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APPLIED INORGANIC CHEMISTRY FOR CHEMICAL ENGINEERS Transition Metal Chemistry CHEM261HC/SS1/01 Periodic table q Elements are divided into four categories Main-group elements Transition metals 1.Main-group elements 2.Transition metals 3.Lanthanides 4.Actinides Main-group element s Lanthanides Actinides CHEM261HC/SS1/02 Transition metals vs. Main -group elements Transition metals Main-group elements v Metals • malleable and ductile • conduct heat and electricity • form positive ions v Transition metals • more electronegative than the main group metals • more likely to form covalent compounds • easily form complexes • form stable compounds with neutral molecules q There is some controversy about the classification of the elements i.e. Zinc (Zn), Cadmium (Cd) and Mercury (Hg) CHEM261HC/SS1/03 Electron configuration of Transition-metal ions § The relationship between the electron configurations of transition-metal elements and their ions is complex. Example § Consider the chemistry of cobalt which forms complexes that contain either Co 2+ or Co 3+ ions. Co: [Ar] 4s2 3d7 Co 2+: [Ar] 3d7 Co 3+: [Ar] 3d6 v In general, electrons are removed from the valence shell s orbitals before they are removed from valence d orbitals when transition metals are ionized. CHEM261HC/SS1/04 q How do we determine the electronic configuration of the central metal ion in any complex? • Try to recognise all the entities making up the complex • Need knowing whether the ligands are neutral or anionic • Then you can determine the oxidation state of the metal ion. A simple procedure exists for the M(II) case. 22 23 24 25 26 27 28 29 Ti V Cr Mn Fe Co Ni Cu 4 5 6 7 8 9 Cross off the first 2, 2 3 CHEM261HC/SS1/05 EXAMPLES Elements Configuration Sc 4s23d1 V 4s23d3 Cr 4s13d5 Fe 4s23d6 Ni 4s23d8 Cu 4s13d10 Zn 4s23d10 Evaluating the oxidation state [CoCl(NO 2)(NH 3)4]+ +1 Neutral zero charge x - 2 = +1 x = +3 Co3+ CHEM261HC/SS1/06 Oxidation states and their relative stabilities v Why do these elements exhibit a variety of oxidation states? ü Because of the closeness of the 3d and 4s energy states Sc +3 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 Cu +1 +2 +3 Zn +7 +2 § The most prevalent oxidation numbers are shown in green. CHEM261HC/SS1/07 § An increase in the No. of oxidation states from Sc to Mn. § All seven oxidation states are exhibited by Mn. § There is a decrease in the No. of oxidation states from Mn to Zn. WHY? ü Because the pairing of d-electrons occurs after Mn (Hund's rule) which in turn decreases the number of available unpaired electrons and hence, the number of oxidation states. § The stability of higher oxidation states decreases in moving from Sc to Zn. § Mn(VII) and Fe(VI) are powerful oxidizing agents and the higher oxidation states of Co, Ni and Zn are unknown. CHEM261HC/SS1/09 § The relative stability of +2 state with respect to higher oxidation states increases in moving from left to right. On the other hand +3 state becomes less stable from left to right. § This is justifiable since it will be increasingly difficult to remove the third electron from the d-orbital. Example 22 23 24 25 26 27 28 29 Ti V Cr Mn Fe Co Ni Cu M = [Ar]4s23dx M+2 = [Ar]3dx M+3 = [Ar]3dx-1 loss of the two s electrons more difficult CHEM261HC/SS1/10 Chromium • Oxidized by HCl or H2SO4 to form blue Cr2+ ion • Cr2+ oxidized by O2 in air to form green Cr3+ Assignment 1 ü Write down balance equations that show the two reactions • Cr also found in +6 state as in CrO42− and Cr2O72− are strong oxidizer Assignment 2 ü Use balanced equations to show that and Cr2O72− are strong oxidizing agents CrO42− Assignment 1 Solution 2 Cr(S) + H2SO4(aq) 2 Cr(s) + 4 HCl (aq) CrCl 2(aq) + 2O 2(g) Cr2SO4)(aq) + H2(g) 2 CrCl 2(aq) + 2H 2(g) CrOCl (aq) Iron • Exists in solution in +2 or +3 state • Elemental iron reacts with non-oxidizing acids to form Fe2+, which oxidizes in air to Fe3+ • Brown water running from a faucet is caused by insoluble Fe2O3 • Fe3+ soluble in acidic solution, but forms a hydrated oxide as red-brown gel in basic solution Assignment 3 ü Complete and balance the following equation Fe2O3 + HCl Coordination Chemistry q A coordination compound (complex), contains a central metal atom (or ion) surrounded by a number of oppositely charged ions or neutral molecules (possessing lone pairs of electrons) which are known as ligands. q If a ligand is capable of forming more than one bond with the central metal atom or ion, then ring structures are produced which are known as metal chelates § the ring forming groups are described as chelating agents or polydentate ligands. q The coordination number of the central metal atom or ion is the total number of sites occupied by ligands. § Note: a bidentate ligand uses two sites, a tridentate three sites etc. CHEM261HC/SS1/13 Ligands molecular formula Lewis base/ligand Lewis acid donor atom coordination number [Zn(CN)4]2- CN- Zn2+ C 4 [PtCl6]2- Cl- Pt4+ Cl 6 [Ni(NH3)6]2+ NH 3 Ni2+ N 6 CHEM261HC/SS1/14 Mono-dentate Multidentate ligands Abbreviation en ox2- EDTA4- Name Formula Ethylenediamine Oxalato Ethylenediaminetetraacetato CHEM261HC/SS1/15 q Chelating ligands bond to metal § forms rings – chelate rings Ø Five or six atoms rings are common (i.e. including metal) q Coordination numbers and geometries Linear Square planar Tetrahedral Octahedral CHEM261HC/SS1/16 Nomenclature of Coordination Compounds — The basic protocol in coordination nomenclature is to name the ligands attached to the metal as prefixes before the metal name. — Some common ligands and their names are listed above. — As is the case with ionic compounds, the name of the cation appears first; the anion is named last. — Ligands are listed alphabetically before the metal. Prefixes denoting the number of a particular ligand are ignored when alphabetizing. — The names of anionic ligands end in “o”; the endings of the names of neutral ligands are not changed. — Prefixes tell the number of a type of ligand in the complex. If the name of the ligand itself has such a prefix, alternatives like bis-, tris-, etc., are used. — If the complex is an anion, its ending is changed to -ate. — The oxidation number of the metal is listed as a roman numeral in parentheses immediately after the name of the metal. Exercise 1 Name the following coordination complexes: (i) Cr(NH 3)Cl3 (ii) Pt(en)Cl2 (iii) [Pt(ox)2]2- Exercise 2 Give the structures of the following coordination complexes: (i) Tris(acetylacetanato)iron(III) (ii) Hexabromoplatinate(2-) (iii) Potassium diamminetetrabromocobaltate(III) Solutions (i) Cr(NH3)Cl3 chromium (III) ammine tri chloro Amminetrichlorochromium(III) (ii) Pt(en)Cl2 Platinum (II) ethylenediammine dichloro Dichloroethylenediammineplat inum(II) (iii) [Pt(ox)2]2Platinate (II) Dioxalatoplatinate(II) dioxalato Solutions (i) Tris(acetylacetanato)iron(III) Fe(acac) 3 Fe3+ (acac)3 (ii) Hexabromoplatinate(2-) [PtBr6]2- Pt Br6 [ ]2- (ii) Potassium diamminetetrabromocobaltate(III) K K[Co(NH3)2Br4] (NH3)2 Br4 3+ Co Isomers q Primarily in coordination numbers 4 and 6. q Arrangement of ligands in space and also the ligands themselves. Types Ionization isomers § Isomers can produce different ions in solution e.g. [PtCl2(NH 3)4]Br2 D [PtBr 2(NH 3)4]Cl2 Polymerization isomers § Same stoichiometry, different arrangement in space. § Different compounds with similar formula e.g. [Co(NH 3)3 (NO 2)3 ]° [MX x B b ] n ( n = 1) [Co(NH 3)4 (NO 2)2 ] + [Co(NH 3)2 (NO 2 )4] − ( n = 2) [Co(NH 3)6 ]3+ [Co(NO 2)6 ] 3− ( n = 2) CHEM261HC/SS1/17 Hydration isomers § Hydration isomers exist for crystals of complexes containing water molecules e.g. CrCl3·6H2O v exist in three different crystalline forms, in which the number of water molecules directly attached to the Cr 3+ ion differs [Cr(H2O) 4 Cl2]Cl·2H2O dark green [Cr(H2O) 5 Cl]Cl2·H2O light green [Cr(H2O) 6 ]Cl3 gray-blue § In each case, the coordination number of the chromium cation is 6 Coordination isomers § In compounds, both cation and anion are complex, the distribution of ligands can vary, giving rise to isomers. [Co(NH 3)6]3+ [Cr(CN) 6]-3 and [Cr(NH 3)6]+3 [Co(CN) 6]-3 Linkage isomers e.g. Nitro and nitito N or O coordination possible Yellow (a) [Co(NO 2)(NH3)5]2+ (b) [Co(ONO)(NH 3)5 ]2+ Red CHEM261HC/SS1/18 Geometric isomers § Formula is the same but the arrangement in 3-D space is different. e.g. square planar molecules give cis and trans isomers. CHEM261HC/SS1/19 For hexacoordinate systems CHEM261HC/SS1/20 For M(X)3(Y)3 systems there is facial and meridian CHEM261HC/SS1/21 Stereoisomer § Are “stereoisomers” also possible? § An analogy to organic chirality. § Molecules which can rotate light. § Enantiomers – non-superimposable mirror images CHEM261HC/SS1/22 Complex Stabilities § Generally in aqueous solution, for a given metal and ligand, complexes where the metal oxidation state is +3 are more stable than +2 § Generally the stabilities of complexes of the first row of transition metals vary in reverse of their cationic radii MnII < FeII < CoII < NiII > CuII > ZnII § Hard and soft Lewis acid-base theory v Hard acids and bases tend to have: • small atomic/ionic radius • high oxidation state • low polarizabilty • high electronegativity • hard bases - energy low-lying HOMO • hard acids - energy high-lying LUMO CHEM261HC/SS1/23 § Chelate effect - is the additional stability of a complex containing a chelating ligand, relative to that of a complex containing monodentate ligands with the same type and number of donors as in the chelate. [Cu(H 2O)4(NH3)2]2+ + en [Cu(H 2O)4(en)] 2+ + 2 NH 3 CHEM261HC/SS1/24 § Mainly an entropy effect. Cu(H2O) 4(NH3)2]2+ + en = [Cu(H2O) 4(en)]2+ + 2 NH3 § When ammonia molecule dissociates - swept off in solution and the probability of returning is remote. § When one amine group of en dissociates from complex ligand retained by end still attached so the nitrogen atom cannot move away – swings back and attach to metal again. § Therefore the complex dissociating. has a smaller probability of CHEM261HC/SS1/25 CHEM261HC/SS1/26 Metal carbonyl § Compounds that have the metal bonded to the carbon monoxide, giving a general formula of M(CO) n M + CO M C O M(CO)n ∏-orbitals in CO are very empty Molecular orbital diagram (CO) Molecular orbital diagram (CO) Therefore the bond order is: 4–1=3 Bond order: No. of e- pairs in the bonding orbital — No. of e- pairs in the anti-bonding orbital Back-bonding (back donation) q Formation of ∏-bonding as a result of the overlap of metal d ∏orbitals and the ligand, CO, ∏* orbitals Effects: § It enhances the bonding strength between the metal and the ligand. § The metal-ligand bond is shortened (M § The C O CO) becomes longer, weaker and the bond order decreases Evidence and extent § IR spectra – Vibration frequency – The greater the extent of back bonding the lower the C stretching frequency (bond order decreases) Free C O ≈ 2143 cm -1 M O CO ≈ 1900 - 2125 cm -1 Effect of replacing the CO ligands Non- ∏ accepting ligands (donor ligands) Cr(CO)6 2100 cm-1 2000 cm-1 1985 cm-1 § Trien Cr(triens)(CO)3 1900 cm-1 1760 cm-1 Replacement of the 3 x (CO) groups with donor ligands (trien) increases ∏acidity of the remaining ligands (CO) so as to counter the accumulation of the negative charge on the metal centre Effect of introducing a positive charge on metal complex V(C O) 6 1860 cm -1 § 1 proton V(CO )6 1 proton 2000 cm -1 V(CO) 6 + 2090 cm -1 Introducing a +ve charge on the metal inhibits shift of electrons from metal to empty ∏*- orbital of the CO ligands – This weakens ∏-bonding or decrease stretching frequencies of M-C O increases. (wave number or frequenc y increases) while the C Thought Ø V(CO) - and Cr(CO) are isoelectronic yet stretching frequencies of CO in V(CO) 6 is lower than that of CO in Cr(CO) 6 ? The origin of colour - absorption CHEM261HC/SS1/27 Colours on coordination compounds The colour can change depending on a number of factors e.g. Ø Metal charge Ø Ligand Physical phenomenon CHEM261HC/SS1/29 Are there any simple theories to explain the colours in transition metal complexes? § There is a simple electrostatic model used by chemists to rationalize the observed results This theory is called Crystal Field Theory § It is not a rigorous bonding theory but merely a simplistic approach to understanding the possible origins of photoand electrochemical properties of the transition metal complexes CHEM261HC/SS1/30