Transition Elements. Electronic Configuration- (n-1)d 1-10 ns 1-2 Physical Properties-

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1 Transition Elements Electronic Configuration- (n-1)d 1-10 ns 1-2 Physical Properties- 1. Shows typical metallic character like high tensile strength, ductility, malleability, high thermal and electrical conductivity and metallic lusture (due to greater effective nuclear charge and large no. of valence electrons). 2. Very hard (except Zn, Cd, Hg) and have low volatility. 3. Melting point first increases and then decreases (presence of metallic bond and half filled d- orbitals. Due to the presence of half-filled d-orbitals they have covalent bond [strong inter-atomic bonding]).

2 4. Possess high Enthalpies of Atomization. EOA exhibit maxima at about the middle of each series due to the presence of stronger inter-atomic interaction with increase in half-filled d-orbital. Note: Definition: Enthalpy of atomization is the amount of enthalpy change when a compound's bonds are broken and the component atoms are reduced to individual atoms. Enthalpy of atomization is denoted by the symbol ΔH a.

3 Atomic and Ionic Sizes 1. First size decreases and then increases. Reason- The decrease in size in the beginning is attributed to the increase in nuclear charge. However, the increased nuclear charge is cancelled by the increased shielding effect of electrons in d- orbitals of the penultimate shell. When the increased nuclear charge and increased shielding effect balance each other, the atomic radii becomes almost constant. Increase in atomic radii towards the end may be attributed to the electron-electron repulsions. In fact, pairing occurs after d 5 configuration. The repulsive interaction between the paired electrons in d- orbital become very dominent towards the end of the period and causes the expansion of electron cloud and thus, increase in atomic size. 2. The filling of 4f before 5d orbital results in a regular decrease in atomic radii called LANTHANOID CONTRACTION.

4 3. Atomic radii increases down the group. But the atomic radii of third transition series is nearly the same as metals of second transition series. Explanation: Size increases due to increase in the no. of outermost shells. Metals of 2 nd and 3 rd transition series have similar atomic sizes due to LANTHANOID CONTRACTION. This is associated with the intervention of 4f- orbitals which are filled before 5d series of elements starts. The filling of 4f- orbitals before 5d orbitals results in regular decrease in atomic radii which compensates the expected increase in atomic size with increasing atomic no.

5 Ionization Enthalpy 1. I.P. increases from left to right. Reason-The increase in I.P. is primarily due to E.N.C. As the transition elements involve the gradual filling of (n-1) d-orbital, the effect of increase in nuclear charge is partly cancelled by increase in screening effect. Consequently, the increase in I.P. along the period is very small. 2. To form M 2+ ions from the gaseous atoms the sum of 1 st and 2 nd I.P. is required in addition to the enthalpy of atomisation of each element

6 Significance of I.E. The sum of First two and First four I.E. of Ni and Pt are as under: Elements Sum of First two I.E. Sum of First four I.E.(kJ/mol) Ni 2.49x x10 3 Pt 2.66x x10 3

7 Since the sum of First two I.E. of Ni is less so Ni(II) is thermodynamically stable than Pt(II) compounds. Similarly Pt(IV) compounds are stable than Ni(IV) compounds. In addition to I.E., the other factors that determine the stability of a particular state are the Enthalpy of Sublimation of the metal and the Lattice Energy or Solvation Energy. M(s) M + (aq)h t = M(s) M(g) H sub M + (g)i.e. + e - M + (aq)h hyd

8 Smaller the value of Total Enthalpy Change H t, for a particular oxidation state of metal, greater will be the stability of that oxidation state in aq. Solution.

9 Metallic character and Enthalpy of Atomization All transition elements are metals. Explanation: Metallic character is due to their relatively low ionisation enthalpies and no. of vacant orbitals in the outermost shell. The hardness of the metals suggests the presence of strong bonding due to overlap of unpaired electrons between different metal atoms. Therefore, these elements exhibit high enthalpies of atomisation. The maxima at about the middle of series indicates that 1 unpaired electron per d- orbital is favourable for strong inter- atomic interactions.

10 Greater the no. of unpaired d- electrons, greater is the no. of bonds and therefore, greater is the strength of these bonds. Thus, as we move from left to right, the no. of unpaired electrons increases from 1 to 6 and then decrease to 0 in case of group 12. Cr, Mo and W have max. no. of unpaired electrons and therefore, these are very hard metals and have max. enthalpies of atomization. Zn, Cd and Hg have no unpaired electrons so they are not very hard.

11 Density High density. Explanation: The atomic volumes of the transition elements are low because the electrons are added in (n-1)d subshell. Therefore, the increased nuclear charge is partly screened by the d electrons and the outer electrons are strongly attracted by the nucleus. Moreover the added electrons occupy inner electrons. Consequently, the densities of transition metals are high. Thus decrease in at. Radii coupled with increase in at. Mass results in increase in density.

12 Melting and Boiling points Very high melting and boiling points. Explanation: high values are due to strong metallic bonds between the atoms of the elements. These metals have high enthalpies of atomization. The metallic bond is formed due to the interaction of electrons in the outermost orbitals. The strength of bonding is roughly related to the no. of unpaired electrons. Thus metallic strength increases upto the middle due to increase in the no. of unpaired electrons and then decreases with the decreasing availability of unpaired electrons.

13 Oxidation States Sc +3 Ti V Cr Mn Fe Co Ni Cu Zn +2

14 The variable oxidation states are due to the participation of ns and (n-1)d electrons in bonding. The lower O.S. is exhibited when ns electrons participate in bonding and higher O.S. is exhibited when both ns and (n-1)d electrons participate in bonding. O.S. of these elements differ by unity. Elements with lower O.S. form ionic bonds but with higher O.S. form covalent bonds.

15 Trends in M 2+ /M Standard Electrode Potential Stability of O.S. in solution is explained on the basis of reduction potential. Lower the reduction potential, higher the -ve value, greater the stability of O.S. Due to high +ve value of reduction potential of Cu, it is not able to displace H 2 gas from dilute acids, i.e., the sum of enthalpies of sublimation and ionization is not balanced by its high hydration enthalpies.

16 There is trend of less ve reduction potential across the series due to increase in sum of 1 st and 2 nd I.P. There is no regular trend due to irregular variation in I.P. and sublimation enthalpies. The values for Mn, Ni and Zn are more ve than expected trend.

17 Trend in M 3+ /M 2+ Standard Electrode Potential Sc has lower value as Sc 3+ has noble gas configuration. Highest value of Zn is due to stable d 10 configuration, for Zn 2+. High value of Mn is due to stable Mn 2+ 3d 5 configuration while Fe has lower value as Fe 3+ 3d 5 is very stable.

18 Trends in stability of Higher O.S. Transition elements show higher O.S. in fluorides and oxides. e.g. MnO Reason is its high electro negativity. They easily pull electrons from transition elements and transition elements shows higher oxidation state.

19 Cu in +2 O.S. form all halides except iodides because Cu 2+ oxidizes I - to I 2. Cu I - Cu 2 I 2 +I 2 Cu in +1 O.S. disproportionate to give stable Cu 2+ and Cu. Cu 1+ Cu 2+ + Cu

20 Cr 3+ (3d 3 4s 0 ) is more stable due to higher hydration energy and stable t 2 g 3 configuration.cr 2+ (3d 4 4s 0 ) 2Cr H + 2Cr 3+ + H 2

21 Disproportionation of an Oxidation State Reactions in which a single substance undergoes change to produce products, one of which is in higher oxidation state and the other in lower oxidation state are called disproportionation reactions. It takes place when a particular oxidation state of the element becomes less stable than the other oxidation states, one lower and one higher. A part is oxidized while the other is reduced. e.g. Mn 3+ undergoes disproportionation in aq. Medium.

22 2Mn 3+ (aq)(iii) + 2H 2 O MnO 2 (s)(iv) + Mn 2+ (II) + 4H +

23 Magnetic Properties Mainly two types of magnetic behaviour diamagnetism and paramagnetism. Diamagnetic substances are repelled by the applied field. Paramagnetic substances are attracted by the applied field. Many of the transition metal ions are paramagnetic.

24 Paramagnetism arises due to the presence of unpaired electron. Each unpaired electron having a magetic moment associated with its spin angular momentum and orbital angular momentum. Magnetic moment is determined by the number of unpaired electrons and is calculated by using spin-only formula.

25 Formation of Coloured ions Compounds are coloured both in solid and aqueous state. The colour of these complexes is due to absorption of some radiations from visible light, which is used in promoting an electron from one of the d-orbital to another. Reason- The d-orbitals in the transition elements do not have same energy in their complexes. Under the influence of the ligands attached, the d-orbitals split into two sets of orbitals having slightly different energies.

26 In the transition elements which have partly filled d-orbitals, the transition of electron can take place from one of the lower d-orbital to some higher d-orbital within the same subshell. The energy required for this transition falls in the visible region. So when white light falls on these complexes they absorb a particular color from the radiation for the promotion of electron and the remaining colors are emitted. The color is due to emitted radiation.

27 Formation of Complex Compounds Transition elements form many coordination complexes. Reason- 1. Small size and high charge density of the ions of transition metals 2. Presence of vacant orbitals of appropriate energy which can accept lone pair of electrons donated by ligands.

28 The no. of electron pair accepted by the transition metal ion is called Coordination no.

29 Catalytic Property Many transition metals and their compounds are known to act as Catalysts. E.g. Iron acts as catalyst for the manufacture of ammonia by Haber`s Process. Reasons- 1. Because of their variable valencies transition metals sometimes form unstable intermediate compounds and provide a new path with lower activation energy for the reaction.

30 e.g. V 2 O 5 catalyzes the oxidation of SO 2 to SO 3. 2SO 2 + O 2 2SO 3 (V 2 O 5 ) The catalytic action can be understood as: V 2 O 5 +SO 2 SO 3 + V 2 O 4 V 2 O 4 + ½ O 2 V 2 O 5 2. In some cases transition metals provide a suitable surface of the reaction to take place. The reactants are adsorbed on the surface of the catalyst where reaction occurs.

31 Formation of Interstitial Compounds Interstitial compounds are formed when small atoms like H, C, N are trapped inside the crystal lattices of metals. Non-stoichiometric and neither typically ionic nor covalent. Referred to as Interstitial compounds. Physical and chemical characteristics are as follows:

32 1. Have high melting points, higher than those of pure metals. 2. Very hard. 3. Retain metallic conductivity. 4. Chemically inert.

33 Reason- Alloy Formation The d-block elements have almost similar atomic sizes. These elements can mutually substitute their positions in their crystal lattices. E.g. Brass (Cu-Zn), Bronze (Cu-Sn) Alloys are relatively hard and have high melting points.

34 Oxides and Oxoanions of Metals Metals of first transition series form oxides with oxygen at high temperature. The oxides are formed in the oxidation states +1 to +7 Highest oxidation states in the oxides of any transition metal is equal to its group no., e.g. 7 in Mn 2 O 7. Some metals in higher oxidation state stabilize by forming oxocations, e.g., V v as VO 2+, V iv as VO 2+.

35 All the metals except scandium form the oxides with the formula MO which are ionic in nature. As the oxidation no. of metal increases, ionic character decreases. e.g. MnO Mn 3 O 4 Mn 2 O 3 MnO 2 Mn 2 O 7 IONIC CHARACTER DECREASES

36 The oxides in lower oxidation states of metals are basic and in higher oxidation states are acidic whereas intermediate oxidation states are amphoteric. MnO Mn 3 O 4 Mn 2 O 3 MnO 2 Mn 2 O 7 ACIDIC CHARACTER INCREASES

37 Potassium Dichromate PREPARATION: Prepared from the ore called chromite or ferrochrome or chrome iron, FeO.Cr 2 O 3. Various steps involved are: 1. Preparation of sodium chromate: 4FeCr 2 O 4 +8Na 2 CO 3 +7O 2 8Na 2 CrO 4 +2Fe 2 O 3 +8CO 2 The yellow compounds of sodium chromate is acidified with sulphuric acid to give orange solution of crystallized sodium dichromate solution: 2Na 2 CrO 4 +2H + Na 2 Cr 2 O 7 +2 Na + +H 2 O Sodium dichromate is more soluble than potassium dichromate. The reaction of compounds of sodium dichromate with potassium chloride. Na 2 Cr 2 O 7 +2KCl K 2 Cr 2 O 7 +2 NaCl Some orange crystals of potassium dichromate crystallize out. The chromate and dichromate are intercovertible in aqueous Solution. 2CrO H + Cr 2 O 2-7 +H 2 O Cr 2 O OH - 2 CrO 2-4 +H 2 O

38 Oxidizing character of Dichromates Sodium and potassium dichromates are strong oxidizing agents. In acidic solution the oxidising action can be represented as: Cr 2 O H + + 6e - 2Cr H 2 O Acidified potassium dichromate will oxidise iodides to iodine, sulphides to sulphur, tin(ll) to tin(lv) and iron(ll) to iron(lll)

39 6I - 3I 2 + 6e - 3Sn 2+ 3Sn e - 3H 2 S 6H + + 3S + 6e - 6Fe 2+ 6Fe e - The full ionic equation: Cr 2 O H + +6 Fe 2+ 2Cr Fe H 2 O

40 Potassium Permanganate Potassium permanganate is prepared by mixture of MnO 2 with an alkali metal hydroxide and an oxidizing agent like KNO 3. This produces the dark green K 2 MnO 4 which disproportionate in a neutral or acidic solution give permanganate. 2MnO 2 + 4KOH + O 2 2K 2 MnO 4 + 2H 2 O 3MnO H + 2MnO MnO 2 + 2H 2 O

41 Manufacture: By alkaline oxidative fusion of MnO 2 followed by electrolytic oxidation of manganate(vl). MnO 2 MnO 4 2- MnO 4 2- MnO 4 -

42 Lab preparation: Manganese (ll) ion salt is oxidized by peroxodisulphate to permanganate. 2Mn S 2 O H 2 O 2MnO SO H + 2KMnO 4 K 2 MnO 4 + MnO 2 + O 2

43 Oxidizing Action Acidified permanganate solution oxidizes oxalates to CO 2, Fe(ll) to Fe(lll), nitrites to nitrates and iodides to free iodine. 5(C 2 O 4 ) 2-10 CO 2 +10e - 5Fe 2+ 5Fe e - 5NO H 2 O 5NO H + +10e - 10I - 5I e -

44 Reduction MnO 4 - +e - MnO 4 2- MnO H + +3e - MnO H + + 3e - MnO 2 +2H 2 O Mn H 2 O

45 Some important Oxidizing reactions of KMnO 4 are given below: In acid solution: 1. Iodine is liberated from potassium iodide: 10 I - + 2MnO H + 2Mn H 2 O + 5I 2 2. Fe 2+ ion (green) is converted to Fe 3+ (yellow): 5Fe 2+ + MnO H + Mn H 2 O + 5Fe 3+ In neutral or faintly alkaline solution: 1. A notable reaction is the oxidation of iodide to iodate: 2MnO H 2 O + I - 2MnO 2 + 2OH - + IO 3-2. Thiosulphate is oxidized almost quantitatively to sulphate: 8MnO S 2 O H 2 O 8MnO 2 + 6SO OH -

46 Lanthanoids Atomic and ionic radii- decreases regularly due to Lanthanoid Contraction. Reason: As the atomic no. increases in lanthanoid series, for every proton in the nucleus the extra electron goes to fill 4-f orbitals. The 4-f electrons constitute inner shells and are rather ineffective in screening the nuclear charge. Thus, there is gradual increase in the effective nuclear charge experienced by the outer electrons. Consequently, the attraction of the nucleus for the electrons in the outermost shell increases as the atomic no. of lanthanoids increases and the electron cloud shrinks. Thus size decreases gradually.

47 Oxidation State Exhibit principal oxidation state of O.S. in lanthanum, gadolinium and lutetium are very stable due to empty, halffilled and fully- filled 4f sub- shell. Cerium and terbium exhibit O.S. of +4. Ce 4+ has 4f 0 and Tb 4+ has 4f 7 configuration. Some shows +4 and +2 O.S., yet they have the tendency to attain +3 O.S. because +3 O.S. is most stable state for all lanthanides.

48 E.g. Ce 4+ is a good oxidizing agent while Sm 2+ is a good reducing agent. Ce 4+ + Fe 2+ Ce 3+ + Fe 3+ 2Sm H 2 O 2Sm 3+ +2OH - + H 2

49 General Characteristics Density- Increases with increase in atomic no. Melting and Boiling Points- Fairly high with no regular trend. Ionization Energies- Fairly low. Electropositive Character- High due to low ionization enthalpies.

50 Coloured ions- Coloured in solids as well as solutions. The colour is attributed to f-f transition since they have partly filled f-orbital. Magnetic Behaviour- Paramagnetic due to unpaired electrons.

51 Chemical Reactivity Heat with carbon at 2773 K- Form carbides Burn in oxygen- Form oxides (Ln 2 O 3 ) Heat with sulphur- Form sulphides (Ln 2 S 3 ) Heat with nitrogen- Form nitrides (LnN) With water- Form hydroxide and hydrogen (Ln(OH) 3 + H 2 )

52 With halogens- Form halides (LnX 3 ) With acids- Liberate H 2 gas.

53 Actinoids Ionic Sizes- Gradual decrease due to actinoid contraction. Oxidation States- 1. Shows +3 O.S. 2. Elements in first half series exhibit higher O.S. 3. May be +4, +5, +6 and +7.

54 General Characteristics Physical Appearance- Silvery white. Density- High densities except Thorium and Americium. Colour- Usually coloured. Elements having 2 to 6 electrons are coloured. I.E.- Low Electropositive Character- Highly electropositive.

55 Melting and Boiling Points- High and no regular trend. Magnetic Properties- Paramagnetic due to unpaired electrons. Radioactive nature.