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Unit Objectives The dd-- and f- f- 4 After studyi...

Unit Objectives The dd-- and f- f- 4 After studying this Unit, you will be able to Block Elements learn the positions of the d– and f-block elements in the periodic table; Iron, copper, silver and gold are among the transition elements that know the electronic configurations have played important roles in the development of human civilisation. of the transition (d-block) and the The inner transition elements such as Th, Pa and U are proving inner transition (f-block) elements; excellent sources of nuclear energy in modern times. appreciate the relative stability of various oxidation states in terms of electrode potential values; The d-block of the periodic table contains the elements describe the preparation, of the groups 3-12 in which the d orbitals are properties, structures and uses progressively filled in each of the four long periods. of some important compounds The f-block consists of elements in which 4 f and 5 f such as K2Cr2O7 and KMnO4; orbitals are progressively filled. They are placed in a understand the general separate panel at the bottom of the periodic table. The characteristics of the d– and names transition metals and inner transition metals f–block elements and the general are often used to refer to the elements of d-and horizontal and group trends in f-blocks respectively. them; There are mainly four series of the transition metals, describe the properties of the 3d series (Sc to Zn), 4d series (Y to Cd), 5d series (La f-block elements and give a and Hf to Hg) and 6d series which has Ac and elements comparative account of the from Rf to Cn. The two series of the inner transition lanthanoids and actinoids with metals; 4f (Ce to Lu) and 5f (Th to Lr) are known as respect to their electronic lanthanoids and actinoids respectively. configurations, oxidation states and chemical behaviour. Originally the name transition metals was derived from the fact that their chemical properties were transitional between those of s and p-block elements. Now according to IUPAC, transition metals are defined as metals which have incomplete d subshell either in neutral atom or in their ions. Zinc, cadmium and 10 mercury of group 12 have full d configuration in their ground state as well as in their common oxidation states and hence, are not regarded as transition metals. However, being the end members of the 3d, 4d and 5d transition series, respectively, their chemistry is studied along with the chemistry of the transition metals. The presence of partly filled d or f orbitals in their atoms makes transition elements different from that of Reprint 2024-25  the non-transition elements. Hence, transition elements and their compounds are studied separately. However, the usual theory of valence as applicable to the non- transition elements can be applied successfully to the transition elements also. Various precious metals such as silver, gold and platinum and industrially important metals like iron, copper and titanium belong to the transition metals series. In this Unit, we shall first deal with the electronic configuration, occurrence and general characteristics of transition elements with special emphasis on the trends in the properties of the first row (3d) transition metals along with the preparation and properties of some important compounds. This will be followed by consideration of certain general aspects such as electronic configurations, oxidation states and chemical reactivity of the inner transition metals. THE TRANSITION ELEMENTS (d-BLOCK) 4.1 Position in the The d–block occupies the large middle section of the periodic table Periodic Table flanked between s– and p– blocks in the periodic table. The d–orbitals of the penultimate energy level of atoms receive electrons giving rise to four rows of the transition metals, i.e., 3d, 4d, 5d and 6d. All these series of transition elements are shown in Table 4.1. In general the electronic configuration of outer orbitals of these elements 4.2 Electronic 1– 10 1–2 is (n-1)d ns except for Pd where its electronic configuration is 4d105s0. Configurations The (n–1) stands for the inner d orbitals which may have one to ten of the d-Block electrons and the outermost ns orbital may have one or two electrons. However, this generalisation has several exceptions because of very Elements little energy difference between (n-1)d and ns orbitals. Furthermore, half and completely filled sets of orbitals are relatively more stable. A consequence of this factor is reflected in the electronic configurations of Cr and Cu in the 3d series. For example, consider the case of Cr, 5 1 4 2 which has 3d 4s configuration instead of 3d 4s ; the energy gap between the two sets (3d and 4s) of orbitals is small enough to prevent electron entering the 3d orbitals. Similarly in case of Cu, the 10 1 9 2 configuration is 3d 4s and not 3d 4s. The ground state electronic configurations of the outer orbitals of transition elements are given in Table 4.1. Table 4.1: Electronic Configurations of outer orbitals of the Transition Elements (ground state) 1st Series Sc Ti V Cr Mn Fe Co Ni Cu Zn Z 21 22 23 24 25 26 27 28 29 30 4s 2 2 2 1 2 2 2 2 1 2 3d 1 2 3 5 5 6 7 8 10 10 Chemistry 90 Reprint 2024-25  2nd Series Y Zr Nb Mo Tc Ru Rh Pd Ag Cd Z 39 40 41 42 43 44 45 46 47 48 5s 2 2 1 1 1 1 1 0 1 2 4d 1 2 4 5 6 7 8 10 10 10 3rd Series La Hf Ta W Re Os Ir Pt Au Hg Z 57 72 73 74 75 76 77 78 79 80 6s 2 2 2 2 2 2 2 1 1 2 5d 1 2 3 4 5 6 7 9 10 10 4th Series Ac Rf Db Sg Bh Hs Mt Ds Rg Cn Z 89 104 105 106 107 108 109 110 111 112 7s 2 2 2 2 2 2 2 2 1 2 6d 1 2 3 4 5 6 7 8 10 10 The electronic configurations of outer orbitals of Zn, Cd, Hg and Cn 10 2 are represented by the general formula (n-1)d ns. The orbitals in these elements are completely filled in the ground state as well as in their common oxidation states. Therefore, they are not regarded as transition elements. The d orbitals of the transition elements protrude to the periphery of an atom more than the other orbitals (i.e., s and p), hence, they are more influenced by the surroundings as well as affect the atoms or molecules n surrounding them. In some respects, ions of a given d configuration (n = 1 – 9) have similar magnetic and electronic properties. With partly filled d orbitals these elements exhibit certain characteristic properties such as display of a variety of oxidation states, formation of coloured ions and entering into complex formation with a variety of ligands. The transition metals and their compounds also exhibit catalytic property and paramagnetic behaviour. All these characteristics have been discussed in detail later in this Unit. There are greater similarities in the properties of the transition elements of a horizontal row in contrast to the non-transition elements. However, some group similarities also exist. We shall first study the general characteristics and their trends in the horizontal rows (particularly 3d row) and then consider some group similarities. On what ground can you say that scandium (Z = 21) is a transition Example 4.1 element but zinc (Z = 30) is not? On the basis of incompletely filled 3d orbitals in case of scandium atom 1 Solution in its ground state (3d ), it is regarded as a transition element. On the 10 other hand, zinc atom has completely filled d orbitals (3d ) in its ground state as well as in its oxidised state, hence it is not regarded as a transition element. 91 The d- and f- Block Elements Reprint 2024-25  Intext Question 4.1 Silver atom has completely filled d orbitals (4d10) in its ground state. How can you say that it is a transition element? We will discuss the properties of elements of first transition series only in the following sections. 4. 3 General 4.3.1 Physical Properties Properties of Nearly all the transition elements display typical metallic properties the Transition such as high tensile strength, ductility, malleability, high thermal and electrical conductivity and metallic lustre. With the exceptions of Zn, Elements Cd, Hg and Mn, they have one or more typical metallic structures at (d-Block) normal temperatures. Lattice Structures of Transition Metals Sc Ti V Cr Mn Fe Co Ni Cu Zn hcp hcp bcc bcc X bcc ccp ccp ccp X (bcc) (bcc) (bcc, ccp) (hcp) (hcp) (hcp) Y Zr Nb Mo Tc Ru Rh Pd Ag Cd hcp hcp bcc bcc hcp hcp ccp ccp ccp X (bcc) (bcc) (hcp) La Hf Ta W Re Os Ir Pt Au Hg hcp hcp bcc bcc hcp hcp ccp ccp ccp X (ccp,bcc) (bcc) 4 (bcc = body centred cubic; hcp = hexagonal close packed; ccp = cubic close packed; X = a typical metal structure). W The transition metals (with the exception Re Ta of Zn, Cd and Hg) are very hard and have low volatility. Their melting and boiling points are 3 Mo Os high. Fig. 4.1 depicts the melting points of Nb Ru transition metals belonging to 3d, 4d and 5d Ir series. The high melting points of these metals Hf Tc are attributed to the involvement of greater M.p./10 K 3 Cr Rh number of electrons from (n-1)d in addition to Zr V Pt 2 the ns electrons in the interatomic metallic Ti Fe bonding. In any row the melting points of these Co Pd 5 Ni metals rise to a maximum at d except for Mn anomalous values of Mn and Tc and fall Cu regularly as the atomic number increases. Au 1 Ag They have high enthalpies of atomisation which Atomic number are shown in Fig. 4.2. The maxima at about Fig. 4.1: Trends in melting points of the middle of each series indicate that one transition elements unpaired electron per d orbital is particularly Chemistry 92 Reprint 2024-25  favourable for strong interatomic interaction. In general, greater the number of valence electrons, stronger is the resultant bonding. Since the enthalpy of atomisation is an important factor in determining the standard electrode potential of a metal, metals with very high enthalpy of atomisation (i.e., very high boiling point) tend to be noble in their reactions (see later for electrode potentials). Another generalisation that may be drawn from Fig. 4.2 is that the metals of the second and third series have greater enthalpies of atomisation than the corresponding elements of the first series; this is an important factor in accounting for the occurrence of much more frequent metal – metal bonding in compounds of the heavy transition metals. –1 D H /kJ mol V a Fig. 4.2 Trends in enthalpies of atomisation of transition elements 4.3.2 Variation in In general, ions of the same charge in a given series show progressive Atomic and decrease in radius with increasing atomic number. This is because the Ionic Sizes new electron enters a d orbital each time the nuclear charge increases of by unity. It may be recalled that the shielding effect of a d electron is Transition not that effective, hence the net electrostatic attraction between the Metals nuclear charge and the outermost electron increases and the ionic radius decreases. The same trend is observed in the atomic radii of a given series. However, the variation within a series is quite small. An interesting point emerges when atomic sizes of one series are compared with those of the corresponding elements in the other series. The curves in Fig. 4.3 show an increase from the first (3d) to the second (4d) series of the elements but the radii of the third (5d) series are virtually the same as those of the corresponding members of the second series. This phenomenon is associated with the intervention of the 4f orbitals which must be filled before the 5d series of elements begin. The filling of 4f before 5d orbital results in a regular decrease in atomic radii called Lanthanoid contraction which essentially compensates for the expected 93 The d- and f- Block Elements Reprint 2024-25  increase in atomic size with increasing atomic number. The net result of the lanthanoid contraction is that the second and the third d series exhibit similar radii (e.g., Zr 160 pm, Hf 159 pm) and have very similar physical and chemical properties much more than that expected on the basis of usual family relationship. 19 The factor responsible for the lanthanoid 18 contraction is somewhat similar to that observed in an ordinary transition series and is attributed 17 Radius/nm to similar cause, i.e., the imperfect shielding of 16 one electron by another in the same set of orbitals. However, the shielding of one 4f electron by 15 another is less than that of one d electron by 14 another, and as the nuclear charge increases 13 along the series, there is fairly regular decrease in the size of the entire 4f n orbitals. 12 Sc Ti V Cr Mn Fe Co Ni Cu Zn The decrease in metallic radius coupled with Y Zr Nb Mo Tc Ru Rh Pd Ag Cd increase in atomic mass results in a general La Hf Ta W Re Os Ir Pt Au Hg increase in the density of these elements. Thus, from titanium (Z = 22) to copper (Z = 29) the Fig. 4.3: Trends in atomic radii of significant increase in the density may be noted transition elements (Table 4.2). Table 4.2: Electronic Configurations and some other Properties of the First Series of Transition Elements Element Sc Ti V Cr Mn Fe Co Ni Cu Zn Atomic number 21 22 23 24 25 26 27 28 29 30 Electronic configuration M 3d14s2 3d24s2 3d34s2 3d 54s 1 3d54s2 3d64s2 3d74s2 3d84s2 3d 104s 1 3d 104s 2 + 1 1 2 1 3 1 5 5 1 6 1 7 1 8 1 M 3d 4s 3d 4s 3d 4s 3d 3d 4s 3d 4s 3d 4s 3d 4s 3d 10 3d 104s 1 2+ 1 2 3 4 5 6 7 8 9 M 3d 3d 3d 3d 3d 3d 3d 3d 3d 3d 10 3+ 1 2 3 4 5 6 7 M [Ar] 3d 3d 3d 3d 3d 3d 3d – – –1 Enthalpy of atomisation, Da H o/kJ mol 326 473 515 397 281 416 425 430 339 126 –1 Ionisation enthalpy/D Di H o/kJ mol D iH o I 631 656 650 653 717 762 758 736 745 906 D iH o II 1235 1309 1414 1592 1509 1561 1644 1752 1958 1734 D iH o III 2393 2657 2833 2990 3260 2962 3243 3402 3556 3837 Metallic/ionic M 164 147 135 129 137 126 125 125 128 137 radii/pm M 2+ – – 79 82 82 77 74 70 73 75 3+ M 73 67 64 62 65 65 61 60 – – Standard electrode M 2+/M – –1.63 –1.18 –0.90 –1.18 –0.44 –0.28 –0.25 +0.34 -0.76 potential E o/V M 3+/M 2+ – –0.37 –0.26 –0.41 +1.57 +0.77 +1.97 – – – –3 Density/g cm 3.43 4.1 6.07 7.19 7.21 7.8 8.7 8.9 8.9 7.1 Chemistry 94 Reprint 2024-25  Why do the transition elements exhibit higher enthalpies of Example 4.2 atomisation? Because of large number of unpaired electrons in their atoms they Solution have stronger interatomic interaction and hence stronger bonding between atoms resulting in higher enthalpies of atomisation. Intext Question 4.2 In the series Sc (Z = 21) to Zn (Z = 30), the enthalpy of atomisation –1 of zinc is the lowest, i.e., 126 kJ mol. Why? 4.3.3 Ionisation There is an increase in ionisation enthalpy along each series of the Enthalpies transition elements from left to right due to an increase in nuclear charge which accompanies the filling of the inner d orbitals. Table 4.2 gives the values of the first three ionisation enthalpies of the first series of transition elements. These values show that the successive enthalpies of these elements do not increase as steeply as in the case of non-transition elements. The variation in ionisation enthalpy along a series of transition elements is much less in comparison to the variation along a period of non-transition elements. The first ionisation enthalpy, in general, increases, but the magnitude of the increase in the second and third ionisation enthalpies for the successive elements, is much higher along a series. The irregular trend in the first ionisation enthalpy of the metals of 3d series, though of little chemical significance, can be accounted for by considering that the removal of one electron alters the relative energies of 4s and 3d orbitals. You have learnt that when d-block elements form ions, ns electrons are lost before (n – 1) d electrons. As we move along the period in 3d series, we see that nuclear charge increases from scandium to zinc but electrons are added to the orbital of inner subshell, i.e., 3d orbitals. These 3d electrons shield the 4s electrons from the increasing nuclear charge somewhat more effectively than the outer shell electrons can shield one another. Therefore, the atomic radii decrease less rapidly. Thus, ionization energies increase only slightly n along the 3d series. The doubly or more highly charged ions have d configurations with no 4s electrons. A general trend of increasing values of second ionisation enthalpy is expected as the effective nuclear charge increases because one d electron does not shield another electron from the influence of nuclear charge because d-orbitals differ in direction. However, the trend of steady increase in second and third ionisation enthalpy breaks for the formation of Mn2+ and Fe3+ respectively. In both the cases, ions have d5 configuration. Similar breaks occur at corresponding elements in the later transition series. The interpretation of variation in ionisation enthalpy for an electronic configuration dn is as follows: The three terms responsible for the value of ionisation enthalpy are attraction of each electron towards nucleus, repulsion between the 95 The d- and f- Block Elements Reprint 2024-25  electrons and the exchange energy. Exchange energy is responsible for the stabilisation of energy state. Exchange energy is approximately proportional to the total number of possible pairs of parallel spins in the degenerate orbitals. When several electrons occupy a set of degenerate orbitals, the lowest energy state corresponds to the maximum possible extent of single occupation of orbital and parallel spins (Hunds rule). The loss of exchange energy increases the stability. As the stability increases, the ionisation becomes more difficult. There is no loss of exchange energy at d6 configuration. Mn+ has 3d54s1 configuration and configuration of Cr+ is d5, therefore, ionisation enthalpy of Mn+ is lower than Cr+. In the same way, Fe2+ has d6 configuration and Mn2+ has 3d5 configuration. Hence, ionisation enthalpy of Fe2+ is lower than the Mn2+. In other words, we can say that the third ionisation enthalpy of Fe is lower than that of Mn. The lowest common oxidation state of these metals is +2. To 2+ form the M ions from the gaseous atoms, the sum of the first and second ionisation enthalpy is required in addition to the enthalpy of atomisation. The dominant term is the second ionisation enthalpy + which shows unusually high values for Cr and Cu where M ions 5 10 have the d and d configurations respectively. The value for Zn is correspondingly low as the ionisation causes the removal of one 4s 10 electron which results in the formation of stable d configuration. The trend in the third ionisation enthalpies is not complicated by the 4s orbital factor and shows the greater difficulty of removing an 5 2+ 10 2+ electron from the d (Mn ) and d (Zn ) ions. In general, the third ionisation enthalpies are quite high. Also the high values for third ionisation enthalpies of copper, nickel and zinc indicate why it is difficult to obtain oxidation state greater than two for these elements. Although ionisation enthalpies give some guidance concerning the relative stabilities of oxidation states, this problem is very complex and not amenable to ready generalisation. 4.3.4 Oxidation One of the notable features of a transition elements is the great variety States of oxidation states these may show in their compounds. Table 4.3 lists the common oxidation states of the first row transition elements. Table 4.3: Oxidation States of the first row Transition Metal (the most common ones are in bold types) Sc Ti V Cr Mn Fe Co Ni Cu Zn +2 +2 +2 +2 +2 +2 +2 +1 +2 +3 +3 +3 +3 +3 +3 +3 +3 +2 +4 +4 +4 +4 +4 +4 +4 +5 +5 +5 +6 +6 +6 +7 Chemistry 96 Reprint 2024-25  The elements which give the greatest number of oxidation states occur in or near the middle of the series. Manganese, for example, exhibits all the oxidation states from +2 to +7. The lesser number of oxidation states at the extreme ends stems from either too few electrons to lose or share (Sc, Ti) or too many d electrons (hence fewer orbitals available in which to share electrons with others) for higher valence (Cu, Zn). Thus, early in the series scandium(II) is virtually unknown and titanium (IV) is more stable than Ti(III) or Ti(II). At the other end, the only oxidation state of zinc is +2 (no d electrons are involved). The maximum oxidation states of reasonable stability correspond in value IV V + to the sum of the s and d electrons upto manganese (Ti O2, V O2 , V1 2– VII – Cr O4 , Mn O4 ) followed by a rather abrupt decrease in stability of II,III higher oxidation states, so that the typical species to follow are Fe , II,III II I,II II Co , Ni , Cu , Zn. The variability of oxidation states, a characteristic of transition elements, arises out of incomplete filling of d orbitals in such a way II III that their oxidation states differ from each other by unity, e.g., V , V , IV V V , V. This is in contrast with the variability of oxidation states of non transition elements where oxidation states normally differ by a unit of two. An interesting feature in the variability of oxidation states of the d– block elements is noticed among the groups (groups 4 through 10). Although in the p–block the lower oxidation states are favoured by the heavier members (due to inert pair effect), the opposite is true in the groups of d-block. For example, in group 6, Mo(VI) and W(VI) are found to be more stable than Cr(VI). Thus Cr(VI) in the form of dichromate in acidic medium is a strong oxidising agent, whereas MoO3 and WO3 are not. Low oxidation states are found when a complex compound has ligands capable of p-acceptor character in addition to the s-bonding. For example, in Ni(CO)4 and Fe(CO)5, the oxidation state of nickel and iron is zero. Name a transition element which does not exhibit variable Example 4.3 oxidation states. Scandium (Z = 21) does not exhibit variable oxidation states. Solution Intext Question 4.3 Which of the 3d series of the transition metals exhibits the largest number of oxidation states and why? 97 The d- and f- Block Elements Reprint 2024-25 4.3.5 Trends in the Table 4.4 contains the thermochemical parameters related to the 2+ 2+ M /M transformation of the solid metal atoms to M ions in solution and their V Standard standard electrode potentials. The observed values of E and those Electrode calculated using the data of Table 4.4 are compared in Fig. 4.4. V Potentials The unique behaviour of Cu, having a positive E , accounts for its inability to liberate H2 from acids. Only oxidising acids (nitric and hot concentrated sulphuric) react with Cu, the acids being reduced. The high energy to transform Cu(s) to Cu2+(aq) is not balanced by its hydration V enthalpy. The general trend towards less negative E values across the Fig. 4.4: Observed and calculated values for the standard electrode potentials (M2+ ® M°) of the elements Ti to Zn series is related to the general increase in the sum of the first and second V ionisation enthalpies. It is interesting to note that the value of E for Mn, Ni and Zn are more negative than expected from the trend. 2+ 3+ 4 Why is Cr reducing and Mn oxidising when both have d configuration? Example 4.4 Cr is reducing as its configuration changes from d to d , the latter Solution 2+ 4 3 having a half-filled t2g level (see Unit 5). On the other hand, the change 3+ 2+ 5 from Mn to Mn results in the half-filled (d ) configuration which has extra stability. Intext Question o 2+ 4.4 The E (M /M) value for copper is positive (+0.34V). What is possible reason for this? (Hint: consider its high DaH o and low DhydH o) Chemistry 98 Reprint 2024-25  Table 4.4: Thermochemical data (kJ mol-1) for the first row Transition Elements and the Standard Electrode Potentials for the Reduction of MII to M. o o o o o Element (M) D aH (M) D iH 1 D 1H2 D hydH (M2+) E /V Ti 469 656 1309 -1866 -1.63 V 515 650 1414 -1895 -1.18 Cr 398 653 1592 -1925 -0.90 Mn 279 717 1509 -1862 -1.18 Fe 418 762 1561 -1998 -0.44 Co 427 758 1644 -2079 -0.28 Ni 431 736 1752 -2121 -0.25 Cu 339 745 1958 -2121 0.34 Zn 130 906 1734 -2059 -0.76 2+ The stability of the half-filled d sub-shell in Mn and the completely 10 2+ filled d configuration in Zn are related to their E o values, whereas E o o for Ni is related to the highest negative DhydH. 3+ 2+ 4.3.6 Trends in An examination of the E o(M /M ) values (Table 4.2) shows the varying 3+ the M3+/M2+ trends. The low value for Sc reflects the stability of Sc which has a Standard noble gas configuration. The highest value for Zn is due to the removal 10 2+ Electrode of an electron from the stable d configuration of Zn. The 2+ 5 Potentials comparatively high value for Mn shows that Mn (d ) is particularly stable, whereas comparatively low value for Fe shows the extra stability 3+ 5 of Fe (d ). The comparatively low value for V is related to the stability 2+ of V (half-filled t2g level, Unit 5). 4.3.7 Trends in Table 4.5 shows the stable halides of the 3d series of transition metals. Stability of The highest oxidation numbers are achieved in TiX4 (tetrahalides), VF5 Higher and CrF6. The +7 state for Mn is not represented in simple halides but Oxidation MnO3F is known, and beyond Mn no metal has a trihalide except FeX3 States and CoF3. The ability of fluorine to stabilise the highest oxidation state is due to either higher lattice energy as in the case of CoF3, or higher bond enthalpy terms for the higher covalent compounds, e.g., VF5 and CrF6. +5 Although V is represented only by VF5, the other halides, however, undergo hydrolysis to give oxohalides, VOX3. Another feature of fluorides is their instability in the low oxidation states e.g., VX2 (X = CI, Br or I) Table 4.5: Formulas of Halides of 3d Metals Oxidation Number + 6 CrF6 + 5 VF5 CrF5 + 4 TiX4 VXI4 CrX4 MnF4 + 3 TiX3 VX3 CrX3 MnF3 FeXI3 CoF3 III II + 2 TiX2 VX2 CrX2 MnX2 FeX2 CoX2 NiX2 CuX2 ZnX2 III + 1 CuX Key: X = F ® I; XI = F ® Br; XII = F, CI; XIII = CI ® I 99 The d- and f- Block Elements Reprint 2024-25  II and the same applies to CuX. On the other hand, all Cu halides are 2+ – known except the iodide. In this case, Cu oxidises I to I2: 2Cu2  4I Cu2 I2  s   I2 However, many copper (I) compounds are unstable in aqueous solution and undergo disproportionation. + 2+ 2Cu ® Cu + Cu 2+ + The stability of Cu (aq) rather than Cu (aq) is due to the much 2+ + more negative DhydH o of Cu (aq) than Cu , which more than compensates for the second ionisation enthalpy of Cu. The ability of oxygen to stabilise the highest oxidation state is demonstrated in the oxides. The highest oxidation number in the oxides (Table 4.6) coincides with the group number and is attained in Sc2O3 to Mn2O7. Beyond Group 7, no higher oxides of Fe above Fe2O3, are known, although ferrates (VI)(FeO4)2–, are formed in alkaline media but they readily decompose to Fe2O3 and O2. Besides the oxides, oxocations stabilise Vv as VO2+, VIV as VO2+ and TiIV as TiO2+. The ability of oxygen to stabilise these high oxidation states exceeds that of fluorine. Thus the highest Mn fluoride is MnF4 whereas the highest oxide is Mn2O7. The ability of oxygen to form multiple bonds to metals explains its superiority. In the covalent oxide Mn2O7, each Mn is tetrahedrally surrounded by O’s including a Mn–O–Mn bridge. The tetrahedral [MO4]n- ions are known for VV, CrVl, MnV, MnVl and MnVII. Table 4.6: Oxides of 3d Metals Oxidation Groups Number 3 4 5 6 7 8 9 10 11 12 + 7 Mn2O7 + 6 CrO3 + 5 V2O5 + 4 TiO2 V2O4 CrO2 MnO2 + 3 Sc2O3 Ti2O3 V2O3 Cr2O3 Mn2O3 Fe2O3 * * * Mn3O4 Fe3O4 Co3O4 + 2 TiO VO (CrO) MnO FeO CoO NiO CuO ZnO + 1 Cu2O * mixed oxides How would you account for the increasing oxidising power in the Example 4.5 series VO2+ < Cr2O72– < MnO4 – ? This is due to the increasing stability of the lower species to which they Solution are reduced. Intext Question 4.5 How would you account for the irregular variation of ionisation enthalpies (first and second) in the first series of the transition elements? Chemistry 100 Reprint 2024-25 4.3.8 Chemical Transition metals vary widely in their chemical reactivity. Many of Reactivity them are sufficiently electropositive to dissolve in mineral acids, although and Eo a few are ‘noble’—that is, they are unaffected by single acids. Values The metals of the first series with the exception of copper are relatively more reactive and are oxidised by 1M H+, though the actual rate at + which these metals react with oxidising agents like hydrogen ion (H ) is sometimes slow. For example, titanium and vanadium, in practice, are passive to dilute non oxidising acids at room temperature. The E o values for M2+/M (Table 4.2) indicate a decreasing tendency to form divalent cations across the series. This general trend towards less negative E o values is related to the increase in the sum of the first and second ionisation enthalpies. It is interesting to note that the E o values for Mn, Ni and Zn are more negative than expected from the general trend. 5 2+ Whereas the stabilities of half-filled d subshell (d ) in Mn and completely filled d subshell (d ) in zinc are related to their E values; for nickel, E o 10 e value is related to the highest negative enthalpy of hydration. 3+ 2+ An examination of the E o values for the redox couple M /M (Table 3+ 3+ 4.2) shows that Mn and Co ions are the strongest oxidising agents 2+ 2+ 2+ in aqueous solutions. The ions Ti , V and Cr are strong reducing agents and will liberate hydrogen from a dilute acid, e.g., 2+ + 3+ 2 Cr (aq) + 2 H (aq) ® 2 Cr (aq) + H2(g) Example 4.6 For the first row transition metals the E o values are: o E V Cr Mn Fe Co Ni Cu 2+ (M /M) –1.18 – 0.91 –1.18 – 0.44 – 0.28 – 0.25 +0.34 Explain the irregularity in the above values. Solution The Eo (M2+/M) values are not regular which can be explained from the irregular variation of ionisation enthalpies ( i H1   i H 2 ) and also the sublimation enthalpies which are relatively much less for manganese and vanadium. Example 4.7 Why is the E o value for the Mn3+/Mn2+ couple much more positive 3+ 2+ 3+ 2+ than that for Cr /Cr or Fe /Fe ? Explain. Solution Much larger third ionisation energy of Mn (where the required change 5 4 is d to d ) is mainly responsible for this. This also explains why the +3 state of Mn is of little importance. Intext Questions 4.6 Why is the highest oxidation state of a metal exhibited in its oxide or fluoride only? 2+ 2+ 4.7 Which is a stronger reducing agent Cr or Fe and why ? 4.3.9 Magnetic When a magnetic field is applied to substances, mainly two types of Properties magnetic behaviour are observed: diamagnetism and paramagnetism. Diamagnetic substances are repelled by the applied field while the paramagnetic substances are attracted. Substances which are 101 The d- and f- Block Elements Reprint 2024-25  attracted very strongly are said to be ferromagnetic. In fact, ferromagnetism is an extreme form of paramagnetism. Many of the transition metal ions are paramagnetic. Paramagnetism arises from the presence of unpaired electrons, each such electron having a magnetic moment associated with its spin angular momentum and orbital angular momentum. For the compounds of the first series of transition metals, the contribution of the orbital angular momentum is effectively quenched and hence is of no significance. For these, the magnetic moment is determined by the number of unpaired electrons and is calculated by using the ‘spin-only’ formula, i.e.,   n n  2 where n is the number of unpaired electrons and µ is the magnetic moment in units of Bohr magneton (BM). A single unpaired electron has a magnetic moment of 1.73 Bohr magnetons (BM). The magnetic moment increases with the increasing number of unpaired electrons. Thus, the observed magnetic moment gives a useful indication about the number of unpaired electrons present in the atom, molecule or ion. The magnetic moments calculated from the ‘spin-only’ formula and those derived experimentally for some ions of the first row transition elements are given in Table 4.7. The experimental data are mainly for hydrated ions in solution or in the solid state. Table 4.7: Calculated and Observed Magnetic Moments (BM) Ion Configuration Unpaired Magnetic moment electron(s) Calculated Observed 3+ 0 Sc 3d 0 0 0 3+ 1 Ti 3d 1 1.73 1.75 2+ 2 Tl 3d 2 2.84 2.76 2+ 3 V 3d 3 3.87 3.86 2+ 4 Cr 3d 4 4.90 4.80 2+ 5 Mn 3d 5 5.92 5.96 2+ 6 Fe 3d 4 4.90 5.3 – 5.5 2+ 7 Co 3d 3 3.87 4.4 – 5.2 2+ 8 Ni 3d 2 2.84 2.9 – 3, 4 2+ 9 Cu 3d 1 1.73 1.8 – 2.2 2+ 10 Zn 3d 0 0 Calculate the magnetic moment of a divalent ion in aqueous solution Example 4.8 if its atomic number is 25. With atomic number 25, the divalent ion in aqueous solution will have Solution 5 d configuration (five unpaired electrons). The magnetic moment, µ is   5  5  2   5.92 BM Chemistry 102 Reprint 2024-25  Intext Question 2+ 4.8 Calculate the ‘spin only’ magnetic moment of M (aq) ion (Z = 27). 4.3.10 Formation When an electron from a lower energy d orbital is excited to a higher of Coloured energy d orbital, the energy of excitation corresponds to the frequency Ions of light absorbed (Unit 5). This frequency generally lies in the visible region. The colour observed corresponds to the complementary colour of the light absorbed. The frequency of the light absorbed is determined by the nature of the ligand. In aqueous solutions where water molecules are the ligands, the colours of the ions observed are listed in Table 4.8. A few coloured solutions of Fig. 4.5: Colours of some of the first row d–block elements are transition metal ions in aqueous solutions. From illustrated in Fig. 4.5. left to right: V4+,V3+,Mn2+,Fe3+,Co2+,Ni2+and Cu2+. Table 4.8: Colours of Some of the First Row (aquated) Transition Metal Ions Configuration Example Colour 0 3+ 3d Sc colourless 3d0 Ti4+ colourless 3d1 Ti3+ purple 3d1 V4+ blue 2 3+ 3d V green 3 2+ 3d V violet 3 3+ 3d Cr violet 4 3+ 3d Mn violet 4 2+ 3d Cr blue 5 2+ 3d Mn pink 5 3+ 3d Fe yellow 6 2+ 3d Fe green 6 7 3+ 2+ 3d 3d Co Co bluepink 3d8 Ni2+ green 3d9 Cu2+ blue 3d10 Zn2+ colourless 4.3.11 Formation Complex compounds are those in which the metal ions bind a number of Complex of anions or neutral molecules giving complex species with 3– 4– Compounds characteristic properties. A few examples are: [Fe(CN)6] , [Fe(CN)6] , 2+ 2– [Cu(NH3)4] and [PtCl4]. (The chemistry of complex compounds is 103 The d- and f- Block Elements Reprint 2024-25  dealt with in detail in Unit 5). The transition metals form a large number of complex compounds. This is due to the comparatively smaller sizes of the metal ions, their high ionic charges and the availability of d orbitals for bond formation. 4.3.12 Catalytic The transition metals and their compounds are known for their catalytic Properties activity. This activity is ascribed to their ability to adopt multiple oxidation states and to form complexes. Vanadium(V) oxide (in Contact Process), finely divided iron (in Haber’s Process), and nickel (in Catalytic Hydrogenation) are some of the examples. Catalysts at a solid surface involve the formation of bonds between reactant molecules and atoms of the surface of the catalyst (first row transition metals utilise 3d and 4s electrons for bonding). This has the effect of increasing the concentration of the reactants at the catalyst surface and also weakening of the bonds in the reacting molecules (the activation energy is lowering). Also because the transition metal ions can change their oxidation states, they become more effective as catalysts. For example, iron(III) catalyses the reaction between iodide and persulphate ions. – 2– 2– 2 I + S2O8 ® I2 + 2 SO4 An explanation of this catalytic action can be given as: 2 Fe3+ + 2 I– ® 2 Fe2+ + I2 2+ 2– 3+ 2– 2 Fe + S2O8 ® 2 Fe + 2SO4 4.3.13 Formation Interstitial compounds are those which are formed when small atoms of like H, C or N are trapped inside the crystal lattices of metals. They are Interstitial usually non stoichiometric and are neither typically ionic nor covalent, Compounds for example, TiC, Mn4N, Fe3H, VH0.56 and TiH1.7, etc. The formulas quoted do not, of course, correspond to any normal oxidation state of the metal. Because of the nature of their composition, these compounds are referred to as interstitial compounds. The principal physical and chemical characteristics of these compounds are as follows: (i) They have high melting points, higher than those of pure metals. (ii) They are very hard, some borides approach diamond in hardness. (iii) They retain metallic conductivity. (iv) They are chemically inert. 4.3.14 Alloy An alloy is a blend of metals prepared by mixing the components. Formation Alloys may be homogeneous solid solutions in which the atoms of one metal are distributed randomly among the atoms of the other. Such alloys are formed by atoms with metallic radii that are within about 15 percent of each other. Because of similar radii and other characteristics of transition metals, alloys are readily formed by these metals. The alloys so formed are hard and have often high melting points. The best known are ferrous alloys: chromium, vanadium, tungsten, molybdenum and manganese are used for the production of a variety of steels and stainless steel. Alloys of transition metals with non transition metals such as brass (copper-zinc) and bronze (copper-tin), are also of considerable industrial importance. Chemistry 104 Reprint 2024-25  Example 4.9 What is meant by ‘disproportionation’ of an oxidation state? Give an example. Solution When a particular oxidation state becomes less stable relative to other oxidation states, one lower, one higher, it is said to undergo disproportionation. For example, manganese (VI) becomes unstable relative to manganese(VII) and manganese (IV) in acidic solution. 3 MnVIO4 2– + 4 H+ ® 2 MnVIIO–4 + MnIVO2 + 2H2O Intext Question + 4.9 Explain why Cu ion is not stable in aqueous solutions? 4. 4 Some 4.4.1 Oxides and Oxoanions of Metals Important These oxides are generally formed by the reaction of metals with Compounds of oxygen at high temperatures. All the metals except scandium form MO oxides which are ionic. The highest oxidation number in the Transition oxides, coincides with the group number and is attained in Sc2O3 to Elements Mn2O7. Beyond group 7, no higher oxides of iron above Fe2O3 are V + IV known. Besides the oxides, the oxocations stabilise V as VO2 , V as 2+ IV 2+ VO and Ti as TiO. As the oxidation number of a metal increases, ionic character decreases. In the case of Mn, Mn2O7 is a covalent green oil. Even CrO3 and V2O5 have low melting points. In these higher oxides, the acidic character is predominant. Thus, Mn2O7 gives HMnO4 and CrO3 gives H2CrO4 and H2Cr2O7. 3– V2O5 is, however, amphoteric though mainly acidic and it gives VO4 as + well as VO2 salts. In vanadium there is gradual change from the basic V2O3 to less basic V2O4 and to amphoteric V2O5. V2O4 dissolves in acids 2+ to give VO salts. Similarly, V2O5 reacts with alkalies as well as acids to give VO34 and VO4 respectively. The well characterised CrO is basic but Cr2O3 is amphoteric. Potassium dichromate K2Cr2O7 Potassium dichromate is a very important chemical used in leather industry and as an oxidant for preparation of many azo compounds. Dichromates are generally prepared from chromate, which in turn are obtained by the fusion of chromite ore (FeCr2O4) with sodium or potassium carbonate in free access of air. The reaction with sodium carbonate occurs as follows: 4 FeCr2O4 + 8 Na2CO3 + 7 O2 ® 8 Na2CrO4 + 2 Fe2O3 + 8 CO2 The yellow solution of sodium chromate is filtered and acidified with sulphuric acid to give a solution from which orange sodium dichromate, Na2Cr2O7. 2H2O can be crystallised. + + 2Na2CrO4 + 2 H ® Na2Cr2O7 + 2 Na + H2O 105 The d- and f- Block Elements Reprint 2024-25  Sodium dichromate is more soluble than potassium dichromate. The latter is therefore, prepared by treating the solution of sodium dichromate with potassium chloride. Na2Cr2O7 + 2 KCl ® K2Cr2O7 + 2 NaCl Orange crystals of potassium dichromate crystallise out. The chromates and dichromates are interconvertible in aqueous solution depending upon pH of the solution. The oxidation state of chromium in chromate and dichromate is the same. 2– + 2– 2 CrO4 + 2H ® Cr2O7 + H2O 2– - 2– Cr2O7 + 2 OH ® 2 CrO4 + H2O The structures of 2– chromate ion, CrO4 and 2– the dichromate ion, Cr2O7 are shown below. The chromate ion is tetrahedral whereas the dichromate ion consists of two tetrahedra sharing one corner with Cr–O–Cr bond angle of 126°. Sodium and potassium dichromates are strong oxidising agents; the sodium salt has a greater solubility in water and is extensively used as an oxidising agent in organic chemistry. Potassium dichromate is used as a primary standard in volumetric analysis. In acidic solution, its oxidising action can be represented as follows: 2– + – 3+ Cr2O7 + 14H + 6e ® 2Cr + 7H2O (E o = 1.33V) Thus, acidified potassium dichromate will oxidise iodides to iodine, sulphides to sulphur, tin(II) to tin(IV) and iron(II) salts to iron(III). The half-reactions are noted below: – – 2+ 4+ – 6 I ® 3I2 + 6 e ; 3 Sn ® 3Sn + 6 e + – 2+ 3+ – 3 H2S ® 6H + 3S + 6e ; 6 Fe ® 6Fe + 6 e The full ionic equation may be obtained by adding the half-reaction for potassium dichromate to the half-reaction for the reducing agent, for e.g., 2– + 2+ 3+ 3+ Cr2O7 + 14 H + 6 Fe ® 2 Cr + 6 Fe + 7 H2O Potassium permanganate KMnO4 Potassium permanganate is prepared by fusion of MnO2 with an alkali metal hydroxide and an oxidising agent like KNO3. This produces the dark green K2MnO4 which disproportionates in a neutral or acidic solution to give permanganate. 2MnO2 + 4KOH + O2 ® 2K2MnO4 + 2H2O 3MnO42– + 4H+ ® 2MnO4– + MnO2 + 2H2O Commercially it is prepared by the alkaline oxidative fusion of MnO2 followed by the electrolytic oxidation of manganate (Vl). Fused with KOH, oxidised Electrolytic oxidation in with air or KNO3 MnO2 → MnO24− 2 alkaline solution ; MnO  4 MnO4 manganate ion manganate permanganate ion Chemistry 106 Reprint 2024-25  In the laboratory, a manganese (II) ion salt is oxidised by peroxodisulphate to permanganate. 2+ 2– – 2– + 2Mn + 5S2O8 + 8H2O ® 2MnO4 + 10SO4 + 16H Potassium permanganate forms dark purple (almost black) crystals which are isostructural with those of KClO4. The salt is not very soluble in water (6.4 g/100 g of water at 293 K), but when heated it decomposes at 513 K. 2KMnO4 ® K2MnO4 + MnO2 + O2 It has two physical properties of considerable interest: its intense colour and its diamagnetism along with temperature-dependent weak paramagnetism. These can be explained by the use of molecular orbital theory which is beyond the present scope. The manganate and permanganate ions are tetrahedral; the p- bonding takes place by overlap of p orbitals of oxygen with d orbitals of manganese. The green manganate is paramagnetic because of one unpaired electron but the permanganate is diamagnetic due to the absence of unpaired electron. Acidified permanganate solution oxidises oxalates to carbon dioxide, iron(II) to iron(III), nitrites to nitrates and iodides to free iodine. The half-reactions of reductants are: COO– 5 10 CO2 + 10e– – COO 5 Fe2+ ® 5 Fe3+ + 5e– 5NO2– + 5H2O ® 5NO3– + 10H+ + l0e– 10I– ® 5I2 + 10e– The full reaction can be written by adding the half-reaction for KMnO4 to the half-reaction of the reducing agent, balancing wherever necessary. If we represent the reduction of permanganate to manganate, manganese dioxide and manganese(II) salt by half-reactions, MnO4– + e– ® MnO42– (E o = + 0.56 V) – + – MnO4 + 4H + 3e ® MnO2 + 2H2O (E o = + 1.69 V) o MnO4– + 8H+ + 5e– ® Mn2+ + 4H2O (E = + 1.52 V) We can very well see that the hydrogen ion concentration of the solution plays an important part in influencing the reaction. Although many reactions can be understood by consideration of redox potential, kinetics of the reaction is also an important factor. Permanganate at + [H ] = 1 should oxidise water but in practice the reaction is extremely slow unless either manganese(ll) ions are present or the temperature is raised. A few important oxidising reactions of KMnO4 are given below: 1. In acid solutions: (a) Iodine is liberated from potassium iodide : – – + 2+ 10I + 2MnO4 + 16H ® 2Mn + 8H2O + 5I2 2+ 3+ (b) Fe ion (green) is converted to Fe (yellow): 2+ – + 2+ 3+ 5Fe + MnO4 + 8H ® Mn + 4H2O + 5Fe 107 The d- and f- Block Elements Reprint 2024-25  (c) Oxalate ion or oxalic acid is oxidised at 333 K: 5C2O42– + 2MnO4– + 16H+ ——> 2Mn2+ + 8H2O + 10CO2 (d) Hydrogen sulphide is oxidised, sulphur being precipitated: + 2– H2S —> 2H + S 2– – + 2+ 5S + 2MnO 4 + 16H ——> 2Mn + 8H2O + 5S (e) Sulphurous acid or sulphite is oxidised to a sulphate or sulphuric acid: 2– – + 2+ 2– 5SO3 + 2MnO4 + 6H ——> 2Mn + 3H2O + 5SO4 (f) Nitrite is oxidised to nitrate: 5NO2– + 2MnO4– + 6H+ ——> 2Mn2+ + 5NO3– + 3H2O 2. In neutral or faintly alkaline solutions: (a) A notable reaction is the oxidation of iodide to iodate: – – – – 2MnO4 + H2O + I ——> 2MnO2 + 2OH + IO3 (b) Thiosulphate is oxidised almost quantitatively to sulphate: 8MnO4– + 3S2O32– + H2O ——> 8MnO2 + 6SO42– + 2OH– (c) Manganous salt is oxidised to MnO2; the presence of zinc sulphate or zinc oxide catalyses the oxidation: 2MnO4– + 3Mn2+ + 2H2O ——> 5MnO2 + 4H+ Note: Permanganate titrations in presence of hydrochloric acid are unsatisfactory since hydrochloric acid is oxidised to chlorine. Uses Uses: Besides its use in analytical chemistry, potassium permanganate is used as a favourite oxidant in preparative organic chemistry. Its uses for the bleaching of wool, cotton, silk and other textile fibres and for the decolourisation of oils are also dependent on its strong oxidising power. THE INNER TRANSITION ELEMENTS ( f-BLOCK) The f-block consists of the two series, lanthanoids (the fourteen elements following lanthanum) and actinoids (the fourteen elements following actinium). Because lanthanum closely resembles the lanthanoids, it is usually included in any discussion of the lanthanoids for which the general symbol Ln is often used. Similarly, a discussion of the actinoids includes actinium besides the fourteen elements constituting the series. The lanthanoids resemble one another more closely than do the members of ordinary transition elements in any series. They have only one stable oxidation state and their chemistry provides an excellent opportunity to examine the effect of small changes in size and nuclear charge along a series of otherwise similar elements. The chemistry of the actinoids is, on the other hand, much more complicated. The complication arises partly owing to the occurrence of a wide range of oxidation states in these elements and partly because their radioactivity creates special problems in their study; the two series will be considered separately here. 4.5 The The names, symbols, electronic configurations of atomic and some Lanthanoids ionic states and atomic and ionic radii of lanthanum and lanthanoids (for which the general symbol Ln is used) are given in Table 4.9. Chemistry 108 Reprint 2024-25 4.5.1 Electronic It may be noted that atoms of these elements have electronic 2 Configurations configuration with 6s common but with variable occupancy of 4f level (Table 4.9). However, the electronic configurations of all the tripositive ions (the most stable oxidation state of all the lanthanoids) are of the form 4f n (n = 1 to 14 with increasing atomic number). 4.5.2 Atomic and The overall decrease in atomic and ionic radii from lanthanum to Ionic Sizes lutetium (the lanthanoid contraction) is a unique feature in the chemistry of the lanthanoids. It has far reaching Sm 2+ consequences in the chemistry of the third 110 2+ transition series of the elements. The decrease Eu in atomic radii (derived from the structures of La 3+ metals) is not quite regular as it is regular in 3+ 3+ M ions (Fig. 4.6). This contraction is, of Ce course, similar to that observed in an ordinary Pr 3+ transition series and is attributed to the same 100 Nd 3+ cause, the imperfect shielding of one electron Pm 3+ by another in the same sub-shell. However, the Ionic radii/pm Sm 3+ shielding of one 4 f electron by another is less Eu 3+ than one d electron by another with the increase 2+ Tm 3+ Gd in nuclear charge along the series. There is 2+ Tb Yb 3+ Ce 4+ fairly regular decrease in the sizes with 3+ Dy Pr 4+ increasing atomic number. 3+ 90 Ho Er The cumulative effect of the contraction of 3+ Tm the lanthanoid series, known as lanthanoid 3+ 3+ Yb contraction, causes the radii of the members 3+ Lu 4+ Tb of the third transition series to be very similar to those of the corresponding members of the second series. The almost identical radii of Zr 57 59 61 63 65 67 69 71 (160 pm) and Hf (159 pm), a consequence of Atomic number the lanthanoid contraction, account for their occurrence together in nature and for the Fig. 4.6: Trends in ionic radii of lanthanoids difficulty faced in their separation. 4.5.3 Oxidation In the lanthanoids, La(II) and Ln(III) compounds are predominant States species. However, occasionally +2 and +4 ions in solution or in solid compounds are also obtained. This irregularity (as in ionisation enthalpies) arises mainly from the extra stability of empty, half-filled IV or filled f subshell. Thus, the formation of Ce is favoured by its noble gas configuration, but it is a strong oxidant reverting to the 4+ 3+ common +3 state. The E o value for Ce / Ce is + 1.74 V which suggests that it can oxidise water. However, the reaction rate is very slow and hence Ce(IV) is a good analytical reagent. Pr, Nd, Tb and Dy 2+ also exhibit +4 state but only in oxides, MO2. Eu is formed by losing 7 the two s electrons and its f configuration accounts for the formation 2+ of this ion. However, Eu is a strong reducing agent changing to the 2+ 14 common +3 state. Similarly Yb which has f configuration is a IV reductant. Tb has half-filled f-orbitals and is an oxidant. The behaviour of samarium is very much like europium, exhibiting both +2 and +3 oxidation states. 109 The d- and f- Block Elements Reprint 2024-25  Table 4.9: Electronic Configurations and Radii of Lanthanum and Lanthanoids Electronic configurations* Radii/pm 2+ 3+ 4+ 3+ Atomic Name Symbol Ln Ln Ln Ln Ln Ln Number 1 2 1 0 57 Lanthanum La 5d 6s 5d 4f 187 106 1 1 2 2 1 0 58 Cerium Ce 4f 5d 6s 4f 4f 4f 183 103 3 2 3 2 1 59 Praseodymium Pr 4f 6s 4f 4f 4f 182 101 4 2 4 3 2 60 Neodymium Nd 4f 6s 4f 4f 4f 181 99 5 2 5 4 61 Promethium Pm 4f 6s 4f 4f 181 98 6 2 6 5 62 Samarium Sm 4f 6s 4f 4f 180 96 7 2 7 6 63 Europium Eu 4f 6s 4f 4f 199 95 7 1 2 7 1 7 64 Gadolinium Gd 4f 5d 6s 4f 5d 4f 180 94 9 2 9 8 7 65 Terbium Tb 4f 6s 4f 4f 4f 178 92 10 2 10 9 8 66 Dysprosium Dy 4f 6s 4f 4f 4f 177 91 11 2 11 10 67 Holmium Ho 4f 6s 4f 4f 176 89 12 2 12 11 68 Erbium Er 4f 6s 4f 4f 175 88 13 2 13 12 69 Thulium Tm 4f 6s 4f 4f 174 87 14 2 14 13 70 Ytterbium Yb 4f 6s 4f 4f 173 86 14 1 2 14 1 14 71 Lutetium Lu 4f 5d 6s 4f 5d 4f – – – * Only electrons outside [Xe] core are indicated 4.5.4 General All the lanthanoids are silvery white soft metals and tarnish rapidly in air. Characteristics The hardness increases with increasing atomic number, samarium being steel hard. Their melting points range between 1000 to 1200 K but samarium melts at 1623 K. They have typical metallic structure and are good conductors of heat and electricity. Density and other properties change smoothly except for Eu and Yb and occasionally for Sm and Tm. Many trivalent lanthanoid ions are coloured both in the solid state and in aqueous solutions. Colour of these ions may be attributed to 3+ 3+ the presence of f electrons. Neither La nor Lu ion shows any colour but the rest do so. However, absorption bands are narrow, probably because of the excitation within f level. The lanthanoid ions other 0 3+ 4+ 14 2+ 3+ than the f type (La and Ce ) and the f type (Yb and Lu ) are all paramagnetic. The first ionisation enthalpies of the lanthanoids are around –1 –1 600 kJ mol , the second about 1200 kJ mol comparable with those of calcium. A detailed discussion of the variation of the third ionisation enthalpies indicates that the exchange enthalpy considerations (as in 3d orbitals of the first transition series), appear to impart a certain degree of stability to empty, half-filled and completely filled orbitals f level. This is indicated from the abnormally low value of the third ionisation enthalpy of lanthanum, gadolinium and lutetium. In their chemical behaviour, in general, the earlier members of the series are quite reactive similar to calcium but, with increasing atomic number, they behave more like aluminium. Values for E o for the half-reaction: 3+ – Ln (aq) + 3e ® Ln(s) Chemistry 110 Reprint 2024-25  Ln2O3 H2 are in the range of –2.2 to –2.4 V except for Eu for which the value is s – 2.0 V. This is, of course, a small id bu ac variation. The metals combine with rn th s hydrogen when gently heated in the wi in O2 gas. The carbides, Ln3C, Ln2C3 and LnC2 are formed when the metals are heated with halogen heated with S Ln s with carbon. They liberate hydrogen Ln2S3 LnX 3 from dilute acids and burn in halogens to form halides. They form oxides M2O3 N wi th th and hydroxides M(OH) 3. The wi 2773 K H2 with C ed hydroxides are definite compounds, not O at just hydrated oxides. They are basic he like alkaline earth metal oxides and LnN LnC2 Ln(OH)3 + H2 hydroxides. Their general reactions are depicted in Fig. 4.7. Fig 4.7: Chemical reactions of the lanthanoids. The best single use of the lanthanoids is for the production of alloy steels for plates and pipes. A well known alloy is mischmetall which consists of a lanthanoid metal (~ 95%) and iron (~ 5%) and traces of S, C, Ca and Al. A good deal of mischmetall is used in Mg-based alloy to produce bullets, shell and lighter flint. Mixed oxides of lanthanoids are employed as catalysts in petroleum cracking. Some individual Ln oxides are used as phosphors