f orbitals, the Sf electrons become involved in bonding and make a covalent contribution in the metal-ligand bonding. The lanthanide and actinide ions generally do not form complexes with 7c bonding ligands because of absence of 5d electrons and non-participation of 4f electrons in bonding.

In transition metal con~plexesthe maximum coordination number is usually restricted to nine but high coordination nunmbers are more common in lanthanide and actinide complexes. For example, the coordination number of metal ions in [Th(acac)& [~d(~20)6]-",[Ce(N03)4(0PPh,)2] and [c~(N)~)~]~-is 8, 9, 10 and 12, respectively. This is because the transition elements have only nine valence shell orbitals (five d, one s and threep) whereaa in lanthanide and actinides f orbitals also exist along with thesc nine valence orbitals. Thus there are a few complexes of these ions having coordination number six but those with coorGnatian number from seven to nine are nmost common. However, the lanthanide ions can be induced to form complexes with lower coordination number by use of bulky ligands. For higher coordination number several alternative geometries are possible and the energy difference between them is commonly very small. Chemistry of Elements 11.5.9 Formation of Interstitial Compounds (Interstitial Solid Solutions) and Alloys (SubstiPutional Solid Solutions) We know that nearly all the transition metals have close packed (cubic or hexagonal) crystal structures while some of them have body-centred cubic (bcc) lattices. In both of the close packed lattices the metals occupy 74% of the total space while in body- centred cubic structures 67% of the available space is occupied by metal atoms. The remaining space is unoccupied and is regarded as a hole or vacant site. In close packed crystal structure there is one octahedral hole and two tetrahedral holes per metal atam. These octahedral holes can be occupied by other atoms of atomic radius lesser than 0.41 times that of metal atom. Similarly, tetrahedral hole may be occupied by atoms of radius lesser than 0.23 times that of the metal atom.

The transition metals including probably some lanthanide and actinide elements can trap some small non-metal atoms such as hydrogen, , and nitrogen etc. in the interstitial spaces or holes in their crystal lattices, thus forming interstitial compounds or interstitial solid solutions. Because of their comparatively large size carbon, boron and nitrogen always occupy octahedral interstitial spaces whereas smaller hydrogen atoms always occupy tetrahedral holes. The crystal structure of the metal ofteh changes during the formation of interstitial compounds. The composition of these compounds is generally non-stoichiometric e.g., TiHl 73, P~Ho.~~,YbHz 5s, ThH375, etc., but in many carbides and nitrides it may approach a 1 :1 stoichiometry and a regular structure when all octahedral holes are occupied e.g. Tic, CrN, UC, ThN etc. The later transition metals of 3d series (Cr, Mn, Fe, Co, Ni) form non- stoichiometric carbides with irregular structures with incomplete occupancy of octahedral holes e.g. Mn7C3,Cr7C3, Co3C , Ni3C etc, which are more reactive than the interstitial carbides of early transition metals. The interstitial compounds retain their most of mdtallic properties such as metallic luster and conductivity but are much more harder and have higher melting points but are more brittle than the parent metals.

The presence of these atoms in their interstitial spaces results in decrease in malleability and ductility of the metals, but increases their tensile strength. Thus carbon are harder due to carbon atoms, which prevent the iron atoms from sliding eve$ one another.

Most of theltransition metals have identical crystal structures and almost similar atomic radius. Therefore, the metal atoms of one element can easily substitute the position f other in the crystal lattice of another transition metals forming a 4 0 substitutional solid solutions or alloys. For example, nickel (ccp, r =1.25 A ) and zinc 0 (ccp, r = 1.28 A ) form alloys with each other. On the other hand, (bcc, r = 0 0 1.86 A ) and potassium (bcc but r = 2.27 A ) do not form alloy. The alloys so formed are relatively hard and have higher melting points. These alloys are of much importance, e.g. brass and bronze are alloys of copper-nickel and copper- respectively. Monel metal is an alloy of nickel-copper. The transition metals form alloys of variable stoichiometry, which are stable over limited range of compositions. An alloy shows a variety of phases. For example, a binary alloy of copper-zinc (brass) has three important phases, (3-phase, y-phase and &-phase.These three phases are called Hume~Rotherycompounds. These phases have different compositions and structure. For example, P-brass phase has bcc structure and CuZn composition, y- brass phase has cubic structure and composition CusZns and &-brassphase has hcp structure and composition CuZn3.Even the phase with same crystal structure has different composition. There are some other phases also besides the Hume-Rothery phases. Chemistry of d- and .f-Block Elements In this unit, w have studied what are d andf block elements, the IUPAC naming of the elements o\ fourth transition series, the electronic configuration of the d block elements as well as f block elements and how the flling of orbitals takes place with the increase in atomic number, the periodicity in their properties, the variation of size, melting and boiling points, enthalpy of ionization , electronegativity, electrode potential, oxidation states. Besides these, a few of their properties like colour, magnetic properties, complex formation, catalytic properties and formation of interstitial compounds have also been discussed. These can be summarized as following: In d block elements the differentiating electin enters the penultimate (n-l)d atomic orbitals (where n = 4, 5,6 or 7). The f block elements- lanthanides and actinides are characterized by the filling of ante penultimate (n-22f orbitals (n = 6 or 7). In the building up of the electronic configurations of the atoms of transition elements, lanthanides and actinides, the filling of atomic orbitals take place in the order 4s, 3d, 4p, 5s, 4d, 5p, 6s, 5d'-*, 4J; 5d3-lo,6p, 7s, 6d - 5J Configurations with half filled and completely filled d or f sub shells are more stable mainly due to exchange energy. The atomic and ionic radii of transition elements decreases slowly with increase in atomic number. The decrease in radii of lanthanide and actinide ions is called lanthanide and actinide contraction, respectively. The enthalpies of ionization of transition elements are in between s- andp- block elements. The transition elements including lanthanide and actinide elements show variable oxidation states due to the participation of variable number of d or f electrons in bonding. For first row transition series elements the lower oxidation states are more stable but higher oxidation states are more stable for second and third series including early actinide elements. For the lanthanides, actinium and trans americium elements, 3+ oxidation states are more stable. All the transition elements, lanthanide and actinide ions which have unpaired d . or f electrons are paramagnetic. The paramagnetism of these ions depends on both spin and orbital angular momentum of the unpaired electron. However, for 3d transition metal ions the paramagnetism depends only on spin momentum. Most of their compounds are coloured in solution or in solid state due to incomplete d or f sub shell so that d-d orfif electronic transitions are possible. All transition elements, lanthanides and actinides have remarkable ability to form complexes.

1. The of lanthanum and the lanthanide elements gradually decrease from lanthanum to lutetium except for europium and ytterbium, which have abnormally low densities. Explain.

2. In the separation of nuclear fission products by elution from a column of ion- exchange resin, the following sequence,of ions is observed: yb3+,~m~+, E~~', HO~+,y3+, D?, ~b~+,Gd? etc. Explain the appearance of y3+.inthis series.

3. Explain the term lanthanide contraction. What are its consequences?

4. Why most off block compounds are paramagnetic in nature? How do the following properties vary in the transition elements (a) Enthalpy of ionization (b) Melting and boiling points (c) Electronegativity

Explain why zinc and cadmium are soft metals.

[T~(H~o)~]~+is coloured while [sc(H~o)~]~+is colourless.

Explain the following: (a) ~n~+ion has highest magnetic moment amongst the bivalent ions of 3d transition elements. (b) The species [cuc1412-is known but [CUI~]*-is not.

Arrange the following ions in order of: (a) increasing size: La3', AC~',y3+, co3+, Lu3*, I?, Fe3+ (b) decreasing paramagnetic character: cr3+,Fe3+, CO~~, ~i~~