Transition Metal Chemistry

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1 APPLIED INORGANIC CHEMISTRY FOR CHEMICAL ENGINEERS Transition Metal Chemistry CHEM261HC/SS1/01 Periodic table Elements are divided into four categories Main-group elements Transition metals 1. Main-group elements 2. Transition metals 3. Lanthanides 4. Actinides Main-group elements Lanthanides Actinides CHEM261HC/SS1/02

2 Transition metals vs. Main-group metals Transition metals Main-group elements Main-group metals malleable and ductile conduct heat and electricity form positive ions Transition metals more electronegative than the main group metals more likely to form covalent compounds easily form complexes form stable compounds with neutral molecules forms one or more stable ions which have incompletely filled d orbitals There is some controversy about the classification of the elements i.e. Zinc (Zn), Cadmium (Cd) and Mercury (Hg) 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 has 27 electrons Co: [Ar] 4s 2 3d 7 [Ar] has 18 electrons Co 2+ : [Ar] 3d 7 Co 3+ : [Ar] 3d 6 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

3 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 Ti V Cr Mn Fe Co Ni Cu Cross off the first 2, d 2 d 3 d 4 d 5 d 6 d 7 d 8 d 9 CHEM261HC/SS1/05 Elements Outer e- configuration EXAMPLES Sc 4s 2 3d 1 V 4s 2 3d 3 Pauli exclusion principle Hand's rule Half-filled or Filled subshell Cr 4s 1 3d 5 Fe 4s 2 3d 6 Ni 4s 2 3d 8 Cu 4s 1 3d 10 Zn 4s 2 3d 10

4 Evaluating the oxidation state [CoCl(NO 2 )(NH 3 ) 4 ] + +1 x - 2 = +1 x = +3 Co 3+ Neutral zero charge CHEM261HC/SS1/06 Exercise Question 1 School of Chemistry UNIVERSITY OF KWAZULU-NATAL, HOWARD COLLEGE April 2008 Test CHEM261: APPLIED INORGANIC CHEMISTRY FOR CHEMICAL ENGINEERS Marks : 40 Time : 45 minutes NAME : STUDENT NO. : (a) For [Co(Br)(NO 2 )(H 2 O) 4 ] +, write (i) the coordination number for Co (1) (ii) the oxidation state for Co (1) (b) Write the outer electron configuration of (i) Co (1) (ii) Mn 2+ (1)

5 Oxidation states and their relative stabilities Why do these elements exhibit a variety of oxidation states? Because of the closeness of the 3d and 4s energy states Sc +3 Ti V Cr Mn Fe Co Ni Cu Zn +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? 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. Increase in effective nuclear charge across (L R) Mn(VII) and Fe(VI) are powerful oxidizing agents and the higher oxidation states of Ni, Cu and Zn are unknown. CHEM261HC/SS1/09

6 Therelativestabilityof+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 Ti V Cr Mn Fe Co Ni Cu M =[Ar]4s 2 3d x M +2 =[Ar]3d x M +3 =[Ar]3d x-1 loss of the two s electrons more difficult CHEM261HC/SS1/10 Chromium Oxidized by HCl or H 2 SO 4 to form blue Cr 2+ ion Cr 2+ oxidized by O 2 in air to form green Cr 3+ Cr also found in +6 state as in CrO 2 4 and Cr 2 O 2 7 are strong oxidizer

7 Iron Exists in solution as +2 or +3 state Elemental iron reacts with non-oxidizing acids to form Fe 2+, which oxidizes in air to Fe 3+ Brown water running from a faucet is caused by insoluble Fe 2 O 3 Fe 3+ soluble in acidic solution, but forms a hydrated oxide as red-brown gel in basic solution Coordination Chemistry 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. 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. 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

8 Ligands molecular formula Lewis base/ligand Lewis acid donor atom coordination number [Zn(CN) 4 ] 2- CN - Zn 2+ C 4 [PtCl 6 ] 2- Cl - Pt 4+ Cl 6 [Ni(NH 3 ) 6 ] 2+ NH 3 Ni 2+ N 6 CHEM261HC/SS1/14 Mono-dentate Multidentate ligands Abbreviation Name Formula en Ethylenediamine ox 2- Oxalato EDTA 4- Ethylenediaminetetraacetato CHEM261HC/SS1/15

9 Chelating ligands bond to metal forms rings chelate rings Five or six atoms rings are common (i.e. including metal) Coordination numbers and geometries Linear Square planar Tetrahedral Octahedral CHEM261HC/SS1/16 Nomenclature of Coordination Compounds Some common ligands and their names are listed below. The basic protocol in coordination nomenclature is to name the ligands attached to the metal as prefixes before the metal name.

10 Asisthecasewithionic 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. Example 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. Example [Co(NH 2 CH 2 CH 2 NH 2 ) 2 Cl 2 ] + Dichlorobis(ethylenediammine)cobalt(III) )

11 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. Example Exercise 1 Name the following coordination complexes: (i) Cr(NH 3 )Cl 3 (ii) Pt(en)Cl 2 (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)

12 Isomers Primarily in coordination numbers 4 and 6. Different arrangement of ligands in space and also can be the ligands themselves. Types Ionization isomers Isomers can produce different ions in solution e.g. [PtCl 2 (NH 3 ) 4 ]Br 2 [PtBr 2 (NH 3 ) 4 ]Cl 2 Polymerization isomers Same empirical formula or stoichiometry, but different molar mass. Different compounds with similar formula [MX x B b ] n e.g. [Co(NH 3 ) 3 (NO 2 ) 3 ] ( 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. CrCl 3 6H 2 O exist in three different crystalline forms, in which the number of water molecules directly attached to the Cr 3+ ion differs [Cr(H 2 O) 4 Cl 2 ]Cl 2H 2 O [Cr(H 2 O) 5 Cl]Cl 2 H 2 O [Cr(H 2 O) 6 ]Cl 3 In each case, the coordination number of the chromium cation is 6

13 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 nitrito N or O coordination possible (a) [Co(NO 2 )(NH 3 ) 5 ] 2+ (b) [Co(ONO)(NH 3 ) 5 ] 2+ 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. Cl NH 3 H 3 N Cl Pt Pt Cl NH 3 Cl NH 3 cis-[ptcl 2 (NH 3 ) 2 ] trans-[ptcl 2 (NH 3 ) 2 ] CHEM261HC/SS1/19

14 For hexacoordinate systems Purple Green CHEM261HC/SS1/20 For M(X) 3 (Y) 3 systems there is facial and meridional CHEM261HC/SS1/21

15 Example Co Octahedral geometry Cis-[CoCl 2 (NH 3 ) 4 ] + Trans-[CoCl 2 (NH 3 ) 4 ] + Cis/Trans Vs. Fac/Mer Fac- [CoCl 3 (NH 3 ) 3 ] + Mer- [CoCl 3 (NH 3 ) 3 ] + Cis- [CoCl 2 (NH 3 ) 4 ] + Trans- [CoCl 2 (NH 3 ) 4 ] +

16 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 Mn II <Fe II <Co II <Ni II >Cu II >Zn II pm Hard and soft Lewis acid-base theory Properties of hard acids and bases: 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 2 O) 4 (NH 3 ) 2 ] 2+ + en [Cu(H 2 O) 4 (en)] NH 3 CHEM261HC/SS1/24

17 Mainly an entropy effect. Cu(H 2 O) 4 (NH 3 ) 2 ] 2+ + en = [Cu(H 2 O) 4 (en)] 2+ +2NH 3 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 chelate complex has a smaller probability of dissociating. Thus, more stable CHEM261HC/SS1/25 CHEM261HC/SS1/26

18 Metal carbonyl Compounds that have the metal bonded to the carbon monoxide, giving a general formula of M(CO) n M + CO M(CO) n M C O -orbitals in CO are very empty Molecular orbital diagram (CO) What is the bond order? Bond order: No.ofe - pairs in the bonding orbital No. of e - pairs in the anti-bonding orbital

19 Back-bonding (back donation) 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 CO) The C O becomes longer, weaker and the bond order decreases Evidence and extent Infra red (IR) spectra Vibration frequency The greater the extent of back bonding the lower the stretching frequency (bond order decreases) C O Free C O 2143 cm -1 M CO cm -1 Effect of replacing the CO ligands Non- accepting ligands (donor ligands) Cr(CO) 6 3NH 3 Cr(NH 3 ) 3 (CO) cm cm cm cm cm -1 Replacement of the 3 x (CO) groups with donor ligands, 3 x (NH 3 ) 3 increases -acidity of the remaining ligands (CO) so as to counter the accumulation of the negative charge on the metal centre. Metal-carbon (M CO) bond enhanced while carbon-oxygen (C O) bond is weakened, hence, lower wavenumbers on IR spectra.

20 Effect of introducing a positive charge on metal complex V(CO) 6-1 proton 1proton + V(CO) 6 V(CO) cm cm 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 while the C O increases. (wave number or frequency increases) Thought 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? Why? The origin of colour - absorption CHEM261HC/SS1/27

21 Colours on coordination compounds The colour can change depending on a number of factors e.g. Metal charge Ligand Physical phenomenon CHEM261HC/SS1/29

22 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