Transition Metal Coordination Chemistry What is a transition metal?
Prof S.M.Draper
2.05 SNIAMS Building [email protected]
Recommended books
M.J. Winter, d-block Chemistry, Oxford Chemistry Primers, OUP, 2001 M.S. Silberberg, Chemistry, 3rd Ed, McGrawHill, 2003 (chapter 23)
C.E. Housecroft, A.G. Sharpe, Inorganic Chemistry, 1st Ed, PrenticeHall, 2001 J.E. Huheey, E.A. Keiter, R.L. Keiter, Inorganic Chemistry, 4th Ed., HarperCollins, 1993
1 Electrons in atoms z Shapes of d-orbitals 4s 3d 4p y x Sc Ti z z V z y y Cr y Mn x x x Fe
Co
yzxz xy x2-y2 z2 Ni
Cu Zn
2 Working out numbers of d-electrons from oxidation states: How many d-electrons has the metal? 1 2
3 4 5 6 7 8 9 10 11 12 1st: how many electrons are there in the shell? - count along the periodic table O O en = ox = H2N NH2 e.g. Mn = 7 electrons Cu = 11 electrons -OO-
2nd: how many electrons are lost? - oxidation state complex O.S. of L O.S. of M no. d electrons 2- 0 e.g. Mn(VII) = 7 electrons lost Cu(II) = 2 electrons lost [Cr2O7] -2 +6 d - 0 [MnO4] -2 +7 d + 3rd: how many electrons left over? [Ag(NH3)2] - subtract 3+ 1 [Ti(H2O)6] 0+3d e.g. Mn(VII) = Cu(II) = 3+ [Co(en)3] 8 [PtCl2(NH3)2]-1, 0 +2 d 4- [V(CN)6] 3- 5 Hence the only valance electrons available in a transition metal ion are d-electrons [Fe(ox)3] -2 +3 d
3 Colour of transition metal complexes
biological magnetic behaviour activity colour Ruby Corundum
Al2O3 with impurities geometry What’s interesting about Sapphire octahedral metal centre Transition Metal Complexes?? Corundum coordination number 6 Al2O3 with and impurities
coordination Emerald number medical applications Beryl oxidation states AlSiO3 containing Be with impurities
4 Haemoglobin Cisplatin Oxygen carrier in blood O 2 Porphyrin-Fe transition metal complex [PtCl2(NH3)2] Fe(II) ion is octahedrally coordinated Coordination number 6 NN NN NN FeFe NN HO C HO 2C2 NRNR
HO2C HO 2C
square planar Pt(II) coordination number 4 cis-isomer
the first of a series of platinum coordination complex-based anti-cancer drugs (Platinol-AQ)
5 Alfred Werner - Nobel Prizewinner 1913 [Co(NH3)6]Cl3 [Co(NH3)5Cl]Cl2 [Co(NH3)4Cl2]Cl
3+ 2+ +
+ CoCl3 . 6NH3 yellow xs Ag 3 moles AgCl
+ CoCl3 . 5NH3 purple xs Ag 2 moles AgCl Werner's conclusions
+ CoCl3 . 4NH3 green xs Ag 1 mole AgCl 1. The metal is in a particular (primary valancy) CoCl . 3NH xs Ag+ 0 moles AgCl 3 3 2. The complex has a fixed (secondary valancy) 3. The ligands are bound to the metal via a bond which resembles a covalent bond
6 Lewis Acid-Base Concept: parallel with main group What is a coordination complex? chemistry • Adducts formed by the reaction of metal ions (Lewis acids) with several ligands (Lewis bases). n+/-
H 2+ X+/- Ni2+ + H O O : Ni n 2 H . .
NH3 2+ NH3
2+ Ni + 6 :NH3 H3N: Ni :NH3
Central metal ion or atom surrounded by a set of ligands NH3 NH3
The ligand donates two electrons to the d-orbitals around the metal forming a • Properties of products differ from those of separate reactants (i.e., colour, composition, magnetic properties, etc.). 1 3
7 F
F
zz
zz z
z
z z FFB
z z z 3+ z z FFB z e.g. [Co(NH ) ]
HHN z FFB z 3 6 zz zz + zz
z z
HHz N z H F zz HHN H H _ NH3 3+ zz H3N > BF3 NH3 BF3
zz
z z z z z 3+ H3N z z NH
6 HHN z Co 3 L zz + H z z H N z z L L 3 NH3
zz
NH3 L L
L
= coordinate or dative bond
1 CO CO Group 14 carbon monoxide - 2- e.g. NH3, H2O, OH , CO3 - - CO, PPh3, C2H4, SRH, CN , SCN Small donor atoms Larger donor atoms Electronegative H Less electronegative H Not very polarisable NH P PPh Group 15 N 3 3 Easily polarisable ammonia phosphine H
R H H O S SR Group 16 O 2 2 e.g. Fe(III), Mn(II), Cr(III) e.g. Ag(I), Cu(I) water thioether R Small metals (1st row) Larger metals (2nd + 3rd row) H High oxidation state Low oxidation state
2 π - bonded ligands
- z CN - - Group 14 CNz Ph H C CH cyanide 2 2 phenyl Cl CH2 K+ Cl Pt
Cl CH2 RC CR NO- - Group 15 z z NCS NOz nitrous NCS z isocyanate 2 - [PtCl3(η -C2H4)]
Group 16 OH- SCN- z O H hydroxide S C N z thiocyanate
Monodentate donor atom per ligand Bidentate donor atoms per ligand Group 17 X halide H hydride Tridentate three donor atoms per ligand Multidentate many donor atoms per ligand
3 Neutral bidentate ligands: 2 donor atoms Anionic bidentate ligands π-donor bidentate ligand
O O
H2NNH2 oxalate = ox2- zz e.g. [PtCl2(en)] zz - - OO Fe C C C O O 1,2-diaminoethane = ethylene diamine = en O O 4 [Fe(CO)3(η -C4H6)] - H3C acetate = ac O
Ph PPPh 2 zz 2 OH zz NN R zz zz R' NN N O O O zz zz Pd M R' O N C N N C 1,2-diphenylphosphineethane 2,2'-bipyridine 1,10-phenanthroline HO R H H dppe bpy phen Pd(II)-oxime complex
4 Tridentate ligands: three donor atoms Tetradentate ligands Hexadentate ligand
tetraanion of
NH2 ethylenediaminetetraacetic acid N H NNHNH2 zz N 2 zz zz zz EDTA N N zz zz NH2 NH2 O O tris(2-aminoethyl)amine - - diethylenetriamine 2,2':6',2"-terpyridine O O tren - NN - dien tpy O O
N O O HN NH NH N NH N 1,2,4-triazacyclonane zz zz N N macrocyclic ligand N HN N HN
zz N N M H porphyrin phthalocyanin
5 Inner coordination sphere = Outer coordination sphere =
SO 2- H2O 4 H2O
H2O H2O OSO3 2+ H2O Mn OH2 Mn2+
OH2 OH H2O 2 H O 2 H2O
[Mn(OH2)6][SO4]: outer sphere complex [Mn(OH2)5(SO4)5]: inner sphere complex
6 Ligand Exchange Reactions 2+ - + [Fe(OH2)6] +CN [Fe(OH2)5(CN)] + H2O
2+ +
+ +
hexaaquo iron(II) cyanide + complex ion const = [Fe(OH2)5(CN) ][H2O] K1 = [H O] 2+ - 2 [Fe(OH2)6 ][CN]
1 The reaction continues…. The reaction continues….
- + - + [Fe(OH2)3(CN)3 ] [Fe(OH ) (CN)] +CN [Fe(OH ) (CN) ] +HO - - 2 5 2 4 2 2 β3 = [Fe(OH2)4(CN)2]+CN [Fe(OH2)3(CN)3] +H2O 2+ - 3 [Fe(OH2)6 ] [CN ]
K =[Fe(OH) (CN) ] 2- 2 2 4 2 [Fe(OH2)2(CN)4 ] - - 2- [Fe(OH ) (CN) ] +CN [Fe(OH ) (CN) ] +HO β4 = 2 3 3 2 2 4 2 [Fe(OH ) 2+] [CN-]4 + - 2 6 [Fe(OH2)5(CN) ] [CN ]
3- [Fe(OH2)(CN)5 ] 2- - 3- β5 = [Fe(OH2)2(CN)4] +CN [Fe(OH2)(CN)5] +H2O 2+ - 5 [Fe(OH2)6 ] [CN ] + [Fe(OH2)5(CN) ] [Fe(OH2)4(CN)2] K K = β = x 1 2 2 [Fe(CN) 4-] + - 6 2+ - - 4- [Fe(OH ) ][CN] [Fe(OH2)5(CN) ] [CN ] β6 = 2 6 [Fe(OH2)(CN)5]+CN [Fe(CN)6] +H2O 2+ - 6 [Fe(OH2)6 ] [CN ]
[ML ] Overall Stability Constant βn = n [M] [L]n
Stability constants are often expressed in log form ie log βn
2 Values of βn The Chelate Effect
2+ 4-
2+ [M(NH3)6] n+ n+ [M(en) ]2+ + 6 +6 NH3 H2N 3 H NH 2 NH H N 3 N 2 3 M M NH NH H3N 3 N 2 H2 NH3 H2N
4- [Fe(CN)6 ] β = ~ 1035 6 [Fe(OH ) 2+] [CN-]6 2 6 n+ n+ [M(OH2)6] +6 NH3 [M(NH3)6] +6 H2O
[M(OH ) ]n+ + 3 en [M(en) ]n+ +6 HO log β6 = 35 2 6 3 2
Kn < … < K3 < K2 < K1
3 Entropy of chelate formation ΔGo = ΔHo -TΔSo ΔGo = -RT ln K 2+ 2+ [Cu(OH2)6] +2 NH3 [Cu(OH2)4(NH3)2] +2 H2O
β 107.7 log β = 7.7 ΔGo = ΔHo -TΔSo 2 2 ΔHo = - 46 kJ mol-1 ΔSo = - 8.4 J K-1mol-1
Enthalphy changes similar Entropy changes differ
2+ 2+ [Cu(OH2)6] +en [Cu(OH2)4(en)] +2 H2O n+ n+ [M(OH2)6] +6 NH3 [M(NH3)6] +6 H2O
n+ n+ β 1010.6 log β = 10.6 [M(OH2)6] + 3 en [M(en)3] +6 H2O 2 1 ΔHo = - 54 kJ mol-1 ΔSo = + 23 J K-1mol-1
ΔSo is large and positive - TΔSo is large and negative
4 Thermodynamic vs Kinetic Stability The Macrocyclic Effect Complexes containing macrocyclic rings have enhanced stability compared to acyclic ligands
2+ 2+ H H H H N N N N Cu Cu N N N N H2 H2 H H 3+ 3 e.g [Cr(OH2)6] d kinetically inert slow substitution of Ls
3+ 5 cf. [Fe(OH2)6] hs d kinetically labile rapid substitution of Ls log K1 = 23.9 log K1 = 28.0 acyclic chelating ligand macrocyclic ligand (thermodynamic stability similar)
ΔG always favours formation of macrocyclic complex
5 Tour of Coordination Numbers and their Geometries Most Common Geometries of Transition Metal Complexes
The Kepert Model The shape of a complex is usually dictated by the number of coordinated atoms. Tetrahedral 109o 28' C.N. 4 The coordinated atoms are as far away from each other as possible. Non- bonding electrons are ignored. Square Planar 90o C.N. 4
Note: Trigonal bipyramidal 120o + 90o C.N. 5 • regular geometries are often distorted • structural features of multinuclear complexes are described in terms used for individual metal centres Square based pyramidal 90o C.N. 5 • when energy differences between different structures are small, fluxional behaviour may be observed
Octahedral 90o C.N. 6
6 Coordination number 2 Coordination number 5 Cu(I), Ag(I), Au(I), Hg(II)
- - trigonal bipyramid square-based pyramid [CuCl2] [Au(CN)2] 180o
90o 90o
Coordination number 3 120o
- [Cu(CN)2] - [HgI3]
CN CN
Cu Cu C o C N N 120 N N C C The two structures have very similar energy Cu Cu
CN CN n
7 trigonal bipyramid square-based pyramid
3- Cl 3- CN Cl Cl Cu NC CN Cl Co Cl NC CN
3- [Co(CN) ]3- [CuCl5] 5
+ N O N O O NCo V N O O Br
+ [Co(Me6tren)Br] [VO(acac)2]
8 Tetrahedral complexes Favoured by steric requirements
optical isomerism
1 1
2 2 4 4 3 3 2- [CoCl4] - non-superimposable mirror images [MnO4] 2- 90o [NiCl4] 109o
1 Square planar geometry Square Planar Geometry
2- e.g. [PtCl4] - [AuBr4] 2- [Co(CN)4] cisplatin
Square planar complexes are formed by cis-[PtCl (NH ) ] trans-[PtCl2(NH3)2] i.e. group 10 Ni2+, Pd2+, Pt2+ 2 3 2 cis-diamminedichloroplatinum(II) trans-diamminedichloroplatinum(II) Au3+
Square planar complexes are formed by e.g. CN-
2 Coordination number 6 Octahedral geometry Geometrical Isomerism
[ML4X2]
2+ e.g. [Mn(OH2)6]
[Cr(CO)6]
+ + trans-[Co(NH3)4Cl2] cis-[Co(NH3)4Cl2] Staggared ligands vs. Eclipsed ligands green violet
C+ C+ do metals
e.g. WMe6
3 Octahedral geometry Geometrical Isomerism Octahedral geometry Geometrical and Optical Isomerism [ML3X3] 2 + [M(Κ L)2X2] e.g. [Co(en)2Cl2]
trans geometrical isomer Cl H H N N Co N N H H Cl
mer-[Co(NH3)3(NO2)3] cis geometrical isomer two optical isomers
Non-superimposable mirror images = Δ and Λ enantiomers fac-[Co(NH3)3(NO2)3]
4 Octahedral geometry Optical Isomerism Tetragonal Distortion:
Tetragonal elongation z-axis 2+ 2 long axial & 4 shorter 2 N [M(Κ L)3] N equatorial bonds N Ru N 2+ 2+ N [Cu(OH ) ] , [Cr(OH ) ] N 2 6 2 6
Tetragonal compression z-axis
2 short axial & 4 longer equatorial bonds M + trans-[Co(NH3)4(CN)2] 2+ [Ru(bpy)3]
* Small distortion but symmetry lower than octahedral.
5 Isomerism Trigonal Distortion:
• Two faces on opposite sides of octahedron move either towards Structural Isomerism Stereoisomerism or away from each other Isomers have the same empirical Isomers have the same M-L bonds, formula but the atoms connectivities but the atoms are arranged differently N differ in space N N Elongation Rotate M by 60° N N N Ionisation isomerism Geometrical isomerism Hydration isomerism cis/trans 2+ Complexes of tacn can show this type of distortion [M(tacn)2] Coordination isomerism mer/fac Linkage isomerism Optical isomerism N tacn = 1,3,7-triazacyclononane Polymerisation isomerism Δ and Λ N N
6 Ionisation isomerism Coordination isomerism Exchange of a ligated anion with a counterion
[Cu(NH3)4][PtCl4][Pt(NH3)4][CuCl4] square planar
2+ Ba BaSO no ppt [Co(NH3)5Br]SO4 4 [Co(NH3)6][Cr(CN)6] [Cr(NH3)6][Co(CN)6] octahedral + Ag Ba2+ AgBr [Co(NH3)5(SO4)]Br no ppt
Polymerisation isomerism
Solvate isomerism
Exchange of a neutral ligand for an anionic ligand e.g. [Pt(NH3)4][PtCl4][Pt(NH3)2Cl2]
e.g. [Co(OH2)6]Cl3 violet Both polymers have the empirical formula [Pt(NH3)2Cl2]n [Co(OH2)5Cl]Cl2 light green
[Co(OH2)4Cl2]Cl dark green
7 Linkage isomerism Higher Coordination Numbers
Coordination Number 7 hν 2+ 2+ [Co(NH3)5(NO2)] [Co(NH3)5(NO2)] Δ yellow red nitro-complex nitrito-complex
- 2- [WBr3(CO)4)] [TaF7] [Pd(NCS)2(PPh3)2] [Pd(SCN)2(PPh3)2] (distorted) [ZrF ]3- isocyanate thiocyanate 7
8 Coordination Number 8
n Na3[Mo(CN)8](Bu4N)3[Mo(CN)8]
Coordination Number 9
2- [ReH9]
9 Symmetrical ligand field z Shapes of d-orbitals orthogonal
y z
x M x and y z z z y y y
x x x x2-y2 yz z2 xz xy
Energy yzxz xy x2-y2 z2
x2-y2 yz z2 xz xy
1 Effect of ligand field on metal d-orbitals The difference in energy between the eg and the t2g energy levels is the . This is given the value 10 Dq.
symmetrical field octahedral ligand field
eg
Δo or 10 Dq
t2g
2 1 Electron configurations: d ion 3+ The nature of Δoct
490-580 nm 3+ [Ti(OH ) ]3+ e.g. [Ti(OH2)6] 2 6 violet solution in water
one d-electron in a t2g orbital
the complex has a Crystal Field Stabilisation Energy (CFSE) of - 0.4 Δoct
-1 Absorption spectrum: λmax = 510 nm = 243 kJ mol
3 d4 ions
3+ e.g. [V(OH2)6]
+ 0.6 Δoct CFSE =
-0.4 Δoct = - 0.6 Δoct
+ 0.6 Δoct e.g. [Cr(OH ) ]3+ 2 6 CFSE =
-0.4 Δoct = - 1.6 Δoct + P
4 High and Low spin Complexes Factors affecting the magnitude of Δoct
High spin complex Low spin complex
2+ -1 3+ -1 e.g. [Fe(OH2)6] Δoct = 10 000 cm [Fe(OH2)6] Δoct = 14 000 cm 2+ -1 3+ -1 [Co(OH2)6] Δoct = 9 700 cm [Co(OH2)6] Δoct = 18 000 cm
Δ is small Δ is large
3+ -1 e.g. [Co(NH3)6] Δoct = 22 900 cm electrons occupy e and t orbitals electrons pair in t oribtals before g 2g 2g 3+ -1 [Rh(NH3)6] Δoct = 34 100 cm singly before pairing occupying e orbitals g 3+ -1 [Ir(NH3)6] Δoct = 41 000 cm
5 eg
eg
- - 2------2- I < Br < S < SCN ≈ Cl < NO3 < F < OH < ox
< H O < NCS- < CH CN < NH ≈ en < bpy t2g 2 3 3
- - < phen ≈ NO2 < PR3 < CN ≈ CO
t2g
M.J. Winter, d-block Chemistry, p.83
6 Octahedral Crystal Field ions, Oh field
d-orbital energies in…
+ 0.6 ∆oct free ion spherical ligand field octahedral ligand field
-0.4 ∆oct
x2-y2 yz z2 xz xy
+ 0.6 ∆oct
x2-y2 yz z2 xz xy -0.4 ∆oct
the complex has a Crystal Field Stabilisation Energy (CFSE)
ions, O field h ions, Oh field
High Spin
eg eg + 0.6 ∆oct + 0.6 ∆oct -0.4 ∆oct t2g -0.4 ∆oct t2g
Low Spin
eg
+ 0.6 ∆oct + 0.6 ∆oct
-0.4 ∆oct -0.4 ∆oct t2g
1 ions, O field h ions, O field 3- h eg 1. What is the CFSE of [Fe(CN)6] ? e g eg + 0.6 ∆ oct + 0.6 ∆oct -0.4 ∆ t oct -0.4 ∆oct 2g t2g 2+ eg. 2. If the CFSE of [Co(H2O)6] is -0.8 ∆oct,
what spin state is it in?
ions, Oh field
+ 0.6 ∆oct Only configurations can be high or low spin -0.4 ∆oct t2g
Jahn-Teller Distortion Jahn-Teller Distortion z z-out: d-orbitals with z-component are stabilised y O field x h
1 + /2 δ1 eg
1 - /2 δ1
∆oct
2 + /3 δ2
t2g 1 - /3 δ2 The ligands contract or elongate giving a complex
2 Jahn-Teller Distortion Jahn-Teller Theorum z-in: d-orbitals with z-component are destabilised "For a nonlinear molecule in an electronically degenerate state distortion must occur to lower the symmetry, remove the degeneracy and lower the energy" Oh field
x2-y2 1 e + /2 δ1 g x2-y2 z2 eg e.g. K2[CuF4]Oh z-in x2-y2 z2 z2 - 1/ δ 2 1 Na2[CuF4]Oh z-out xy ∆oct t2g Cr2F5 Oh z-out yz xz xy Cr(II) = high spin d4 + 1/ δ 3 2 yz xz t2g yz xz xy 2 2+ - /3 δ2 e.g. [Cu(OH2)6] Cu(II) = d9 z-out
Square planar geometry: ML4
Oh z-out square planar
x2-y2
x2-y2 eg 2 2 2 x -y z z2
xy
z2 xy t2g yz xz xy yz xz
yz xz
d8 complexes: some Ni(II), and Pd(II), Pt(II)
3 Tetrahedral crystal field splitting Splitting of d-orbitals in a tetrahedral crystal field
eg
Δoct
metal ion in symmetrical field in free space
t2g
x2-y2 yz z2 xz xy
1 Effect of crystal fields on complex geometry Spin orbit coupling is particularly significant for metals lower in the periodic table but for 1st row transition metal complexes the formula can be used in a simplified form because:
“The spin contribution outweighs the orbital angular momentum contribution to μeff"
Spin only magnetic moment, μSO
μSO = n (n + 2)
where n = number of unpaired electrons
μSO = 2 S (S + 1)
where S = total spin quantum number = n / 2
2 Spin only magnetic moment, μSO Measurement of magnetic susceptibility where n = number of unpaired electrons μSO = n (n + 2) Gouy Balance
What is the spin only magnetic moment of …..? eg x2-y2 z2 2+ [Co(OH2)6] pivotal balance
t2g yz xz xy n = μSO = = scale
t2 yz xz xy [NiCl ]2- sample 4 e x2-y2 z2 n = μ = = SO N S magnetic field x2-y2
2- [Ni(CN)4] xy
z2 n = μSO = = yz xz
3 Values of μeff for high spin Oh complexes
Effective magnetic moment, μeff
n M ion d S μSO(BM) μobs (BM) From experimentally measured molar magnetic susceptibility, χm Sc3+, Ti4+ d0 00 0
3+ 11 Ti d /2 1.73 1.7 - 1.8 V3+ d2 1 2.83 2.8 - 3.1
2+ 3+ 33 V , Cr d /2 3.87 3.7 - 3.9 Cr2+, Mn3+ d4 2 4.90 4.8 - 4.9
2+ 3+ 55 Mn , Fe d /2 5.92 5.7 - 6.0 Fe2+, Co3+ d6 2 4.90 5.0 - 5.6
2+ 73 Co d /2 3.87 4.3 - 5.2 μeff in Bohr Magnetons: 2.83 is the magnetogyric ratio Ni2+ d8 1 2.83 2.9 - 3.9 -24 -1 μB = e h = 9.27 x 10 J T T is temperature in Kelvins Cu2+ d91/ 1.73 1.9 - 2.1 4 π me 2 Zn2+ d10 00 0
4 Colour in TM complexes Colour of d-d transitions depends on magnitude of Δ
3+ [Ti(OH2)6]
Absorption spectrum: λmax = 510 nm
490-580 nm The Colour Wheel white light 400-800 nm If red light is absorbed the complex appears green Blue: 400-490 nm yellow-green: 490-580 nm If purple light is absorbed Red: 580-800 nm the complex appears yellow wavelength, λ (nm)
5 Effect of magnitude of Δ on colour Effect of magnitude of Δ on colour
1. For a given ligand, the colour 2. For a given metal ion, the colour depends on 3+ 2+ [V(H2O)6] [V(H2O)6]
violet light absorbed yellow light absorbed complex appears yellow complex appears violet
3+ 2+
6 Colour and the Spectrochemical Series
Weak Field Ligands Strong Field Ligands High Spin Complexes Low Spin Complexes
- - 2------2- I < Br < S < SCN < Cl < NO3 < F < OH < ox - small Δ < H2O < NCS < CH3CN < NH3 < en < bpy large Δ - - < phen < NO2 < phosph < CN < CO
7 [Ti(H O) ]3+ Absorption spectrum: λ = 510 nm Transition Metal Coordination Chemistry 2 6 max
490-580 nm Professor Sylvia Draper
eg eg [email protected] hν ∆o WHERE HAVE WE BEEN ? WHERE ARE WE GOING ? t2g t2g Lecture 7: Magnetism and Colour Lecture 8: Tetrahedral crystal field Colour continued, and Ligand Field Magnetism Spin only magnetic moment Theory Magnetism measurements Selection rules Absorption of light Factors affecting magnitude of ∆ d1 complexes and colour Limitations of CFT Effect of ∆ on colour LFSE and ∆ 1. Oxidation state of metal e.g. Thermodynamic aspects of LFT 2. Position of metal in periodic table
3. Type of ligand in spectrochemical series
1 Intensity of colour depends on ….. The Intensity of the colour depends on Selection Rules and is measured by the absorbance
Absorbance is defined by the Beer-Lambert equation: A = log10(I0/I) = ε.c.l
where A = absorbance
I0 = intensity of incident light; I = intensity of light after passing through the cell; ε = molar absorption coefficient; c = concentration; l = cell pathlength.
Typical Spectrum: Octahedral Cr(III) complex 0.25 [Cr] = 0.002M ? 0.2
Calculation of ε: At 590 nm, A = 0.138 0.15 1. Laporte Selection Rule A = εcl 0.1
0.138 = ε x 0.002M x 1 cm Absorbance 2. Spin Selection Rule ε = 0.138/ 0.002 x 1 0.05 -1 -1 ε = 69 M cm 0 300 350 400 450 500 550 600 650 700 Wavelength (nm) Transitions allowed by the selection rules (high ε values > 1000 M-1cm-1).
2 Laporte Selection Rule Transitions between d-orbitals (g) forbidden by the Laporte selection rule
Octahedral complexes
Centrosymmetric: t2g and eg orbitals
Selection rule lifted by molecular vibration
Tetrahedral complexes s-orbital d-orbital p-orbital gerade gerade ungerade non-Centrosymmetric: t2 and e orbitals
e.g. p-orbital to d-orbital by the Laporte selection rule Tetrahedral complexes are usually more strongly coloured than analogous d-orbital to d-orbital by the Laporte selection rule octahedral complexes
3 The Spin Selection Rule Assumptions made in CFT
1. The ligands are 2. Interactions are purely
Explains 1. Geometry 2. Magnetism 3. Colour
Limitations Does not allow for in M-L bonding
Doesn't explain Order of ligands in the Spectrochemical Series Nephelauxetic Effect
4 The Spectrochemical Series The Nephelauxetic Effect
Weak Field Ligands Strong Field Ligands High Spin Complexes Low Spin Complexes "there is in complexes than in the free ion"
- some in M-L bonds – M and L share electrons - - 2------2- - I < Br < S < SCN < Cl < NO3 < F < OH < ox < H2O < NCS < CH3CN - effective size of metal orbitals
- - - electron-electron repulsion < NH3 < en < bpy < phen < NO2 < phosph < CN < CO
Nephelauxetic series of ligands
- 2- - - - - F < H2O < NH3 < en < [oxalate] < [NCS] < Cl < Br < I CFT does not explain why some ligands with charges are Nephelauxetic series of metal ions than analogous ligands which are not charged Mn(II) < Ni(II) Co(II) < Mo(II) > Re (IV) < Fe(III) < Ir(III) < Co(III) < Mn(IV)
5 LFSE:- the EXTRA stabilisation of complexes that occurs as a result of the splitting of orbitals in a ligand field (cf. to CFSE: it is always positive and no account of P is taken)
LFSE Octahedral (always greater than tetrahedral) high/low spin options d4 to d7 only
1 2 3 d d d d1 : LFSE 0.4 eg 3/5∆o d2 : LFSE 0.4 X 2 = 0.8 -2/5∆O t2g d3 : LFSE 0.4 x 3 = 1.2 maximum d4 hs : LFSE 0.4 x 3 – 0.6 = 0.6 d4 ls : LFSE 0.4 x 4 = 1.6 d4 d4 d5 d5 eg d5 hs : LFSE 0.4 x 3 – 0.6 x 2 = 0 minimum 3/5∆o d5 ls : LFSE 0.4 x 5 = 2.0 highest -2/5∆O t2g LFT is identical to CFT, except that it allows for some in M-L bonding. d6 hs : LFSE 0.4 x 4 - 0.6 x 2 = 0.4 d6 ls : LFSE 0.4 x 6 = 2.4 d6 d6 d7 d7 7 LFT is based on parameters. eg d hs : LFSE 0.4 x 5 – 0.6 x 2 = 0.8 3/5∆o d7 ls : LFSE 0.4 x 6 – 0.6 = 1.8 -2/5∆O t2g d8 : LFSE 0.4 x 6 – 0.6 x 2 = 1.2 maximum
ls gives higher LFSE than hs option
6 Variation in LFSE with ∆excluding P) Lattice Energies (from Born Haber cycle)
For 1st row M(II) chlorides: hs, Oh
Ca(II) Ti(II) V(II) Cr(II)….etc
The energy when ions come together from infinite ls d6: LFSE separation to form a crystal hs d0, d5, d10 :LFSE Mn+ + n X- MX 3 8 (g) (g) n(s) hs Oh d , d :LFSE
Discrepancies in lattice energies are a result of
7 Variation in Hydration Enthalpies
Enthalpy of hydration:
2+ 2+ M (g) + xs H2O [M(OH2)6] (aq)
8 Variation in Ionic Radii
Ca2+ Zn2+ high spin
low spin Last Lecture 9 lecture course on Coordination Chemistry Prof. Sylvia Draper Sc3+ Ga3+ high spin
So……. CFSE, LFE and MO theory agree ! low spin - each one building upon and improving the other
6 1 low spin: steady followed by increase from t2g eg 3 1 high spin: steady followed by increase from t2g eg
1 Ionisation Energies of first row TM ions, hs, O h Irving-Williams series
2+ 4- 2- [M(H2O)6] +[EDTA] [M(EDTA)] +6H2O
Position of equilibrium depends on
d5 d6 d7 d8 d9 d10
Mn2+ < Fe2+ < Co2+ < Ni2+ < Cu2+ > Zn2+
2 Valance Bond Theory 2- 2- Tetrahedral e.g. [Zn(OH)4] Square Planar e.g. [Ni(CN)4]
Ligand = Lewis base Metal = Lewis acid s, p and d orbitals give with specific geometries Number and type of M-L hybrid orbitals determines geometry of the complex
Octahedral Complex
3+ e.g. [Cr(NH3)6] Limitations of VB theory
Assumes bonding is Does not account for Can predict wrongly Does not account for
3 Molecular Orbital (MO) theory Ligand Group Orbitals MO theory considers covalent interactions e.g. molecular hydrogen Metal valance shell orbitals = Linear Combination of Atomic Orbitals (LCAO) Ligand valance shell orbitals =
LCAO approach 3+ e.g. [Co(NH3)6] = sigma-bonded complex
1s 2s 2p 1s 1s zz sp3 hybrid orbitals N H H H
3 H atomH2 molecule H atom Six sp hybrid ligand orbitals form a set of
4 3- Complexes with M-L π-bonding e.g. [CoF6] 3+ - t1u* Co = 6 e - - 6 x F = 12 e - - - - e.g. H2O,OH , lower halides Cl , Br , I a * 1g Total = π-bond σ-bond high spin electron density to the metal centre π-bond ligand orbital and metal orbital 4p
4s of electron density from filled p orbitals
eg*
t 3d 2g π-acceptor ligands e.g. CO, N2, NO, alkenes 6 x LGO
eg electron density from the metal centre OC C O
t1u metal orbital and ligand orbitals L
L L a1g L L n+ n+ ML ∗ M 6 L of electron density into emtpy π -antibonding orbitals
5 π-donor ligands: filled ligand orbitals and empty metal orbitals π-acceptor ligands: back-donation to empty ligand antibonding orbitals
t1u* t1u*
small ∆oct large ∆oct - weak field ligand a1g* - strong field ligand a1g*
Ligand π∗-orbitals eg* 4p 4p (vacant) eg* 4s Ligand π-orbitals 4s (occupied) 3d 3d
Ligand σ-orbitals Ligand σ-orbitals (occupied) (occupied)
eg Metal d eg Metal d electrons electrons t t (d6 example) 1u (d6 example) 1u
a1g a1g
n+ n+ n+ n+ Energy M ML6 6L Energy M ML6 6L
6