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The d-Block Elements 1 Introduction • d-block elements locate between the s-block and p-block known as transition elements occur in the fourth and subsequent periods of the Periodic Table 2 d-block elements period 4 period 5 period 6 period 7 3 Introduction Transition elements are elements that contain an incomplete d sub-shell (i.e. d1 to d9) in at least one of the oxidation states of their compounds. 3d0 3d10 4 Introduction Sc and Zn are not transition elements because They form compounds with only one oxidation state in which the d sub-shell are NOT imcomplete. Sc Sc3+ 3d0 5 Zn Zn2+ 3d10 Introduction Cu Cu+ 3d10 not transitional Cu2+ 3d9 transitional 6 The first transition series the first horizontal row of the d-block elements 7 Characteristics of transition elements (d-block metals vs s-block metals) 1. Physical properties vary slightly with atomic number across the series (cf. s-block and p-block elements) 2. Higher m.p./b.p./density/hardness than s-block elements of the same periods. 3. Variable oxidation states (cf. fixed oxidation states of s-block metals) 8 Characteristics of transition elements 4. Formation of coloured compounds/ions (cf. colourless ions of s-block elements) 5. Formation of complexes 6. Catalytic properties 9 Electronic Configurations The building up of electronic configurations of elements follow: Aufbau principle Pauli exclusion principle Hund’s rule 10 Electronic Configurations 11 • 3d and 4s sub-shells are very close to each other in energy. • Relative energy of electrons in subshells depends on the effective nuclear charge they experience. • Electrons enter 4s sub-shell first • Electrons leave 4s sub-shell first Cu Cu2+ Relative energy levels of orbitals in atom and in ion 12 Electronic Configurations • Valence electrons in the inner 3d orbitals • Examples: The electronic configuration of scandium: 1s22s22p63s23p63d14s2 The electronic configuration of zinc: 1s22s22p63s23p63d104s2 13 Electronic configurations of the first series of the d-block elements 14 Element Atomic number Electronic configuration Scandium 21 [Ar] 3d 14s2 Titanium 22 [Ar] 3d 24s2 Vanadium 23 [Ar] 3d 34s2 Chromium 24 [Ar] 3d 54s1 Manganese 25 [Ar] 3d 54s2 Iron 26 [Ar] 3d 64s2 Cobalt 27 [Ar] 3d 74s2 Nickel 28 [Ar] 3d 84s2 Copper 29 [Ar] 3d 104s1 Zinc 30 [Ar] 3d 104s2 • A half-filled or fully-filled d sub-shell has extra stability 15 d -Block Elements as Metals • d-Block elements are typical metals Physical properties of d-Block elements : good conductors of heat and electricity hard and strong malleable and ductile 16 d -Block Elements as Metals • Physical properties of d-Block elements: lustrous high melting points and boiling points • Exceptions : Mercury low melting point liquid at room temperature and pressure 17 d -Block Elements as Metals • d-block elements extremely useful as construction materials strong and unreactive 18 d -Block Elements as Metals • Iron used for construction and making machinery nowadays abundant easy to extract 19 cheap d -Block Elements as Metals • Iron corrodes easily often combined with other elements to form steel harder and more resistant to corrosion 20 d -Block Elements as Metals • Titanium Corrosion resistant, light, strong and withstand large temperature changes used to make aircraft and space shuttles expensive 21 d -Block Elements as Metals • The similar atomic radii of the transition metals facilitate the formation of substitutional alloys the atoms of one element to replace those of another element modify their solid structures and physical properties 22 d -Block Elements as Metals • Chromium confers inertness to stainless steel • Manganese confers hardness & wearing resistance to its alloys e.g. duralumin : alloy of Al with Mn/Mg/Cu 23 Atomic Radii and Ionic Radii • Two features can be observed: 1. The d-block elements have smaller atomic radii than the s-block elements 2. The atomic radii of the d-block elements do not show much variation across the series 24 Atomic Radii and Ionic Radii Variation in atomic radius of the first 36 elements 25 26 27 On moving across the Period, (i) Nuclear charge (ii) Shielding effect (repulsion between e-) (i) (ii) (i) > (ii) 28 (ii) > (i) Atomic Radii and Ionic Radii • At the beginning of the series atomic number effective nuclear charge the electron clouds are pulled closer to the nucleus atomic size 29 • In the middle of the series more electrons enter the inner 3d sub-shell The inner 3d electrons shield the outer 4s electrons effectively the effective nuclear charge experienced by 4s electrons increases very slowly only a slow decrease in atomic radius in this region 30 Atomic Radii and Ionic Radii • At the end of the series the screening and repulsive effects of the electrons in the 3d subshell become even stronger Atomic size 31 Comparison of Some Physical and Chemical Properties between the d-Block and s-Block Elements • Many of the differences in physical and chemical properties between the d-block and s-block elements explained in terms of their differences in electronic configurations and atomic radii 32 1. Density Densities (in g cm–3) of the s-block elements and the first series of the d-block elements at 20C 33 1. Density • d-block > s-block the atoms of the d-block elements 1. are generally smaller in size 2. are more closely packed (fcc/hcp vs bcc in group 1) 3. have higher relative atomic masses 34 1. Density • The densities generally increase across the first series of the d-block elements 1. general decrease in atomic radius across the series 2. general increase in atomic mass across the series 35 2. Ionization Enthalpy K Ca (sharp ) ; Ca Sc (slight ) Ionization enthalpy (kJ mol–1) Element 36 1st 2nd 3rd 4th K 418 3 070 4 600 5 860 Ca 590 1 150 4 940 6 480 Sc 632 1 240 2 390 7 110 Ti 661 1 310 2 720 4 170 V 648 1 370 2 870 4 600 Cr 653 1 590 2 990 4 770 2. Ionization Enthalpy Sc Cu (slight ) ; Cu Zn (sharp ) Ionization enthalpy (kJ mol–1) Element 37 1st 2nd 3rd 4th Cr 653 1 590 2 990 4 770 Mn 716 1 510 3 250 5 190 Fe 762 1 560 2 960 5 400 Co 757 1 640 3 230 5 100 Ni 736 1 750 3 390 5 400 Cu 745 1 960 3 550 5 690 Zn 908 1 730 3 828 5 980 2. Ionization Enthalpy • The first ionization enthalpies of the d-block elements greater than those of the s-block elements in the same period of the Periodic Table 1. The atoms of the d-block elements are smaller in size 2. greater effective nuclear charges 38 Sharp across periods 1, 2 and 3 Slight across the transition series 39 2. Ionization Enthalpy • Going across the first transition series the nuclear charge of the elements increases additional electrons are added to the ‘inner’ 3d sub-shell 40 2. Ionization Enthalpy • The screening effect of the additional 3d electrons is significant • The effective nuclear charge experienced by the 4s electrons increases very slightly across the series • For 2nd, 3rd, 4th… ionization enthalpies, slight and gradual across the series are observed. 41 Electron has to be removed from completely filled 3p subshell 3d5 3d5 3d5 42 Cr+ Fe3+ Mn2+ 3d10 d10/s2 2. Ionization Enthalpy • The first few successive ionization enthalpies for the d-block elements do not show dramatic changes 4s and 3d energy levels are close to each other 43 3. Melting Points and Hardness d-block >> s-block 1. both 4s and 3d e- are involved in the formation of metal bonds 2. d-block atoms are smaller 1541 1668 1910 1907 1246 1538 1495 1455 1084 419 44 3. Melting Points and Hardness K has an exceptionally small m.p. because it has an more open b.c.c. structure. 1541 1668 1910 1907 1246 1538 1495 1455 1084 419 45 Sc Ti V Cr 1541 1668 1910 1907 Mn Fe Co Ni Cu Zn 1246 1538 1495 1455 1084 419 Unpaired electrons are relatively more involved in the sea of electrons 46 Sc Ti V Cr 1541 1668 1910 1907 Sc Ti V Mn Fe Co Ni Cu Zn 1246 1538 1495 1455 1084 419 3d 4s 1. m.p. from Sc to V due to the of unpaired d-electrons (from d1 to d3) 47 Sc Ti V Cr 1541 1668 1910 1907 Mn Co Ni Cu Zn 1246 1538 1495 1455 1084 419 3d Fe Co Ni Fe 4s 2. m.p. from Fe to Zn due to the of unpaired d-electrons (from 4 to 0) 48 Sc Ti V Cr 1541 1668 1910 1907 Mn Fe Co Ni Cu Zn 1246 1538 1495 1455 1084 419 3. Cr has the highest no. of unpaired electrons but its m.p. is lower than V. 3d Cr 4s It is because the electrons in the half-filled d-subshell are relatively less involved in the sea of electrons. 49 Sc Ti V Cr 1541 1668 1910 1907 Mn Fe Co Ni Cu Zn 1246 1538 1495 1455 1084 419 4. Mn has an exceptionally low m.p. because it has the very open cubic structure. Why is Hg a liquid at room conditions ? All 5d and 6s electrons are paired up and the size of the atoms is much larger than that of Zn. 50 3. Melting Points and Hardness • The hardness of a metal depends on the strength of the metallic bonds • The metallic bonds of the d-block elements are stronger than those of the s-block elements much harder than the s-block elements 51 Mohs scale : - A measure of hardness Talc 0 K Diamond 10 Ca Sc Ti V Cr 0.5 1.5 3.0 4.5 6.1 9.0 5.0 52 Mn Fe Co Ni Cu Zn 4.5 -- -- 2.8 2.5 4. Reaction with Water • In general, the s-block elements react vigorously with water to form metal hydroxides and hydrogen • The d-block elements react very slowly with cold water react with steam to give metal oxides and hydrogen 53 4. Reaction with Water 2K(s) + 2H2O(l) 2KOH(aq) + H2(g) 2Na(s) + 2H2O(l) 2NaOH(aq) + H2(g) Ca(s) + 2H2O(l) Ca(OH)2(aq) + H2(g) Zn(s) + H2O(g) ZnO(s) + H2(g) 3Fe(s) + 4H2O(g) Fe3O4(s) + 4H2(g) 54 d-block compounds vs s-block compounds A Summary : Ions of d-block metals have higher charge density more polarizing 1. more covalent in nature 2. less soluble in water 3. less basic (more acidic) Basicity : Fe(OH)3 < Fe(OH)2 << NaOH Charge density : Fe3+ > Fe2+ > Na+ 55 d-block compounds vs s-block compounds A Summary : 4. less thermally stable e.g. CuCO3 << Na2CO3 5. tend to exist as hydrated salts e.g. CuSO4.5H2O, CoCl2.2H2O 6. hydrated ions undergo hydrolysis more easily e.g. [Fe(H2O)6]3+(aq) + H2O [Fe(OH)(H2O)5]2+(aq) + H3O+ acidic 56 Variable Oxidation States • One of the most striking properties variable oxidation states • The 3d and 4s electrons are in similar energy levels available for bonding 57 Variable Oxidation States • Elements of the first transition series form ions of roughly the same stability by losing different numbers of the 3d and 4s electrons 58 Oxidation states of the elements of the first transition series in their oxides and chlorides Oxidation Oxides / Chloride states Cu2O +1 +2 +3 +4 +5 +6 +7 59 Cu2Cl2 TiO VO CrO MnO FeO TiCl2 VCl2 CrCl2 MnCl2 FeCl2 CoCl2 NiCl2 Sc2O3 Ti2O3 V2O3 Cr2O3 Mn2O3 Fe2O3 ScCl3 TiCl3 VCl3 CrCl3 MnCl3 FeCl3 TiO2 VO2 MnO2 TiCl4 VCl4 CrCl4 V2O5 CrO3 Mn2O7 CoO NiO Ni2O3 • xH2O CuO ZnO CuCl2 ZnCl2 Oxidation states of the elements of the first transition series in their compounds Element Possible oxidation state Sc Ti +1 +2 +3 +4 V +1 +2 +3 +4 +5 Cr +1 +2 +3 +4 +5 +6 Mn +1 +2 +3 +4 +5 +6 Fe +1 +2 +3 +4 +5 +6 Co +1 +2 +3 +4 +5 Ni +1 +2 +3 +4 +5 Cu +1 +2 +3 Zn 60 +3 +2 +7 1. Scandium and zinc do not exhibit variable oxidation states • Scandium of the oxidation state +3 the stable electronic configuration of argon (i.e. 1s22s22p63s23p6) • Zinc of the oxidation state +2 the stable electronic configuration of [Ar] 3d10 61 2. (a) All elements of the first transition series (except Sc) can show an oxidation state of +2 (b) All elements of the first transition series (except Zn) can show an oxidation state of +3 62 3. Manganese has the highest oxidation state +7 E.g. MnO4-, Mn2O7 Mn7+ ions do not exist. 63 The +7 state of Mn does not mean that all 3d and 4s electrons are removed from Mn to give Mn7+. Instead, Mn forms covalent bonds with oxygen atoms by making use of its half O filled orbitals Mn O O O 64 - Draw the structure of Mn2O7 O O Mn Mn O O O 65 O O 3. Manganese has the highest oxidation state +7 • The highest possible oxidation state = the total no. of the 3d and 4s electrons inner electrons (3s, 3p…) are not involved in covalent bond formation 66 4. For elements after manganese, there is a reduction in the number of possible oxidation states • The 3d electrons are held more firmly the decrease in the number of unpaired electrons the increase in nuclear charge 67 5. The relative stability of various oxidation states is correlated with the stability of electronic configurations Stability : - Mn2+(aq) Major factor > [Ar] 3d5 Fe3+(aq) [Ar] 3d4 > [Ar] 3d5 Major factor 68 H Mn3+(aq) o : hydration Fe2+(aq) [Ar] 3d6 Fe3+ > Fe2+ 5. The relative stability of various oxidation states is correlated with the stability of electronic configurations Stability : - Zn2+(aq) > [Ar] 3d10 Major factor 69 H o : hydration Zn+(aq) [Ar] 3d104s1 Zn2+ > Zn+ 1. Variable Oxidation States of Vanadium and their Interconversions • The compounds of vanadium, vanadium oxidation states of +2, +3, +4 and +5 forms ions of different oxidation states show distinctive colours in aqueous solutions 70 Colours of aqueous ions of vanadium of different oxidation states Ion Oxidation state of Colour in vanadium in the ion aqueous solution V2+(aq) +2 Violet V3+(aq) +3 Green VO2+(aq) +4 Blue VO2+(aq) +5 Yellow 71 1. Variable Oxidation States of Vanadium and their Interconversions • In an acidic medium the vanadium(V) state usually occurs in the form of VO2+(aq) dioxovanadium(V) ion the vanadium(IV) state occurs in the form of VO2+(aq) oxovanadium(IV) ion 72 1. Variable Oxidation States of Vanadium and their Interconversions • In an alkaline medium the stable form of the vanadium(V) state is VO3–(aq), metavanadate(V) or VO43–(aq), orthovanadate(V), in strongly alkaline medium 73 1. Variable Oxidation States of Vanadium and their Interconversions • Compounds with vanadium in its highest oxidation state (i.e. +5) strong oxidizing agents 74 1. Variable Oxidation States of Vanadium and their Interconversions • Vanadium of its lowest oxidation state (i.e. +2) in the form of V2+(aq) strong reducing agent easily oxidized when exposed to air 75 1. Variable Oxidation States of Vanadium and their Interconversions • Interconversions of the common oxidation states of vanadium can be carried out readily in the laboratory • The most convenient starting material ammonium metavanadate(V) (NH4VO3) a white solid the oxidation state of vanadium is +5 76 1. Variable Oxidation States of Vanadium and their Interconversions 1. Interconversions of Vanadium(V) species VO2+(aq) Yellow OH H+ V2O5(s) orange OH H+ VO3(aq) yellow OH H+ VO43(aq) colourless Vanadium(V) can exist as cation as well as anion 77 1. Variable Oxidation States of Vanadium and their Interconversions 1. Interconversions of Vanadium(V) species VO2+(aq) Yellow OH H+ In acidic medium 78 V2O5(s) orange OH H+ Amphoteric VO3(aq) yellow OH H+ VO43(aq) colourless In alkaline medium 1. Variable Oxidation States of Vanadium and their Interconversions 1. Interconversions of Vanadium(V) species VO2+(aq) Yellow OH H+ In acidic medium V2O5(s) orange OH H+ Amphoteric VO3(aq) yellow OH H+ VO43(aq) colourless In alkaline medium Give the equation for the conversion : V2O5 VO2+ 79 V2O5(s) + 2H+(aq) 2VO2+(aq) + H2O(l) 1. Variable Oxidation States of Vanadium and their Interconversions 1. Interconversions of Vanadium(V) species VO2+(aq) Yellow OH H+ In acidic medium V2O5(s) orange OH H+ Amphoteric VO3(aq) yellow OH H+ VO43(aq) colourless In alkaline medium Give the equation for the conversion : V2O5 VO3 80 V2O5(s) + 2OH(aq) 2VO3(aq) + H2O(l) 1. Variable Oxidation States of Vanadium and their Interconversions 1. Interconversions of Vanadium(V) species VO2+(aq) Yellow OH H+ In acidic medium V2O5(s) orange OH H+ Amphoteric VO3(aq) yellow OH H+ VO43(aq) colourless In alkaline medium Give the equation for the conversion : VO3 VO2+ 81 VO3(aq) + 2H+(aq) VO2+(aq) + H2O(l) H H H O H O H V5+ O H H O H H O H 8H2O orthovanadate(V) ion VO43(aq) + 8H3O+ V5+ ions does not exist in water since it undergoes vigorous hydrolysis to give VO43 The reaction is favoured in highly alkaline solution 82 V VO43(aq) orthovanadate(V) ion Cr CrO42(aq) chromate(VI) ion Mn MnO4(aq) manganate(VII) ion Draw the structures of VO43, CrO42 and MnO4 O O Cr O 83 O O- - Mn O OO H H H O H O H V5+ O H H O H H O H 6H2O Metavanadate(V) ion VO3(aq) + 6H3O+ The reaction is favoured in alkaline solution VO3 is a polymeric anion like SiO32 84 Metavanadate(V) ion, (VO3)nn 85 H H H O H O H V5+ O H H O H H O H 4H2O VO2+(aq) + 4H3O+ The reaction is favoured in acidic solution 86 1. Variable Oxidation States of Vanadium and their Interconversions 2. The action of zinc powder and concentrated hydrochloric acid vanadium(V) ions can be reduced sequentially to vanadium(II) ions 87 1. Variable Oxidation States of Vanadium and their Interconversions VO2 +(aq) yellow Zn conc. HCl VO2+(aq) Zn conc. HCl blue V3+(aq) green 88 V2+(aq) Zn conc. HCl violet (a) (b) VO2+(aq) VO2+(aq) (c) (d) V3+(aq) V2+(aq) Colours of aqueous solutions of compounds containing vanadium in four different oxidation states: (a) +5; (b) +4; (c) +3; (d) +2 89 • The feasibility of the changes in oxidation state of vanadium can be predicted using standard electrode potentials Half reaction Zn2+(aq) + 2e– (V) –0.76 Zn(s) VO2+(aq) + 2H+(aq) + e– VO2+(aq) + H2O(l) +1.00 VO2+(aq) + 2H+(aq) + e– V3+(aq) + H2O(l) +0.34 V3+(aq) + e– 90 V2+(aq) –0.26 1. Variable Oxidation States of Vanadium and their Interconversions • Under standard conditions zinc can reduce 91 1. VO2+(aq) to VO2+(aq) >0 2. VO2+(aq) to V3+(aq) >0 3. V3+(aq) to V2+(aq) >0 1. Variable Oxidation States of Vanadium and their Interconversions 2 × (VO2+(aq) + 2H+(aq) + e– VO2+(aq) + H2O(l)) –) Zn2+(aq) + 2e– Zn(s) = +1.00 V = –0.76 V 2VO2+(aq) + Zn(s) + 4H+(aq) 2VO2+(aq) + Zn2+(aq) + 2H2O(l) = +1.76 V 92 1. Variable Oxidation States of Vanadium and their Interconversions 2 × (VO2+(aq) + 2H+(aq) + e– V3+(aq) + H2O(l)) –) Zn2+(aq) + 2e– Zn(s) = +0.34 V = –0.76 V 2VO2+(aq) + Zn(s) + 4H+(aq) 2V3+(aq) + Zn2+(aq) + 2H2O(l) = +1.10 V 93 1. Variable Oxidation States of Vanadium and their Interconversions 2 × (V3+(aq) + e– –) Zn2+(aq) + 2e– 2V3+(aq) + Zn(s) V2+(aq)) = –0.26 V Zn(s) = –0.76 V 2V2+(aq) + Zn2+(aq) = +0.50 V 94 2. Variable Oxidation States of Manganese and their Interconversions • Manganese show oxidation states of +2, +3, +4, +5, +6 and +7 in its compounds • The most common oxidation states +2, +4 and +7 95 Colours of compounds or ions of manganese in different oxidation states 96 Ion Oxidation state of manganese in the ion Colour Mn2+ +2 Very pale pink Mn(OH)3 +3 Dark brown Mn3+ +3 Red MnO2 +4 Black MnO43 +5 Bright blue MnO42– +6 Green MnO4– +7 Purple (a) (b) Mn2+(aq) Mn(OH)3(aq) (c) MnO2(s) Colours of compounds or ions of manganese in differernt oxidation states: (a) +2; (b) +3; (c) +4 97 (d) MnO42–(aq) (e) MnO4–(aq) Colours of compounds or ions of manganese in differernt oxidation states: (d) +6; (e) +7 98 2. Variable Oxidation States of Manganese and their Interconversions • Manganese of the oxidation state +2 the most stable at pH 0 Mn3+ +1.50V +1.51V MnO4 99 Mn2+ 1.18V +1.23V MnO2 Mn Mn(VII) Explosive on heating and extremely oxidizing +7 2 2KMnO4 heat +4 0 K2MnO4 + MnO2 + O2 in ON = 2(+2) = +4 in ON = (1) + (3) = 4 100 +6 Mn(VII) +7 2 4MnO4 + 4H+ light +4 0 4MnO2 + 2H2O + 3O2 in ON = 6(+2) = +12 in ON = 4(3) = 12 The reaction is catalyzed by light Acidified KMnO4(aq) is stored in amber bottle 101 Oxidizing power of Mn(VII) depends on pH of the solution In an acidic medium (pH 0) MnO4–(aq) + 8H+(aq) + 5e– Mn2+(aq) + 4H2O(l) = +1.51 V In a neutral or alkaline medium (up to pH 14) MnO4–(aq) + 2H2O(l) + 3e– MnO2(s) + 4OH (aq) = +0.59 V 102 Why is the Eo of MnO4 MnO42 Eo = +0.56V not affected by pH ? MnO4(aq) + e MnO42 Eo = +0.56V The reaction does not involve H+(aq) nor OH(aq) 103 In an acidic medium (pH 0) MnO4–(aq) + 8H+(aq) + 5e– Mn2+(aq) + 4H2O(l) = +1.51 V In a neutral or alkaline medium (up to pH 14) MnO4–(aq) + 2H2O(l) + 3e– MnO2(s) + 4OH (aq) = +0.59 V Under what conditions is the following conversion favoured? MnO4(aq) + e MnO42 Eo = +0.56V When [OH(aq)] > 1M 104 Predict if Mn(VI) Mn(VII) + Mn(IV) is feasible at (i) pH 0 and (ii) pH 14 (1) MnO42(aq) + 4H+(aq) + 2e MnO2(s) + 2H2O(l) Eo = +2.26V (2) MnO42(aq) + 2H2O(l) + 2e MnO2(s) + 4OH(aq) Eo = +0.60V (3) MnO4 + e MnO42 Eo = +0.56V At pH 0 (1) 2(3) 3MnO42(aq) + 4H+(aq) Eocell = +1.70V (feasible) At pH 14 (2) 2(3) 3MnO42(aq) + 2H2O(l) 2MnO4(aq) + MnO2(s) + 2H2O(l) Mn(VI) is unstable in acidic medium 2MnO4(aq) + MnO2(s) + 4OH(aq) Eocell = +0.04V (much less feasible) 105 Mn(IV) Oxidizing in acidic medium MnO2(s) + 4H+(aq) + 2e– Mn2+(aq) + 2H2O(l) = 1.23 V • Used in the laboratory production of chlorine MnO2(s) + 4HCl(aq) MnCl2(aq) + 2H2O(l) + Cl2(g) 106 Mn(IV) Reducing in alkaline medium • Oxidized to Mn(VI) in alkaline medium 2MnO2 + 4OH + O2 2MnO42 + 2H2O 107 MnO2 is oxidized to MnO42 in alkaline medium 2MnO2 + 4OH + O2 2MnO42 + 2H2O Suggest a scheme to prepare MnO4 from MnO2 1. 2MnO2 + 4OH + O2 2MnO42 + 2H2O 2. 3MnO42 + 4H+ 2MnO4 + MnO2 + 2H2O 3. Filter the resulting mixture to remove MnO2 7B 108 Cu+(aq) + e Cu(s) Eo = +0.52V Cu2+(aq) + 2e Cu(s) Eo = +0.34V Cu2+(aq) is more stable than Cu+(aq) The only copper(I) compounds which can be stable in water are those which are (i) insoluble (e.g. Cu2O, CuI, CuCl) (ii) complexed with ligands other than water e.g. [Cu(NH3)4]+ Cu+(aq) + e Cu(s) Under these conditions, [Cu+(aq)] 109 Equil. Position shifts to left Estimation of Cu2+ ions 2Cu2+(aq) + 4I(aq) 2CuI(s) + I2(aq) unknown excess white fixed I2(aq) + 2S2O32(aq) 2I(aq) + S4O62(aq) standard solution 110 Formation of Complexes • 111 Another striking feature of the d-block elements is the formation of complexes Formation of Complexes A complex is formed when a central metal atom or ion is surrounded by other molecules or ions which form dative covalent bonds with the central metal atom or ion. The molecules or ions that donate lone pairs of electrons to form the dative covalent bonds are called ligands. 112 Formation of Complexes • A ligand can be an ion or a molecule having at least one lone pair of electrons that can be donated to the central metal atom or ion to form a dative covalent bond 113 Formation of Complexes Complexes can be electrically neutral Ni(CO)4 positively charged [Co(H2O)6]3+ negatively charged [Fe(CN)6]3 114 A co-ordination compound is either a neutral complex e.g. Ni(CO)4 or made of a complex ion and another ion e.g. [Co(H2O)6]Cl3 [Co(H2O)6]3+ + 3Cl K3[Fe(CN)6] 3K+ + [Fe(CN)6]3 115 Criteria for complex formation 1. Presence of vacant and low-energy 3d, 4s, 4p and 4d orbitals in the metal atoms or ions to accept lone pairs from ligands. 2. High charge density of the central metal ions. 116 Diagrammatic representation of the formation of a complex 117 [Co(H2O)6]2+ Co : 3d Co2+ : 4s 3d 4d 4p 4d 4s 4p The six sp3d2 orbitals accept six lone pairs from six H2O. Arranged octahedrally to minimize repulsion between dative bonds. 118 sp3d2 hybridisation 1. Complexes with Monodentate Ligands A ligand that forms one dative covalent bond only is called a monodentate ligand. • 119 Examples: neutral CO, H2O, NH3 anionic Cl–, CN–, OH– 120 In the formation of complexes, classify the transition metal ion and the ligand as a Lewis acid or base. Explain your answer briefly. The transition metal ion is the Lewis acid since it accepts lone pairs of electrons from the ligands in forming dative covalent bonds. The ligand is the Lewis base since it donates a lone pair of electrons to the transition metal ion in forming dative covalent bonds. 121 What is the oxidation state of the central metal ? Cr3+ 122 Zn2+ What is the oxidation state of the central metal ? Co3+ 123 What is the oxidation state of the central metal ? Fe3+ 124 Co2+ 2. Complexes with Bidentate Ligands A ligand that can form two dative covalent bonds with a metal atom or ion is called a bidentate ligand. A ligand that can form more than one dative covalent bond with a central metal atom or ion is called a chelating ligand. 125 Ethylenediamine (H2NCH2CH2NH2) Oxalate (C2O42–) ethylenediamine oxalate ion The term chelate is derived from Greek, meaning ‘claw’. The ligand binds with the metal like the great claw of the lobster. 126 ethylenediamine 127 oxalate ion 3. Complexes formed by Multidentate Ligands Ligands that can form more than two dative covalent bonds to a metal atom or ion are called multidentate ligands. Some ligands can form as many as six bonds to a metal atom or ion. • Example: ethylenediaminetetraacetic acid (abbreviated as EDTA) 128 EDTA forms six dative covalent bonds with the metal ion through six atoms giving a very stable complex. hexadentate ligand ethylenediaminetetraacetate ion 129 Fe2+ EDTA4 2 ? [FeEDTA]2 Structure of the complex ion formed by iron(II) ions and EDTA 130 Uses of EDTA 1. Determining concentrations of metal ions by complexometric titrations e.g. determination of water hardness 2. In chelation therapy for mercury poisoning and lead poisoning Poisonous Hg2+ and Pb2+ ions are removed by forming stable complexes with EDTA. 3. Preparing buffer solutions ( K a to K a ) 4. As preservative to prevent catalytic oxidation of food by metal ions. 131 1 4 The coordination number of the central metal atom or ion in a complex is the number of dative covalent bonds formed by the central metal atom or ion in a complex. Complex The central metal atom Coordination or ion in the complex number [Ag(NH3)2]+ Ag+ 2 [Cu(NH3)4]2+ Cu2+ 4 [Fe(CN)6]3– Fe3+ 6 132 4. Nomenclature of Transition Metal Complexes with Monodentate Ligands IUPAC conventions 1. (a) For any ionic compound the cation is named before the anion (b) If the complex is neutral the name of the complex is the name of the compound 133 1. (c) In naming a complex (which may be neutral, a cation or an anion) the ligands are named before the central metal atom or ion the liqands are named in alphabetical order (prefixes not counted) (d) The number of each type of ligands are specified by the Greek prefixes 134 1 mono- 2 di 3 tri 4 tetra- 5 penta- 6 hexa- 1. (e) The oxidation number of the metal ion in the complex is indicated immediately after the name of the metal using Roman numerals K3[Fe(CN)6] potassium hexacyanoferrate(III) [CrCl2(H2O)4]Cl tetraaquadichlorochromium(III) chloride [CoCl3(NH3)3] triamminetrichlorocobalt(III) 135 2. (a) The root names of anionic ligands always end in “-o” CN– cyano OH hydroxo Cl– chloro NO2 nitro Br bromo SO42 sulphato I iodo H hydrido (b) The names of neutral ligands are the names of the molecules except NH3, H2O, CO and NO 136 137 Neutral ligand Ammonia (NH3) Name of ligand Water (H2O) Aqua Carbon monoxide (CO) Carbonyl Nitrogen monoxide (NO) Nitrosyl Ammine 3. (a) If the complex is anionic the suffix “-ate” is added to the end of the name of the metal, followed by the oxidation number of that metal 138 Formula Name of the complex [CoCl4]2 tetrachlorocobaltate(II) ion [Fe(CN)6]3 hexacyanoferrate(III) ion [CuCl4]2– tetrachlorocuprate(II) ion Names of some common metals in anionic complexes 139 Metal Name in anionic complex Titanium Titanate Vanadium Vanadate Chromium Chromate Manganese Manganate Iron Ferrate Cobalt Cobaltate Nickel Nickelate Copper Cuprate Zinc Zincate Platinum Platinate 3. (b) If the complex is cationic or neutral the name of the metal is unchanged followed by the oxidation number of that metal 140 Formula Name of the complex [CrCl2(H2O)4]+ tetraaquadichlorochromium(III) ion [CoCl3(NH3)3] triamminetrichlorocobalt(III) (a) Write the names of the following compounds. (i) [Fe(H2O)6]Cl2 (ii) [Cu(NH3)4]Cl2 (iii) [PtCl4(NH3)2] (iv) K2[CoCl4] (v) [Cr(NH3)4SO4]NO3 (vi) [Co(H2O)2(NH3)3Cl]Cl (vii) K3[AlF6] 141 (i) [Fe(H2O)6]Cl2 Hexaaquairon(II) chloride (ii) [Cu(NH3)4]Cl2 Tetraamminecopper(II) chloride (iii) [PtCl4(NH3)2] Diamminetetrachloroplatinum(IV) (iv) K2[CoCl4] Potassium tetrachlorocobaltate(II) (v) [Cr(NH3)4SO4]NO3 Tetraamminesulphatochromium(III) nitrate 142 (a) (vi) [Co(H2O)2(NH3)3Cl]Cl triamminediaquachlorocobalt(II) chloride (vii) K3[AlF6] potassium hexafluoroaluminate Al has a fixed oxidation state (+3) no need to indicate the oxidation state 143 (b) Write the formulae of the following compounds. (i) pentaamminechlorocobalt(III) chloride [Co(NH3)5Cl]Cl2 (ii) Ammonium hexachlorotitanate(IV) (NH4)2[TiCl6] (iii) Tetraaquadihydroxoiron(II) [Fe(H2O)4(OH)2] 144 Stereo-structures of complexes Coordination number of the central metal atom or ion Shape of complex Example sp hybridized [Ag(NH3)2]+ 2 [Ag(CN)2]– linear 145 Stereo-structures of complexes Coordination number of the central metal atom or ion Shape of complex sp3 Example [Zn(NH3)4]2+ [CoCl4]2+ 4 Tetrahedral dsp2 [Cu(NH3)4]2+ [CuCl4]2– Square planar 146 Tetra-coordinated Complexes (a) Tetrahedral complexes tetrahedral shape 147 [Co(H2O)6]2+ Octahedral, pink blue Tetra-coordinated Complexes (b) Square planar complexes have a square planar structure 148 Tetra-coordinated Complexes • 149 Example: Stereo-structures of complexes Coordination number of the central metal atom or ion Shape of complex Example sp3d2 [Cr(NH3)6]3+ 6 [Fe(CN)6]3– Octahedral 150 Hexa-coordinated Complexes • 151 Example: 6. Displacement of Ligands and Relative Stability of Complex Ions Different ligands have different tendencies to bind with the metal atom/ion ligands compete with one another for the metal atom/ion. A stronger ligand can displace a weaker ligand from a complex. 152 6. Displacement of Ligands and Relative Stability of Complex Ions Stronger ligand [Fe(H2O)6]2+(aq) + 6CN–(aq) Hexaaquairon(II) ion Weaker ligand [Fe(CN)6]4–(aq) + 6H2O(l) Hexacyanoferrate(II) ion Reversible reaction Kst 1024 mol6 dm18 Equilibrium position lies to the right 153 Stronger ligand [Ni(H2O)6]2+(aq) + 6NH3(aq) Hexaaquanickel(II) ion Weaker ligand [Ni(NH3)6]2+(aq) + 6H2O(l) Hexaamminenickel(II) ion The greater the equilibrium constant, the stronger is the ligand on the LHS and the more stable is the complex on the RHS The equilibrium constant is called the stability constant, Kst 154 Consider the general equilibrium system below, [M(H2O)x]m+ + xLn [M(L)x](m-xn)+ + xH2O [[M(L)x ] ] Kst m n x [[M(H2O)x ] ][L ] (m xn) Units = (mol dm3)-x Kst measures the stability of the complex, [M(L)x](m-xn)+, relative to the aqua complex, [M(H2O)x]m+ 155 monodentate bidentate multidentate TAS Expt 6 156 Relative strength of some ligands bonding with copper(II) ions Equilibrium [Cu(H2O)4]2+(aq) + 4Cl–(aq) Kst ((mol dm–3)–n) 4.2 × 105 [CuCl4]2–(aq) + 4H2O(l) [Cu(H2O)4]2+(aq) + 4NH3(aq) 1.1 × 1013 [Cu(NH3)4]2+(aq) + 4H2O(l) [Cu(H2O)4]2+(aq) + 2H2NCH2CH2NH2(aq) 1.0 × 1018.7 [Cu(H2NCH2CH2NH2)2]2+(aq) + 4H2O(l) [Cu(H2O)4]2+(aq) + EDTA4–(aq) 1.0 × 1018.8 [CuEDTA]2–(aq) + 4H2O(l) What is the Kst of the formation of [Cu(H2O)4]2+(aq) ? 157 [Cu(H2O)4]2+ + 4H2O [Cu(H2O)4]2+ + 4H2O [[Cu(H2O)4 ]2 ] Kst 1 2 [[Cu(H2O)4 ] ] 158 Factors affecting the stability of complexes 1. The charge density of the central ion Equilibrium [Co(H2O)6]2+(aq) + 6NH3(aq) Kst (mol6 dm18) 7.7 × 104 [Co(NH3)6]2+(aq) + 6H2O(l) [Co(H2O)6]3+(aq) + 6NH3(aq) 4.5 × 1033 [Co(NH3)6]3+(aq) + 6H2O(l) [Fe(H2O)6]2+(aq) + 6CN–(aq) ≈ 1024 [Fe(CN)6]4–(aq) + 6H2O(l) [Fe(H2O)6]3+(aq) + 6CN–(aq) ≈ 1031 [Fe(CN)6]3–(aq) + 6H2O(l) 159 Factors affecting the stability of complexes 2. The nature of ligands Ability to form complex : CN > NH3 > Cl > H2O [Zn(CN)4]2 Kst = 5 1016 mol4 dm12 [Zn(NH3)4]2+ Kst = 3.8 109 mol4 dm12 [Cu(NH3)4]2+ Kst = 1.1 1013 mol4 dm12 [CuCl4]2+ Kst = 4.2 105 mol4 dm12 160 Factors affecting the stability of complexes 3. The pH of the solution In acidic solution, the ligands are protonated lone pairs are not available the complex decomposes [Cu(NH3)4]2+(aq) + 4H2O(l) [Cu(H2O)4]2+(aq) + 4NH3(aq) H+(aq) Equilibrium position shifts to the right 161 NH4+(aq) Consider the stability constants of the following silver complexes: Ag+(aq) + 2Cl–(aq) [AgCl2]–(aq) Kst = 1.1 × 105 mol–2 dm6 Ag+(aq) + 2NH3(aq) [Ag(NH3)2]+(aq) Kst = 1.6 × 107 mol–2 dm6 Ag+(aq) + 2CN–(aq) [Ag(CN)2]–(aq) Kst = 1.0 × 1021 mol–2 dm6 What will be formed when CN–(aq) is added to a solution of [Ag(NH3)2]+? [Ag(CN)2](aq) and NH3 162 Consider the stability constants of the following silver complexes: Ag+(aq) + 2Cl–(aq) [AgCl2]–(aq) Kst = 1.1 × 105 mol–2 dm6 Ag+(aq) + 2NH3(aq) [Ag(NH3)2]+(aq) Kst = 1.6 × 107 mol–2 dm6 Ag+(aq) + 2CN–(aq) [Ag(CN)2]–(aq) Kst = 1.0 × 1021 mol–2 dm6 What will be formed when NH3(aq) is added to a solution of [Ag(CN)2]–? No apparent reaction 163 FeSO4(aq) is used as the antidote for cyanide poisoning [Fe(H2O)6]2+(aq) + 6CN(aq) Kst 1 1024 mol6 dm18 [Fe(CN)6]4 + 6H2O(l) Very stable Only free CN is poisonous Why is Fe2(SO4)3(aq) not used as the antidote ? Fe3+(aq) is too acidic. [Fe(H2O)6]3+(aq) + H2O(l) [Fe(H2O)5OH]2+(aq) + H3O+(aq) 164 [Cu(H2O)4]2+(aq) + Cl(aq) [Cu(H2O)3Cl]+(aq) + H2O(l) [Cu(H2O)3Cl]+(aq) + Cl(aq) [Cu(H2O)2Cl2](aq) + H2O(l) [Cu(H2O)2Cl2](aq) + Cl(aq) K3 = 5.4 mol1 dm3 [Cu(H2O)Cl3](aq) + H2O(l) [Cu(H2O)Cl3](aq) + Cl(aq) K1 = 3.1 mol1 dm3 [CuCl4]2(aq) + H2O(l) [Cu(H2O)4]2+(aq) + 4Cl(aq) [CuCl4]2(aq) + 4H2O(l) K1 = 6.3102 mol1 dm3 K2 = 40 mol1 dm3 Kst = K1 K2 K3 K4 = 4.2 105 mol4 dm12 165 K1 > K2 > K3 > K4 Reasons : 1. Statistical effect On successive displacement, less water ligands are available to be displaced. 166 K1 > K2 > K3 > K4 Reasons : 2. Charge effect On successive displacement, the Cl experiences more repulsion from the complex [Cu(H2O)4]2+ [Cu(H2O)Cl3] 167 Cl Cl attraction repulsion Colours of some copper(II) complexes Formula of copper(II) complex Colour of the complex [Cu(H2O)4]2+ Pale blue [CuCl4]2– Yellow [Cu(NH3)4]2+ Deep blue Violet [Cu(H2NCH2CH2NH2)]2+ [Cu(EDTA)]2– Sky blue The displacement of ligands are usually accompanied with easily observable colour changes 168 Coloured Ions The colours of many gemstones are due to the presence of small quantities of d-block metal ions 169 Coloured Ions • Most of the d-block metals form coloured compounds due to the presence of the incompletely filled d orbitals in the d-block metal ions Which aqueous transition metal ion(s) is/are not coloured ? 3d10 : Zn2+, Cu+; 170 3d0 : Sc3+, Ti4+ Colours of some d-block metal ions in aqueous solutions Number of unpaired electrons in 3d orbitals 0 1 171 d-Block metal ion Colour in aqueous solution Sc3+ Colourless Ti4+ Colourless Zn2+ Colourless Cu+ Colourless Ti3+ Purple V4+ Blue Cu2+ Blue Colours of some d-block metal ions in aqueous solutions Number of unpaired electrons in 3d orbitals 2 3 172 d-Block metal Colour in ion aqueous solution V3+ Green Ni2+ Green V2+ Violet Cr3+ Green Co2+ Pink Colours of some d-block metal ions in aqueous solutions Number of unpaired electrons in 3d orbitals 4 5 173 d-Block metal Colour in ion aqueous solution Cr2+ Blue Mn3+ Violet Fe2+ Green Mn2+ Very pale pink Fe3+ Yellow Co2+(aq) Zn2+(aq) Fe3+(aq) Colours of some d-block metal ions in aqueous solutions 174 Mn2+(aq) Fe2+(aq) Cu2+(aq) Colours of some d-block metal ions in aqueous solutions 175 A substance absorbs visible light of a certain wavelength reflects or transmits visible light of other wavelengths (complimentary colour) appears coloured Light Light reflected or absorbed transmitted [Cu(H2O)4]2+(aq) Yellow Blue [CuCl4]2(aq) Blue Yellow Coloured ion 176 Complimentary colour chart Blue light absorbed Appears yellow Violet Blue Cyan Yellow light absorbed Appears blue 177 Green Magenta Red Yellow Greenish yellow Coloured Ions • The absorption of visible light is due to the d-d electronic transition 3d 3d i.e. an electron jumping from a lower 3d orbital to a higher 3d orbital 178 In gaseous state, the five 3d orbitals are degenerate i.e. they are of the same energy level In the presence of ligands, The five 3d orbitals interact with the orbitals of ligands and split into two groups of orbitals with slightly different energy levels 179 distributes along z axis Interact more strongly with the orbitals of ligands distributes along x and y axes d z 2 , d x2 y 2 eg d xy , d xz , d yz The splitting of the degenerate 3d orbitals of a d-block metal ion in an octahedral complex 180 t 2g d x 2 y 2 181 Higher energy eg Criterion for d-d transition : presence of unpaired d electrons in the dblock metal atoms or ions Or presence of incompletely filled d-subshell d-d transition is possible for 3d1 to 3d9 arrangements d-d transition is NOT possible for 3d0 & 3d10 arrangements 182 H2O as ligand Cu2+ 3d9 : d-d transition is possible 183 Yellow light absorbed, appears blue H2O as ligand *Cu2+ 3d9 : d-d transition is possible 184 Fe2+ 3d6 : d-d transition is possible 185 Magenta light absorbed, appears green *Fe2+ 3d6 : d-d transition is possible 186 Zn2+ 3d10 : d-d transition NOT possible 187 Sc3+ 3d0 : d-d transition NOT possible 188 E E depends on [Fe(H2O)6]2+ green, [Fe(H2O)6]3+ yellow 1. the nature and charge of metal ion 2. the nature of ligand 189 [Cu(H2O)4]2+ blue, [CuCl4]2 yellow Coloured Ions Why does Na+(aq) appear colourless ? 3d0 : d-d transition is NOT possible 2p 3s transition involves absorption of radiation in the UV region. 190 Catalytic Properties of Transition Metals and their Compounds • The d-block metals and their compounds important catalysts in industry and biological systems 191 The use of some d-block metals and their compounds as catalysts in industry d-Block metal Catalyst V V2O5 or vanadate(V) (VO3–) Reaction catalyzed Contact process 2SO2(g) + O2 (g) 2SO3(g) Haber process Fe 192 Fe N2(g) + 3H2(g) 2NH3(g) The use of some d-block metals and their compounds as catalysts in industry d-Block metal Ni Pt 193 Catalyst Reaction catalyzed Ni Hardening of vegetable oil (Manufacture of margarine) RCH = CH2 + H2 RCH2CH3 Pt Catalytic oxidation of ammonia (Manufacture of nitric(V) acid) 4NH3(g) + 5O2(g) 4NO(g) + 6H2O(l) Catalytic Properties of Transition Metals and their Compounds • The d-block metals and their compounds exert their catalytic actions in either heterogeneous catalysis homogeneous catalysis 194 Catalytic Properties of Transition Metals and their Compounds • Generally speaking, the function of a catalyst provides an alternative reaction pathway of lower activation energy enables the reaction to proceed faster than the uncatalyzed one 195 1. Heterogeneous Catalysis • The catalyst and reactants exist in different states • The most common heterogeneous catalysts finely divided solids for gaseous reactions 196 1. Heterogeneous Catalysis A heterogeneous catalyst provides a suitable reaction surface for the reactants to come close together and react. 197 1. Heterogeneous Catalysis • Example: The synthesis of gaseous ammonia from nitrogen and hydrogen (i.e. Haber process) N2(g) + 3H2(g) 198 2NH3(g) 1. Heterogeneous Catalysis • In the absence of a catalyst the formation of gaseous ammonia proceeds at an extremely low rate • The probability of collision of four gaseous molecules (i.e. one nitrogen and three hydrogen molecules) very small 199 1. Heterogeneous Catalysis • The four reactant molecules collide in proper orientation in order to form the product • The bond enthalpy of the reactant (N N), very large the reaction has a high activation energy 200 1. Heterogeneous Catalysis • In the presence of iron as catalyst the reaction proceeds much faster provides an alternative reaction pathway of lower activation energy 201 1. Heterogeneous Catalysis • Fe is a solid • H2, N2 and NH3 are gases • The catalytic action occurs at the interface between these two states • The metal provides an active reaction surface for the reaction to occur 202 1. Heterogeneous Catalysis 1. Gaseous nitrogen and hydrogen molecules diffuse to the surface of the catalyst 2. The gaseous reactant molecules adsorbed (i.e. adhered) on the surface of the catalyst 203 1. Heterogeneous Catalysis 2. The iron metal many 3d electrons and low-lying vacant 3d orbitals form bonds with the reactant molecules adsorb them on its surface 204 weakens the bonds present in the reactant molecules 1. Heterogeneous Catalysis 2. The free nitrogen and hydrogen atoms come into contact with each other readily to react and form the product 3. The weak interaction between the product and the iron surface gaseous ammonia molecules desorb easily 205 The catalytic mechanism of the formation of gaseous ammonia from nitrogen and hydrogen 206 The catalytic mechanism of the formation of gaseous ammonia from nitrogen and hydrogen 207 The catalytic mechanism of the formation of gaseous ammonia from nitrogen and hydrogen 208 The catalytic mechanism of the formation of gaseous ammonia from nitrogen and hydrogen 209 The catalytic mechanism of the formation of gaseous ammonia from nitrogen and hydrogen 210 43.3 Characteristic Properties of the d-Block Elements and their compound (SB p.162) 1. Heterogeneous Catalysis • Sometimes, the reactants in aqueous or liquid state • Other example: The decomposition of hydrogen peroxide 2H2O2(aq) 2H2O(l) + O2(g) MnO2(s) as the catalyst 211 Energy profiles of the reaction of nitrogen and hydrogen to form gaseous ammonia in the presence and absence of a heterogeneous catalyst 212 2. Homogeneous Catalysis • A homogeneous catalyst the same state as the reactants and products the catalyst forms an intermediate with the reactants in the reaction changes the reaction mechanism to an another one with a lower activation energy 213 2. Homogeneous Catalysis In homogeneous catalysis, the ability of the d-block metals to exhibit variable oxidation states enables the formation of the reaction intermediates. • Example: The reaction between peroxodisulphate(VI) ions (S2O82–) and iodide ions (I–) 214 2. Homogeneous Catalysis • Peroxodisulphate(VI) ions oxidize iodide ions to iodine in an aqueous solution themselves being reduced to sulphate(VI) ions S2O82–(aq) + 2I–(aq) 2SO42–(aq) + I2 (aq) Eocell 1.51 V 215 2. Homogeneous Catalysis • • The reaction is very slow due to strong repulsion between like charges. Iron(III) ions take part in the reaction by oxidizing iodide ions to iodine themselves being reduced to iron(II) ions 2I–(aq) + 2Fe3+(aq) I2(aq) + 2Fe2+(aq) 216 = +0.23 V 2. Homogeneous Catalysis • Iron(II) ions subsequently oxidized by peroxodisulphate(VI) ion the original iron(III) ions are regenerated 2Fe2+(aq) + S2O82–(aq) 2Fe3+(aq) + 2SO42–(aq) 217 = +1.28 V 2. Homogeneous Catalysis • The overall reaction: 2I–(aq) + 2Fe3+(aq) I2(aq) + 2Fe2+(aq) +) = +0.23 V 2Fe2+(aq) + S2O82–(aq) 2Fe3+(aq) + 2SO42–(aq) = +1.28 V S2O82–(aq) + 2I–(aq) 2SO42–(aq) + I2(aq) = +1.51 V Feasible reaction 218 43.3 Characteristic Properties of the d-Block Elements and their compound (SB p.164) 2. Homogeneous Catalysis • Iron(III) ions catalyze the reaction acting as an intermediate for the transfer of electrons between peroxodisulphate(VI) ions and iodide ions 219 2. Homogeneous Catalysis • Peroxodisulphate(VI) ions oxidize Fe2+ to Fe3+ • Iodide ions reduce Fe3+ to Fe2+ 220 The End 221 Check Point 43-3E 222 Energy profiles for the oxidation of iodide ions by peroxodisulphate (VI) ions in the presence and absence of a homogeneous catalyst Besides iron(III) ions, iron(II) ions can also catalyze the reaction between peroxodisulphate(VI) ions and iodide ions. Why? Answer 223 Iron(II) ions catalyze the reaction by reacting with the peroxodisulphate(VI) ions first. 2Fe2+(aq) + S2O82–(aq) 2Fe3+(aq) + 2SO42–(aq) The iron(III) ions formed then oxidize the iodide ions. 2Fe3+(aq) + 2I–(aq) 2Fe2+(aq) + I2(aq) In this way, the reaction between peroxodisulphate(VI) ions and iodide ions is catalyzed. Back 224 Which of the following redox systems might catalyze the oxidation of iodide ions by peroxodisulphate(VI) ions in an aqueous solution? Cr2O72–(aq) + 14H+(aq) + 6e– 2Cr3+(aq) + 7H2O(l) Answer = +1.33 V MnO4–(aq) + 8H+(aq) + 5e– Mn2+(aq) + 4H2O(l) = +1.52 V Sn4+(aq) + 2e– = +0.15 V Sn2+(aq) (Given: S2O82–(aq) + 2e– 225 I2(aq) + 2e– 2SO42–(aq) 2I–(aq) = +2.01 V = +0.54 V) Those redox systems with greater than +0.54 V and smaller than +2.01 V are able to catalyze the oxidation of iodide ions by peroxodisulphate(VI) ions in an aqueous solution. Therefore, the following two redox systems are able to catalyze the reaction. Cr2O72–(aq) + 14H+(aq) + 6e– 2Cr3+(aq) + 7H2O(l) MnO4–(aq) + 8H+(aq) + 5e– Mn2+(aq) + 4H2O(l) Back 226