d and f block elements XII (LATEST)

CBSE CHEMISTRY d AND f BLOCK ELEMENTS, CBSE 12 CHEMISTRY

1. d – AND f – BLOCK ELEMENTS   TINTO JOHNS M. Sc., M. Ed

2. The d-block of the periodic table contains the elements of the groups 3-12 in which the d orbitals are progressively filled in each of the four long periods. The elements constituting the f -block are those in which the 4 f and 5 f orbitals are progressively filled in the latter two long periods. There are mainly three series of the transition metals , 3d series (Sc to Zn ), 4d series (Y to Cd ) and 5d series (La to Hg, omitting Ce to Lu). The fourth 6d series which begins with Ac is still incomplete . The two series of the inner transition metals, (4f and 5f) are known as lanthanoids and actinoids respectively .

3. A transition element is defined as the one which has incompletely filled d orbitals in its ground state or in any one of its oxidation states. Zinc, cadmium and mercury of group 12 have full d 10 configuration in their ground state as well as in their common oxidation states and hence, are not regarded as transition metals

4. Position in the Position in the Periodic Table Periodic Table The d–block occupies the large middle section flanked by s– and p– blocks in the periodic table. Electronic Configurations of the d-Block of the d-Block Elements General the electronic configuration (n-1)d 1–10 ns 1–2 . 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. Consider the case of Cr, for example, which has 3d 5 4s 1 instead of 3d 4 4s 2 .

5. Physical properties The transition metals (with the exception of Zn, Cd and Hg) are very much hard and have low volatility. Their melting and boiling points are high.

6. MELTING POINT AND BOILING POINT High M.P and B.P - Due to strong metallic bond and the presence of half filled d- orbitals Involvement of greater number of electrons from (n-1) d in addition to the n s electrons in the inter atomic metallic bonding . Because of stronger interatomic bonding High enthalpy of atomisation transition elements have high M.P and B.P

7. In moving along the period from left to right, the M.P of these metals first INCREASES to MAXIMUM and the DECREASES regularly towards the end of the period.

8. melting points of these metals rise to a maximum at d 5 except for anomalous values of Mn and Tc and fall regularly as the atomic number increases. TRENDS OF M.P OF 3- d , 4-d AND 5-d TRANSITION METALS The strength of interatomic bonds in transition elements is roughly related to the number of half filled d- orbitals In the beginning the no. of half filled d- orbitals increases till the middle of the period causing increase in strength of interparticle bonds But thereafter the pairing of electrons in d – orbitals occurs and the no. of half filled orbitals decreases , which also cause decrease in M.P

9. Trends in enthalpies of atomization of transition elements Greater the number of valence electrons, stronger the inter atomic attraction, hence stronger bonding between atoms resulting in higher enthalpies of atomization . 2 . metals of the second and third series have greater enthalpies of atomization than the corresponding elements of the first series 3d < 4d < 5d

10. Atomic and ionic radii The Atomic/ionic radii first DECREASES till the middle, becomes almost constant and then INCREASES towards the end of the period. New electron enters a d orbital each time the nuclear charge increases by unity, But the shielding effect of a d electron is not that effective, hence the net electrostatic attraction between the nuclear charge and the outermost electron increases and the ionic radius decreases

11. However the increased nuclear charge is partly cancelled by the increased screening effect of electrons in the d – orbitals of penultimate shell. When the increased nuclear charge and increased Screening effect balance each other , the atomic radii becomes almost constant. Increase in atomic radii towards the end may be attributed to the electron – electron repulsion . In fact the pairing of electrons in d – orbitals occurs after d 5 configuration. The repulsive interaction between the paired electron causes Increase in Atomic/ ionic radii

12. Sc = [Ar]4s 2 3d 1 Ti = [Ar]4s 2 3d 2 3d< 4d= 5d

13. There is increase from the first (3 d ) to the second (4 d ) series of the elements. But the radii of the third (5 d ) series are virtually the same as 4d This is due to the intervention of the 4 f orbital which must be filled before the 5 d series of elements begin. There is a steady decrease in atomic radii from La due to the poor shielding of inner core electrons (4f) is known lanthanoid contraction .

14. Why do the transition elements exhibit higher enthalpies of atomisation? Because of large number of unpaired electrons in their atoms they have stronger interatomic interaction and hence stronger bonding between atoms resulting in higher enthalpies of atomisation.

15. IONISATION ENTHALPIES Due to an increase in nuclear charge there is an increase in ionisation enthalpy along each series of the transition elements from left to right. Ionisation enthalpies give some guidance concerning the relative stabilities of oxidation states. Although the first ionisation enthalpy, in general, increases, the magnitude of the increase in the second and third ionisation enthalpies for the successive elements, in general, is much higher. Mostly IE1<IE2 <IE3 in each group

16. The increase in IE is primarily due to increase in nuclear charge . As the transition elements involve the gradual filling of (n-1)d orbitals , the effect of increase in nuclear charge is partly cancelled by the increase in screening effect. Consequently, the increase in I.E along the periods of d – block elements is very small. 3d < 4d < 5d ( in 5d series - ineffective shielding by 4f electrons)

17. Relation between I.E and Stability of a metal in a given oxdn state With the help of I.E, we can predict which of the two metals in a given oxdn state is thermodynamically more stable. Eg When a metal M (0) is converted into M(II), the energy required is equal to I 1 + I 2 Similarly M (IV) = I 1 + I 2 + I 3 + I 4

18. Ni (0) Ni (II) I 1 + I 2 =2.49 x 10 3 kJ mol -1 P t (0) Pt (II) I 1 + I 2 =2.66 x 10 3 kJ mol -1 Ni (0) Ni (IV) I 1 + I 2 + I 3 + I 4 =11.299 x 10 3 kJ mol -1 Pt (0) Pt (IV) I 1 + I 2 + I 3 + I 4 =9.36 x 10 3 kJ mol -1 I 1 + I 2 for Ni (II) is less than I 1 + I 2 for Pt (II). So Ni (II) is more stable Similarly Pt (IV) is more stable

19. OXIDATION STATES +3 One of the notable features of a transition element is the great variety of oxidation states it may show in its compounds Stability of a particular oxdn state depends up on nature of the element with which the transition metals form the compound

20. 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 . Elements in the beginning of the series exhibit fewer oxidation state (have small no. of electrons in which they lose or contribute for sharing). Elements at the end of the series shows fewer oxdn states because they have too many electrons in d – orbitals . So they have few vacant d – orbitals which can be involved in bonding.

21. Lower oxdn state – Covalent character Higher oxdn state – ionic Higher oxdn states are more stable for heavier members. Eg : in group VI, Mo (VI) and W (VI) are more stable than Cr (VI). So Cr (VI) act as strong oxidizing agent. The highest oxdn state - +8 (Ruthenium and Osmium ). Low oxidation states are found when a complex compound has ligands capable of π-acceptor character in addition to the σ-bonding. For example, in Ni(CO) 4 and Fe(CO) 5 , the oxidation state of nickel and iron is zero.

22. Trends in Stability of Higher Oxidation States Stability – compounds with F and Oxygen The ability of Fluorine to stabilize the highest oxidation state is due to either high lattice energy as in case of CoF 3 or high bond enthalpy as in case of VF 5 and CrF 6 . The ability of Oxygen to stabilize the highest oxidation state is due to its ability to form multiple bonds with metals.

23. Stable halides of first transition elements Oxdn no. 4 5 6 7 8 9 10 11 12 +6 Cr F 6 +5 VF 5 Cr F 5 +4 Ti X 4 VX 4 I Cr X 4 MNf 4 +3 Ti X 3 VX 3 Cr X 3 MnF 3 Fe X 3 Co F 3 +2 Ti X 2 III VX 2 I Cr X 2 MnX 2 Fe X 2 Co X 2 Ni X 2 Cu X 2 II ZnX 2 +1 Cu X III X = F to I, X II = F, X I = F to Br , X III = Cl to I

24. The highest oxidation numbers are achieved in TiX 4 ( tetrahalides ), VF 5 and CrF 6 . The +7 state for Mn is not represented in simple halides but MnO 3 F is known, and beyond Mn , no metal has a trihalide except FeX 3 and CoF 3 . Although V(V)is represented only by VF 5 , the other halides, however, undergo hydrolysis to give oxohalides , VOX 3 . Another feature of fluorides is their instability in the low oxidation states e.g., VX 2 (X = CI, Br or I)

25. All Cu(II) halides are known except the iodide. In this case, Cu 2+ oxidises I – to I 2 : 2Cu 2+ + 4I - → Cu 2 I 2 ( s ) + I 2 However, many copper (I) compounds are unstable in aqueous solution and undergo disproportionation . 2Cu 2+ → Cu 2+ + Cu The stability of Cu 2+ ( aq ) rather than Cu + ( aq ) is due to the much more negative Δ hyd H of Cu 2+ ( aq ) than Cu + , which more than compensates for the second ionisation enthalpy of Cu.

26. Transition metals also exhibits the highest Oxdn state in their Oxides. The ability of Oxygen to stabilize higher oxidation states are much higher than Fluorine.. The highest Oxdn state with Fluorine by Mn is +4 in MnF 4 while it is + 7 in Mn 2 O 7 . Oxygen has the ability to form Multiple bonds with Metal atom. The oxides of 3 – d transition elements are given below :

27. Oxdn No 3 4 5 6 7 8 9 10 11 12 +7 Mn 2 O 7 +6 CrO 3 +5 V 2 O 5 MnO 2 +4 Ti O 2 V 2 O 4 CrO 2 Mn 2 O 3 Fe 2 O 3 +3 Sc 2 O 3 Ti 2 O 3 V 2 O 3 Cr 2 O 3 Mn 3 O 4 Fe 3 O 4 Co 3 O 4 +2 Ti O VO CrO MnO FeO CoO NiO CuO ZnO +1 Cu 2 O

28. The highest oxidation number in the oxides coincides with the group number and is attained in Sc 2 O 3 to Mn 2 O 7 . Beyond Group 7, no higher oxides of Fe above Fe 2 O 3 , are known , although ferrates (VI) (FeO4) 2– , are formed in alkaline media but they readily decompose to Fe 2 O 3 and O 2 . Besides the oxides, oxocations stabilise V(v) as VO 2 + , V(IV) as VO 2+ and Ti(IV) as TiO 2+.

29. STANDARD ELECTRODE POTENTIAL ELECTRODE POTENTIALS ARE THE MEASURE OF THE VALUE OF TOTAL ENTHALPY CHANGE. Electrode Potentials value depends enthalpy of atomization Δ H a & hydration Δ H hyd Lower the std E. P ( E o red ), the more stable is the oxdn state of the metal in aqueous state .

30. The E (M 2 + /M) value for copper is positive (+0.34V ) : high Δ H a and low Δ H hyd ). --- GREATER AMNT OF ENERGY REQUIRED TO TRANSFORM Cu INTO Cu 2+

31. Due to + ve E o , Cu does not liberate hydrogen from acids. The general trend towards less negative E o values across the series is related to the general increase in the sum of the first and second ionisation enthalpies. It is interesting to note that the value of E o for Mn , Ni and Zn are more negative than expected from the trend .

32. The stability of the half-filled d sub-shell in Mn 2+ and the completely filled d 10 configuration in Zn 2+ are related to their E o values, whereas E o for Ni is related to the highest negative Δ hyd H o . The low value for Sc reflects the stability of Sc 3+ which has a noble gas configuration . The highest value for Zn is due to the removal of an electron from the stable d 10 configuration of Zn 2+. The comparatively high value for Mn shows that Mn 2+ ( d 5 ) is particularly stable , whereas comparatively low value for Fe shows the extra stability of Fe 3+ ( d 5 ).

33. CHEMICAL REACTIVITY Transition metals vary widely in their chemical reactivity. Many of them are sufficiently electropositive to dissolve in mineral acids, although a few are ‘noble’—that is, they are unaffected by simple acids . 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.

34. The E O values for M 2+/ M indicate a decreasing tendency to form divalentcations 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.

35. E O values for the redox couple M3+/M2+ shows that Mn 3+ and Co 3+ ions are the strongest oxidising agents in aqueous solutions. The ions Ti 2+, V 2+ and Cr 2+ are strong reducing agents and will liberate hydrogen from a dilute acid, e.g ., 2 Cr 2+ (aq) + 2 H + (aq) → 2 Cr 3+ (aq) + H 2 (g)

36. MAGNETIC PROPERTIES Substances which contain species (Atoms/ions/molecules) with unpared electrons in their orbitals – PARAMAGNETIC. PARAMAGNETIC SUBSTANCES are weakly attracted by the magnetic field. Strongly attracted called FERROMAGNETIC. Substances which do not contain any unpaired electrons and are repelled by magnetic field - DIAMAGNETIC.

37. Transition metals usually contains unpaired electrons – so it is paramagnetic. Paramagnetic behavior increases with increase in unpaired electron. Paramagnetism expressed in terms of Magnetic moment., it is related to no. of unpaired electrons. The magnetic moments calculated from the ‘ spin-only’ formula and those derived experimentally. Magnetic moment µ = √ n(n+2) BM

38. n- no. of unpaired electrons BM – Bohr magnetone (unit of M.M) BM = 9.27x10 -21 erg/gauss Single unpaired electronhas a magnetic moment of 1.73 Bohr magnetons (BM ). magnetic moment of an electron is due to its spin angular momentum and orbital angular momentum

39. Formation of Coloured Ions When an electron from a lower energy d orbital is excited to a higher energy d orbital, the energy of excitation corresponds to the frequency of light absorbed. 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 .

40. Zn 2+ / Cd 2+ - all d orbitals are fully filled Ti 4+ - all d orbitals are vacant so, no d – d transition occurs. Therefore they do not absorb radiations. So they are colorless.

41. Formationof Complex Compounds Metal ions bind a number of anions or neutral molecules giving complex [Fe( CN ) 6 ] 3– , [Fe( CN ) 6 ] 4– , [ Cu( NH 3 ) 4 ] 2 + and [ PtCl 4 ] 2 – . 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.

42. Formation of Interstitial Compounds W hen small atoms like H, C or N are trapped inside the crystal lattices of metals They are usually non stoichiometric example, TiC , Mn 4 N , Fe 3 H , VH 0.56 and TiH 1.7 ( 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 .

43. Alloy Formation Because of similar radii and other characteristics of transition metals, The alloys so formed are hard and have often high melting points. 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),

44. CATALYTIC ACTIVITY The transition metals and their compounds are known for their catalytic activity. This activity is ascribed to their ability to adopt multiple oxidation states and to form complexes.

45. DISPROPORTIONATION 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 Mn VI O4 2– + 4 H + → 2 Mn VII O – 4 + Mn IV O 2 + 2H 2 O

46. Oxides and Oxoanions of Metals The elements of first transition series form variety of oxides of different oxidation states having general formula MO, M 2 O 3 , M 3 O 6 , MO 2 , MO 3 . Theses oxides are generally formed by heating the metal with oxygen at high temperature.

47. Sc – Sc 2 O 3 Basic Ti – TiO Basic, Ti 2 O 2 Basic, TiO 2 Amphoteric V – VO Basic, V 2 O 3 Basic, VO 2 Ampho , V 2 O 5 Acidic Cr – CrO Basic, Cr 2 O 3 Ampho , CrO 2 Ampho , CrO 3 Acidic Mn – MnO basic, Mn 2 O 3 Basic, Mn 3 O 4 Ampho , MnO 2 Ampho , Mn 2 O 7 Acidic Fe – FeO Basic, Fe 2 O 3 Amph , Fe 3 O 4 Basic Co – CoO Basic Ni – NiO Basic Cu – Cu 2 O Basic, CuO Ampho Zn – ZnO Ampho

48. In general lower oxidation state metal – BASIC Higher oxidation state metal – ACIDIC Intermediate oxidation state - AMPHOTERIC Example MnO (+2)basic, Mn 2 O 3 (+3)Basic, Mn 3 O 4 (+ 8/3) Ampho , MnO 2 (+4) Ampho , Mn 2 O 7 (+7)Acidic

49. The highest oxidation number in the oxides coincides with the group number and is attained in Sc 2 O 3 to Mn 2 O 7 . Beyond Group 7, no higher oxides of Fe above Fe 2 O 3 , are known, although ferrates (VI) (FeO4) 2– , are formed in alkaline media but they readily decompose to Fe 2 O 3 and O 2 . Besides the oxides, oxocations stabilise V(v) as VO 2 + , V(IV) as VO 2+ and Ti(IV) as TiO 2+.

50. As the oxidation number of a metal increases, ionic character decreases. In the case of Mn , Mn 2 O 7 is a covalent green oil. Even CrO 3 and V 2 O 5 have low melting points. In these higher oxides, the acidic character is predominant.

51. Potassium dichromate K 2 Cr 2 O 7 STEP 1 Dichromates are generally prepared from chromate which in turn are obtained by the fusion of chromite ore (FeCr 2 O 4 ) with sodium or potassium carbonate in free access of air. The reaction with sodium carbonate occurs as follows: 4 FeCr 2 O 4 + 8 Na 2 CO 3 + 7 O 2 → 8 Na 2 CrO 4 + 2Fe 2 O 3 + 8 CO 2

52. STEP 2 The yellow solution of sodium chromate is filtered and acidified with sulphuric acid to give a solution from which orange sodium dichromate, Na 2 Cr 2 O 7 . 2H 2 O can be crystallised. 2Na 2 CrO 4 + H 2 SO 4 → Na 2 Cr 2 O 7 + Na 2 SO 4 + H 2 O STEP 3 Conversion of Sodium dichromate in to Potassium dichromate Na 2 Cr 2 O 7 + 2 KCl → K 2 Cr 2 O 7 + 2 NaCl

53. The oxidation state of chromiumin chromate and dichromate is the same. 2 CrO 4 2– + 2H + → Cr 2 O 7 2– + H 2 O Cr 2 O 7 2– + 2 OH - → 2 CrO 4 2– + H 2 O The chromate ion is tetrahedral whereas the dichromate ion consists of two tetrahedral sharing one corner with Cr–O–Cr bond angle of 126°.

54. Sodium and potassium dichromates are strong oxidising agents Potassium dichromate is used as a primary standard in volumetric analysis. In acidic solution, its oxidising action can be represented as follows: Cr 2 O 7 2– + 14H + + 6e – → 2Cr 3+ + 7H 2 O ( EV = 1.33V)

55. 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: 6 I – → 3I 2 + 6 e – ; 3 H 2 S → 6H + + 3S + 6e – 3 Sn 2+ → 3Sn 4+ + 6 e – 6 Fe 2+ → 6Fe 3+ + 6 e – Cr 2 O 7 2– + 14 H + + 6 Fe 2+ → 2 Cr 3+ + 6 Fe 3+ + 7 H 2 O

56. Potassium permanganate KMnO 4 Potassium permanganate is prepared by fusion of MnO 2 with an alkali metal hydroxide and an oxidising agent like KNO 3 . This produces the dark green K 2 MnO 4 which disproportionates in a neutral or acidic solution to give permanganate. 2MnO 2 + 4KOH + O 2 → 2K 2 MnO 4 + 2H 2 O 3MnO 4 2– + 4H + → 2MnO 4 – + MnO 2 + 2H 2 O

57. The manganate and permanganate ions are tetrahedral; the green manganate is paramagnetic with one unpaired electron but the permanganate is diamagnetic.

58. THE INNER TRANSITION ELEMENTS ( f-BLOCK)

59. The elements in which the additional electrons enters (n-2)f orbitals are called inner transition elements . The valence shell electronic configuration of these elements can be represented as (n – 2)f 0-14 (n – 1)d 0-1 ns 2 . 4f inner transition metals are known as lanthanides because they come immediately after lanthanum and 5f inner transition metals are known as actinoids because they come immediately after actinium.

60. Electronic Configuration Element name Symbol Z Ln Ln 3+ Radius Ln 3+ / pm Lanthanum La 57 [ Xe ]6s 2 5d 1 [ Xe ]4f 116 Cerium Ce 58 [ Xe ]4f 1 6s 2 5d 1 [ Xe ]4f 1 114 Praesodymium Pr 59 [ Xe ]4f 3 6s 2 [ Xe ]4f 2 113 Neodymium Nd 60 [ Xe ]4f 4 6s 2 [ Xe ]4f 3 111 Promethium Pm 61 [ Xe ]4f 5 6s 2 [ Xe ]4f 4 109 Samarium Sm 62 [ Xe ]4f 6 6s 2 [ Xe ]4f 5 108 Europium Eu 63 [ Xe ]4f 7 6s 2 [ Xe ]4f 6 107 Gadolinium Eu 64 [ Xe ]4f 7 6s 2 5d 1 [ Xe ]4f 7 105 Terbium Tb 65 [ Xe ] 4f 9 6s 2 [ Xe ]4f 8 104 Dysprosium Dy 66 [ Xe ] 4f 10 6s 2 [ Xe ]4f 9 103 Holmium Ho 67 [ Xe ] 4f 11 6s 2 [ Xe ]4f 10 102 Erbium Er 68 [ Xe ] 4f 12 6s 2 [ Xe ]4f 11 100 Thulium Tm 69 [ Xe ] 4f 13 6s 2 [ Xe ]4f 12 99 Ytterbium Yb 70 [ Xe ] 4f 14 6s 2 [ Xe ]4f 13 99 Lutetium Lu 71 [ Xe ] 4f 14 6s 2 5d 1 [ Xe ]4f 14 98

61. Atomic and ionic sizes: The Lanthanide Contraction As the atomic number increases, each succeeding element contains one more electron in the 4f orbital and one proton in the nucleus. The 4f electrons are ineffective in screening the outer electrons from the nucleus causing imperfect shielding. As a result, there is a gradual increase in the nucleus attraction for the outer electrons. Consequently gradual decrease in size occur. This is called lanthanide contraction.

62. Consequences of L. C There is close resemblance between 4d and 5d transition series. Ionization energy of 5d transition series is higher than 3d and 4d transition series. Difficulty in separation of lanthanides

63. Ionization Enthalpies Fairly low I. E First ionization enthalpy is around 600 kJ mol -1 , the second about 1200 kJ mol -1 comparable with those of calcium. Due to low I. E, lanthanides have high electropositive character

64. Coloured ions Many of the lanthanoid ions are coloured in both solid and in solution due to f – f transition since they have partially filled f – orbitals . Absorption bands are narrow, probably because of the excitation within f level. La 3+ and Lu 3+ ions do not show any colour due to vacant and fully filled f- orbitals .

65. Magnetic properties The lanthanoid ions other then the f type (La 3+ and Ce 3+ ) and the f 14 type (Yb 2+ and Lu 3+ ) are all paramagnetic. The paramagnetism rises to the maximum in neodymium. Lanthanides have very high magnetic susceptibilities due to their large numbers of unpaired f -electrons.

66. Oxidation States Predominantly +3 oxidation state. +3 oxidation state in La, Gd , Lu are especially stable ( Empty half filled and Completely filled f – subshell respectively) Ce and Tb shows +4 oxdn state ( Ce 4+ - 4f o & Tb 4+ 4f 7 ) Occasionally +2 and +4 ions in solution or in solid compounds are also obtained. This irregularity arises mainly from the extra stability of empty, half filled or filled f subshell .

67. The most stable oxidation state of lanthanides is +3. Hence the ions in +2 oxidation state tend to change +3 state by loss of electron acting as reducing agents whereas those in +4 oxidation state tend to change to +3 oxidation state by gain of electron acting as a good oxidising agent in aqueous solution. Why Sm 2+ , Eu 2+ , and Yb 2+ ions in solutions are good reducing agents but an aqueous solution of Ce 4+ is a good oxidizing agent?

68. properties Silvery white soft metals, tarnish in air rapidly Hardness increases with increasing atomic number, samarium being steel hard. Good conductor of heat and electricity. Promethium - Radioactive

69. Chemical Properties Metal combines with hydrogen when gently heated in the gas. The carbides, Ln 3 C, Ln 2 C 3 and LnC 2 are formed when the metals are heated with carbon. They liberate hydrogen from dilute acids and burn in halogens to form halides. They form oxides and hydroxides, M 2 O 3 and M(OH) 3 , basic like alkaline earth metal oxides and hydroxides.

70. The Actinides All isotopes are radioactive, with only 232 Th, 235 U, 238 U and 244 Pu having long half-lives. Only Th and U occur naturally-both are more abundant in the earth’s crust than tin. The others must be made by nuclear processes.

71. The dominant oxidation state of actinides is +3. Actinides also exhibit an oxidation state of +4. Some actinides such as uranium, neptunium and plutonium also exhibit an oxidation state of +6. The actinides show actinide contraction (like lanthanide contraction) due to poor shielding of the nuclear charge by 5f electrons. All the actinides are radioactive. Actinides are radioactive in nature.

72. Actinoide Contraction The size of atoms / M 3+ ions decreases regularly along actinoid seris . The steady decrease in ionic/ atomic radii with increase in atomic number is called Actinoide Contraction. The contraction is greater from element to element in this series – due to poor shielding effect by 5 f electron.

73. Electronic configuration Element name Symbol Z Ln Ln 3+ Radius Ln 3 + / pm Actinium Ac 89 [ Rn ] 6d 1 7s 2 [ Rn ]4f 111 Thorium Th 90 [ Rn ]5d 2 7s 2 [ Rn ]4f 1 Protactinium Pa 91 [ Rn ]5f 2 6d 1 7s 2 [ Rn ]4f 2 Uranium U 92 [ Rn ]5f 3 6d 1 7s 2 [ Rn ]4f 3 103 Neptunium Np 93 [ Rn ]5f 4 6d 1 7s 2 [ Rn ]4f 4 101 Plutonium Pu 94 [ Rn ]5f 6 7s 2 [ Rn ]4f 5 100 Americium Am 95 [ Rn ]5f 7 7s 2 [ Rn ]4f 6 99 Curium Cm 96 [ Rn ]5f 7 6d 1 7s 2 [ Rn ]4f 7 99 Berkelium Bk 97 [ Rn ]5f 9 7s 2 [ Rn ]4f 8 98 Californium Cf 98 [ Rn ]5f 10 7s 2 [ Rn ]4f 9 98 Einsteinium Es 99 [ Rn ]5f 11 7s 2 [ Rn ]4f 10 Fermium Fm 100 [ Rn ]5f 12 7s 2 [ Rn ]4f 11 Mendelevium Md 101 [ Rn ]5f 13 7s 2 [ Rn ]4f 12 Nobelium No 102 [ Rn ]5f 14 7s 2 [ Rn ]4f 13 Lawrencium Lr 103 [ Rn ]5f 14 6d 1 7s 2 [ Rn ]4f 14

74. Magnetic properties Paramagnetic behaviour Magnetic properties are more complex than those of lanthanoids . M.P and B.P High M.P and B.P Do not follow regular gradation of M.P or B.P with increase in atomic number

75. IONISATION ENTHALPY Low I.E. so electropositiity is High COLOUR Generally coloured Colour depends up on the number of 5 f electrons The ions containing 5 f o and 5 f 7 are colouress Eg – U 3+ ( 5 f 3 ) – Red NP 3+ ( 5 f 4 ) – Bluish

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