General Properties of the Transition Elements (d-Block)

Chemistry THE D-AND F-BLOCK ELEMENTS GENERAL PROPERTIES OF THE TRANSITION ELEMENTS (d-Block)  CHARACTERISTIC PROPERTIES OF TRANSITION ELEMENTS: (a) Variable oxidation state (b) Coloured ions (c)Paramagnetic properties (d) Catalytic properties (e) Formation of alloys (f)Formation of interstitial compounds (g) Formation of complexes. Variable Valency or Variable Oxidation States (a) They exhibit variable valency due to involvement of (ns) and (n-1) d-electrons. Due to less energy difference between these electrons. (b) The oxidation states of all transition elements of '3d' series are as follows – Element Conf. Outer electronic Oxidation states configuration Sc 3d14s2 +3 3d 4s Ti 3d24s2 +2 +3 +4 V 3d34s2 +2 +3 +4 +5 Cr 3d54s1 +1 +2 +3 +4 +5 +6 Mn 3d54s2 +2 +3 +4 +5 +6 +7 Fe 3d64s2 +2 +3 +4 Co 3d74s2 +2 +3 +4 Ni 3d84s2 +2 +3 +4 Cu 3d104s1 +1 +6 +2 1 Chemistry Zn 3d104s2 +2 (c) Highest oxidation state of transition elements can be calculated by n + 2 where (n = number of unpaired electrons) It is not applied for Cr and Cu. (d) The transition metal ions having stable configuration like d 0 d5 or d10 are more stable. Ex. Sc+3, Ti+4, V+5, Fe+3 , Mn+2, Zn+2 etc. (e) In aqueous medium Cr+3 is stable. (f) Co+3 and Ni+2 is stable in complexes. (g) In aqueous medium due to disproportionation Cu+1 is less stable than Cu+2 while its configuration is 3d10 (h) Most common oxidation state among the transition elements is +2. (i) Highest oxidation state shown by transition elements of '4d' and '5d' series is +8 by Ru (44) and Os (76). (j) The common oxidation state shown by elements of IIIB i.e., Sc, Y, La and Ac is +3 as their divalent compounds are highly unstable. (k) In lower oxidation state transition elements form ionic compounds and in higher oxidation state their compounds are covalent. (l) They also shows zero oxidation state in their carbonyl compounds like Ni(CO) 4. (m) Usually transition metal ions in their lower oxidation state act as reducing agents and in higher oxidation state they are oxidising agents. Ex. Sc+2, Ti+2, V+2, Fe+2, Co+2 etc are reducing agents Cr+6, Mn+7, Mn+6, Mn+5, Mn+4 etc are oxidising agents. Standard Electrode Potentials The ionization enthalpy magnitude indicates the energy needed to remove electrons and create a specific oxidation state of the metal in a compound. Consequently, ionization enthalpy values provide insights into the thermodynamic stability of 2 Chemistry transition metal compounds in various oxidation states. A lower ionization enthalpy for the metal implies greater stability for its compound For example, the first four ionisation enthalpies of nickel and platinum are given below: Ionisation enthalpies Ni Pt IE1 + IE2 2.49×103 Kj mol -1 2.66×103 Kj mol -1 IE3 + IE4 8.8 ×103 Kj mol -1 6.70×103 Kj mol -1 11.29 ×103 Kj mol -1 9.36×103 Kj mol -1 Total It is clear from the above table that the sum of first two ionization enthalpies is less for nickel than for platinum. Ni → Ni2+ + 2e– I.E. = 2.49 × 103 kJ mol–1 Pt → Pt2+ + 2e– I.E. = 2.66 × 103 kJ mol–1 Consequently, the ionization of nickel to Ni 2+ is more energetically advantageous in comparison to platinum. This establishes that nickel (II) compounds exhibit greater thermodynamic stability than platinum (II) compounds. Conversely, the total of the first four ionization enthalpies is lower for platinum than for nickel, as demonstrated by: Ni → Ni4+ + 2e– I.E. = 11.29 × 103 kJ mol–1 Pt → Pt4+ + 2e– I.E. = 9.36 × 103 kJ mol–1 Hence, platinum (IV) compounds exhibit comparatively higher stability than nickel (IV) compounds. Consequently, K2PtCl6, which features Pt (IV), is a recognized compound, while the corresponding nickel compound remains unknown. However, in solutions, the stability of these compounds is contingent upon electrode potentials. Electrode potentials: In addition to ionization enthalpy, factors like enthalpy of sublimation, hydration enthalpy, and ionization enthalpy collectively contribute to the stability of a specific oxidation state in a solution. This can be elucidated based on their electrode potential values. The oxidation potential of a metal encompasses the following process: M(s) → M+(aq) + e– 3 Chemistry This process actually takes place in the following three steps as given in following flowchart: (i) In the first step, the atoms get isolated from one another and become independent in the gaseous state. This converts solid metal to the gaseous state. The energy needed for this step is known as enthalpy of sublimation. M(s) → M+(g) Enthalpy of sublimation, subH (ii) In the second step, the outer electron is removed from the isolated atom. The energy required for this change is ionisation enthalpy. M(s) → M+(g) + e– Ionisation enthalpy, IE (iii) In the third step the gaseous ion gets hydrated. In this process, energy known as hydration enthalpy, is liberated. M+(g) + nH2O → M+(aq) Enthalpy of hydration, hydH The oxidation potential which gives the tendency of the overall change to occur, depends upon the net effect of these three steps. The overall energy change is H = subH+ IE + hydH Hence, ΔH represents the enthalpy change necessary to transform the solid metal, M, into the monovalent ion in an aqueous medium, M+ (aq). A similar cycle can be constructed for the formation of an anion in a solution, with the ionization enthalpy potentially replaced by electron gain enthalpy as the gaseous atom transforms into a gaseous anion. ΔH aids in predicting the stability of a specific oxidation state. The smaller the values of total energy change for a given oxidation state in an aqueous solution, the greater the stability of that oxidation state. Electrode potentials serve as a measure of total energy change. Qualitatively, the stability of transition metal ions 4 Chemistry in various oxidation states can be assessed based on electrode potential data. The lower the electrode potential, or the more negative the standard reduction potential of the electrode, the more stable the oxidation state of the transition metal in the aqueous solution. The electrode potentials of different metals can also be measured by forming the cell with standard hydrogen electrode. For the measurement of electrode potential of M2+ | 1M, the e.m.f. of the cell in which the following reaction occurs is measured: 2H+(aq) + M(s) ⇌ M2+ (aq) + H2(g) Knowing the potential of 2H+(aq) | H2(g), it is possible to determine the potential of M2+ (aq) | M. For the first transition series, the E values of M2+(aq) | M are given below: The observed values of E and those calculated using the data are compared in the following figure. Thermochemical data (kJ mol–1) for the first row Transition Elements and the Standard Electrode potentials for the Reduction of MII to M 5 Chemistry Element (M) Ti V Cr Mn Fe Co Ni Cu Zn △a Hq (M) 469 515 398 279 418 427 431 339 130 △f 𝐇𝟏𝛉 661 648 653 716 762 757 736 745 908 △f 𝐇𝟐𝛉 1310 1370 1590 1510 1560 1640 1750 1960 1730 △hyd 𝐇 𝛉 (M2+) -1866 -1895 -1925 -1862 -1998 -2079 -2121 -2121 -2059 𝐄𝛉 / 𝐕 -1.63 -1.18 -0.90 -1.18 -0.44 -0.28 -0.25 0.34 -0.76 The results lead to the following conclusions: (i) There is no regular trend in these values. This is attributed to the irregular variation of ionisation enthalpies (IE1 + IE2) and the sublimation energies in the period. (ii) It may be noted that the electrode potentials of transition metals are low in comparison to elements of group 2 (e.g., Ca = –2.87 V). Compared to group 2 elements, the transition elements have fairly large ionisation enthalpies and very large enthalpies of atomization. These reduce their electrode potentials though their hydration enthalpies are large. (iii) Zinc has low enthalpy of atomization and fairly large hydration energy. But it has also low electrode potential (–0.76 V) because of its very high ionisation enthalpy (IE 1 + IE2). (iv) It is clear from above table and figure that copper has positive reduction potential, E (0.34 V) and this shows that copper is least reactive metal out of the first transition series. This unique behavior (+ve) E value of copper) also accounts for its inability to liberate H2 from acids. It has been observed that only oxidizing acids (such as nitric acid and hot concentrated sulphuric acid) react with copper in which the acids are reduced. The high energy required to convert Cu(s) to Cu2+ (aq) is not balanced by its hydration enthalpy. (v) In general, the value becomes, less negative across the series. This is related to the general increase in the sum of first and second ionisation enthalpies. It is interesting to note that the values of Eº of Mn, Ni and Zn are more negative than expected from 6 Chemistry the general trend. The relatively more negative values of E  for Mn and Zn are due to stability of half-filled d-sub-shell in Mn2+ (3d5) and the completely filled (3d10) configuration in Zn2+. The exceptionally high E value of Ni from regular trend is related to the highest negative enthalpy of hydration of Ni 2+ ion. Trends in the M3+ | M2+ Standard Electrode Potentials Except copper and zinc, all other elements of first transition series show +3 oxidation states also to form M3+ ions in aqueous solutions. The standard reduction potentials for M3+ | M2+ redox couple are given below: 3+ E (aq) | M2+ (aq) (in Volt) Ti –0.37 V –0.26 Cr –0.41 Mn +1.57 Fe +0.77 Co +1.97 These values reveal the following facts: (i) The low value of scandium reflects the stability of Sc 3+ which has a noble gas configuration. (ii) The comparatively high value for Mn shows that Mn 2+ (d5 configuration) is particularly stable. On the other hand, comparatively low value for Fe shows the extra stability of Fe3+ (d5 configuration). (iii) The comparatively low value of V is related to the stability of V 2+ (due to half-filled t2g3 energy level of 3d orbitals in octahedral crystal field spitting). (iv) The E value for Mn3+ /Mn2+ couple much more positive than for Cr3+/Cr2+ or Fe3+ /Fe2+. This is because of the much larger IIIrd ionisation energy of Mn (removal of electron from d5 configuration). Trends in Stability of Higher Oxidation States Data on standard electrode potentials offer valuable insights into the stability of various oxidation states exhibited by an element. The highest oxidation states are typically observed in halides and oxides. 7 Chemistry 1. In metal halides. The transition elements undergo reactions with halogens at elevated temperatures, resulting in the formation of transition metal halides. These reactions exhibit substantial heat of reaction. However, once the reaction initiates, the generated heat is adequate to sustain the reaction. The halogens engage in the reaction in the following descending order F2 > Cl2 > Br2 > I2 Halides of first transition series Oxidation Number +6 +5 +4 +3 +2 Sc Ti V Cr Min Fe Co Ni Cu Zn CrF6 VF5 CrF5 TiX4 VX4a CrF4 MnF4 ScX3 TiX3 VX3 CrF3 MnF3 FeXa3 CoF3 TiX2 VX2c CrF2 MnX2 FeX2 +1 CoF2 NiX2 CuX2b ZnX2 CuXc where X = F, Cl, Br, I, Xa = F, Cl, Br, Xb = F, Cl, Xc = Cl, Br, I Among the transition groups 3-12, variations in the stability of different oxidation states exist. Typically, the second and third transition series elements demonstrate higher coordination numbers, making their higher oxidation states more stable compared to the corresponding elements in the first transition series. The following trends are observed from table regarding transition metal halides: (i) In general, the elements of first transition series tend to exist in low oxidation states. Chromium to zinc form stable difluorides and the other chlorides are also known. (ii) Since fluorine is the most electronegative element, the transition metals show highest oxidation states with fluorine. The highest oxidation states are found in TiX 4 (tetrahalides, X = F, Cl, Br and I), VF5 and CrF6. (iii) The +7-oxidation state for Mn is not shown by simple halides. However, MnO 3F is known in which the oxidation state of Mn is +7. 8 Chemistry (iv) After Mn, the tendency to show higher oxidation states with halogens are uncommon. Iron and cobalt form trihalides FeX3 (X = F, Cl or Br) and CoF3. (v) The tendency of fluorine to stabilize the highest oxidation state is due to either higher lattice enthalpy as in case of CoF3 or higher bond enthalpy due to higher covalent bonds e.g., VF5 and CrF6. (vi) V(V) is shown by VF5 only. However, the other halides undergo hydrolysis to form oxyhalides, VOX3. (vii) Fluorides are relatively unstable in their low oxidation states. For example, vanadium form only VX2 (X = Cl, Br or I) and copper can form CuX (X = Cl, I). All copper (II) halides are known except the iodide. This is because, Cu 2+ oxidizes I– to I2. 2Cu2+ + 4I– → Cu2I2(s) + I2 It has been observed that many copper (I) compounds are unstable in aqueous solution and they undergo disproportionation to Cu(II) and Cu(0) as : 2Cu+ → Cu2+ + Cu Copper in +2 oxidation state is more stable than in +1 oxidation state. This can be explained on the basis of much larger negative hydration enthalpy (hydH) of Cu2+ (aq) than Cu+, which is much more than compensates for the large energy required to remove the second electron i.e., second ionisation enthalpy of copper. 2. In metal oxides and oxo cations. Oxidation Sc Number +7 Ti V Min Fe Co Ni Cu Zn Mn2O7 +6 +5 Cr CrO3 V2 O5 +4 TiO2 V2 CrO2 MnO2 Fe2O3 O4 9 Chemistry +3 Sc2O3 TiO3 V2 Cr2O3 Mn2O3 FeO CoO NiO CuO ZnO O3 +2 TiO VO (CrO) MnO CuO +1 Mixed oxides Mn3O4 Fe3O4 Co3O4 Oxygen's capacity to stabilize the highest oxidation state is evident in its oxides. The highest oxidation states in these oxides correspond to the group number. For instance, scandium in group 3 exhibits a highest oxidation state of +3 in its oxides, Sc2O3, while manganese in group 7 shows +7 in Mn2O7. However, beyond group 7, no higher oxides of iron above Fe2O3 are known. Although higher oxidation states, such as +6, are seen in ferrates like FeO42– in alkaline medium, they readily decompose to Fe2O3 and O2. In addition to oxides, the oxidation of metals also stabilizes higher oxidation states, as seen in VV as VO2+, VIV as VO2+, and TiIV as TiO2+. Notably, oxygen's ability to stabilize these high oxidation states surpasses that of fluorine. For instance, manganese forms its highest fluoride as MnF 4, while its highest oxide is Mn2O7. This is attributed to oxygen's superior ability to form multiple bonds to metals. In covalent oxides like Mn2O7, each Mn is tetrahedrally surrounded by oxygen atoms, forming Mn–O–Mn bridges. Tetrahedral [MO 4]n– ions are also known for vanadium (V), chromium (VI), manganese (VI), and manganese (VII). Transition elements in the +2 and +3 oxidation states predominantly form ionic bonds, while with higher oxidation states, the bonds become essentially covalent. For instance, in MnO4–, all bonds are covalent. As the oxidation number of a metal increases, the ionic character of their oxides decreases. Mn 2O7, for example, is a covalent green oil. In these higher oxides, the acidic character is predominant. Thus, CrO3 yields H2CrO4 and H2Cr2O7, and Mn2O7 yields HMnO4. V2O5 is amphoteric, mainly acidic, and reacts with alkalis to give VO43– and with acids to produce VO2+. 10 Chemistry The relative stability of various oxidation states (a) The relative stabilities of various oxidation states of 3d-series element can be correlated with the extra stability of 3d°,3d 5 & 3 d10 configuration to some extent. Ex. Stability of Ti4+ (3d0) > Ti3 + (3d1 ) Mn2+ (3d5) > Mn3+ (3d4) (b) The higher oxidation state of 4d and 5d series element are generally more stable than the elements of 3d series. Ex. (i) (oxidation state of Mo is +6), (4d series) &(5d series) are more stable due to their maximum oxidation state. (ii) & (3d-series) are strong oxidizing agents. (c) Strongly reducing states probably do not form fluorides or oxides, but may well form the heavier halides Conversely, strong oxidizing state form oxides & fluoride, but not Bromide and lodide. Ex. (i) V (Vanadium) react with halogens to form VF5 VCl5, VBr3,but doesn' t form VBr5 or VI5 because in + 5 oxidation state Vanadium is strong oxidizing agent thus convert Br – & I– to Br2 & I2 respectively, So VBr3 &VI3 are formed but not VBr5 & VI5. (ii) On the other hand, VF5 is formed because V5+ ion unable to oxidize highly electronegative & small anion F – (iii) Similarly, highly electronegative and small O2 – ion formed oxides Ex. VO4 3 –, CrO42– & MnO4– etc. Different oxidation state of chloride & oxides compound +2 +3 +4 2 TiCl3 TiCl 4 VCl 2 VCl 3 VCl 4 TiCl (Ionic, basic) Less ionic (Amphoteric) +5 +6 +7 VOCl 3 Covalent and Acidic (Strong lewis acid) 11 Chemistry (d) TiO Ti2O3 VO V2O3 CrO Cr2O3 MnO Mn2O3 Ionic, basic TiO2 V2O5 CrO3 MnO2 MnO3 Less Ionic (Amphoteric) Mn2O7 Acidic, covalent Such compounds are expected to be unstable except in case where vacant d-orbitals are used for accepting lone-pair from -bonding ligand. Ex. [Ni(CO)4], [Ag(CN)2]– ,[Ag)(NH3)2]+ Colour Property (a) Most of the transition metal ions exhibit colour property. (b) This is due to d-d transition of unpaired electrons in their t 2g and eg sets of 'd' orbitals. (c) They require less amount of energy to undergo excitation of electrons. Hence, they absorb visible region of light exhibiting colour. Ex. Sc+2 : [Ar]3d1, Ti+2 : [Ar]3d2, V+2 : [Ar]3d3 (d) Transition metal ions which do not have any unpaired electrons in their 'd' orbitals like 3d0 and 3d10 configurations, do not exhibit any colour property. Ex. Sc+3 : [Ar]3d0, Cu+1 : [Ar]3d10, Ti+4 : [Ar]3d0 etc are colourless ions. (e) A transition metal ion selectively absorbs a portion of the visible light spectrum and releases the remaining colors, resulting in the emitted light's overall hue. The color of the metal ion corresponds to the emitted light's color. (f) In transition metal ion the 'd' orbitals split into lower energy set t 2g orbitals and higher energy set eg orbitals. The electrons from t2g set get excited to higher energy set e.g., set. This excitation of electrons is called as 'd-d' transition. Due to this 'd -d' transition the transition metal ions exhibit colour property. 12 Chemistry Lower energy set = t2g Higher energy set = eg Presence of ligands ↑ dx2- y2 dz2 (eg ) ↑ light d-d Transition d-orbitals (degenerate) dxy dyz dyz (t2g ) Factors affecting the colour of complex The colour of a transition metal complex depends on(a) The magnitude of energy difference between the two d-levels (0), (b) An increase in the magnitude of 0 decreases the wave length () of the light absorbed by the complexes. 0  1 λ(Wavelength of light absorb) (c) Thus, with a decrease in the  the colour of complex changes from Red to Violet. [Co(H2O)6]3+ Ex. Complex ions < [Co(NH3)6]3+ [Co(CN)6]3– NH3 < CN– Ligand field strength H2O Magnitude of 0 0(H2O) < 0(NH3) < 0(CN–) Magnitude of  (H2O) < (NH3) < (CN–) Colour of the transmitted orange Green-blue Colour of absorbed light Green-blue Orange violet Yellow greenlight (i.e., colour of the complex (d) KMnO4 (dark pink), K2Cr2O7 (orange) having d° configuration but they are coloured due to charge transfer spectrum and charge is transferred from anion to cation. Example of Some coloured metal ions: Ti+3 Purple Mn+2 Light pink Fe+3 Yellow Ni+2 Green 13 Chemistry Sc3+ Colourless Ti3+ Purple V3+ Green Cr2+ Blue Mn3+ Violet Fe2+ Green(light) Co2+ Pink Zn2+ Colourless Cr+3 Green Fe+2 Green Co+3 Pink Cu+2 Blue Ti4+ Colourless V4+ Blue V2+ Violet Cr3+ Green Mn2+ Pink Fe3+ Yellow Ni2+ Blue Ex. Explain the blue colour of CuSO4.5H2O. Sol. Cu2+ ion (3d9) absorbs red light from the visible region, for the promotion of 3d electrons, the ions reflect blue light and appear blue. Magnetic Properties (a) Typically, transition elements display magnetic properties. Para magnetism refers to the attraction of a substance to a magnetic field, primarily attributed to the existence of unpaired electrons in atoms, ions, or molecules. This property varies inversely with temperature. (b) Diamagnetic substances are slightly repelled by a magnetic field, and this characteristic is temperature-independent. (c) As is evident most of the transition metal ions have unpaired electrons in their 'd' orbitals. Hence most of the transition metal ions are paramagnetic in nature. Ex. Ti+2 [Ar]3d2, Ti+3 [Ar]3d1. V+2[Ar]3d3, Cr+3[Ar]3d3 (d) Transition metal ions having 3d0 and 3d10 configuration exhibit diamagnetic nature. (e) The total magnetic moment of a substance is the resultant of the magnetic moments of all the individual electrons. (f) The magnetic moment () created due to spinning of unpaired electrons can be calculated by using 14 Chemistry μ = n(n + 2) Where - 'n' is the number of unpaired electrons in the metal ion.  = Magnetic moment in Bohr Magnetons (B.M.) (g) The magnetic moment of diamagnetic substances will be zero. (h) Transition metal ions with a d5 configuration will exhibit maximum paramagnetic behavior due to the presence of the highest number of unpaired electrons. Catalytic Property (a) Transition elements and their compounds exhibit catalytic properties. This is due to their variable valency as well as due to the free valencies on their surface. (b) When transition elements and their compounds are in powdered state, their catalytic properties exhibited will be to a greater extent. This is due to greater surface area available in the powdered state. (c) Transition metals and their compounds exhibiting catalytic properties in various processes are Catalyst Used TiCl3 Used as the Ziegler-Natta catalyst in the production of polythene. MnO2 Used as a catalyst to decompose KClO3 to give O2 Fe Promoted iron is used in the Haber-Bosch process for making NH3 FeCl3 Used in the production of CCl4 from CS and Cl2 FeSO4 and H2O2 Used as Fenton's reagent for oxidizing alcohols to aldehydes. PdCl2 Wacker process for converting C2H4+H2O+PdCl2 to CH3CHO+2HCl+Pb. pd Used for hydrogenation (e.g., phenol to cyclohexanone). Pt/ptO Adams catalyst, used for reductions. 15 Chemistry Pt Formerly used for SO2 → SO3in the contact process for making H2SO4 Pt/Rh Formerly used in the Ostwald process for making HNO3 to NO Cu Is used in the direct process for manufacture of (CH3)2SiCl2 used to make silicones. Cu/V Oxidation of cyclohexanol/cyclohexanone mixture to adipic acid which is used to make nylone-66 CuCl Decon process of making Cl from HCl Ni Raney nickel, numerous reduction processes (e.g., manufacture of hexamethylenediamine, production of H2 from NH3 reducing anthraquinone to anthraquinol in the production of H2O2) 16