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Transistion review

by: Rachael Chase

Transistion review CHEM 111 - 02

Rachael Chase

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summaries of metals
Principles of Chemistry I
Dr. Wang
Class Notes
inorganic, Chem
25 ?




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This 12 page Class Notes was uploaded by Rachael Chase on Monday July 11, 2016. The Class Notes belongs to CHEM 111 - 02 at Georgia State University taught by Dr. Wang in Summer 2016. Since its upload, it has received 3 views. For similar materials see Principles of Chemistry I in Chemistry at Georgia State University.


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Date Created: 07/11/16
MO Diagrams for O and N 2 2 Drawing MO diagrams - always same number of MOs as AOs. - more electronegative element on the right - lower in energy. ------------------------------------------------------------------------------------------------------------------ Paramagnetic O 2 σ*2p Bond order of 2 (O=O) π*2p 2p 2p π 2p larger ΔE σ2p σ*2s 2s 2s σ 2s O O2 O ------------------------------------------------------------------------------------------------------------------ σ*2p Diamagnetic N 2 Bond order of 3 (N≡N) π*2p 2p Very strong NN bond σ 2p 2p Unlike O2, the 2pand π 2pbitals are smaller ΔE switched in energy (see next page) π2p HOMO = σ , 2pMO = π* 2p σ* 2s 2s 2s σ2s N N2 N MO Diagram for CO σ*2p LUMO 2p π*2p σ 2p 2p HOMO (slightly antibonding) π2p 2s σ* 2s 2s σ 2s C CO O - CO is isoelectronic with N (C has one less electron than N, O has one more) – MO diagram is 2 very similar to that of N2. - Diamagnetic, Bond order of 3 (C≡O) - Important: HOMO = σ (wea2py antibonding), LUMO = π* (strongly an2pbonding) - Since O is far more electronegative than C, would expect a large dipole moment with δ- on O. However, CO actually has only a small dipole moment (0.1 Debye) with the δ- on Carbon! Usually electronegativities are a good indication of the direction and magnitude of the dipole on a molecule, but especially in molecules where orbitals with antibonding character are occupied, things are not so straightforward. ------------------------------------------------------------------------------------------------------------------ Ordering of MOs in O 2is σ2p π 2pπ* ,2p* bu2pin N and 2O, the order is π , σ , 2p ,2p* 2p 2p Reason - O -2arger ΔE between s and p orbitals Æ less orbital mixing - larger ΔE because across the pe riod (B,C,N,O,F) more protons added to the nucleus, which pulls the electrons in clos er to the nucleus (lower energy). This effect is felt more strongly by the s-orbita ls than the p-orbitals, so the energy of the s-orbitals drops more rapidly than that of the p-orbitals. - N2 and CO - mixing between the σ and t2p σ* orbitals2saises the energy of the σ 2p above the π . 2p Crystal Field Theory • Very wrong theory (for transition metal organometallic chemistry), but quite useful. • Easy way to roughly predict the ordering of the energy levels for a TM compound. - Ligand lone pair thought of as a point charge (e.g. for Cl -, H or CN ) or the partial negative charge of a dipole (e.g. for :OH , :NH or :PF ). 2 3 3 - Ligands attracted to the TM catio n since they are point charges Æ purely electrostatic bonding assumed. - Any interactions between the electrons of the ligand and those of the metal are repulsive Æ no covalent bonding (σ, π or δ) exists according to this theory!!! dz2 dx2-y2 axial, g Δ o TM d-orbitals inter-axial, t 2g dxz,dyz,dxy - The axial orbitals (e gsymmetry in an octahedral enviro nment) point directly towards the ligands so according to crysta l field theory, are raised in energy (unfavourable interaction between the electrons of the ligands and those of the metal) - The inter-axial orbitals (t 2g symmetry in an octahedral en vironment) do not point directly towards the ligands, so are lower in energy than the axial orbitals. Crystal Field Theory Square pyramidal and Square Planar Geometry from Octahedral d x2-y2 d , d d z2 x2-y2 x2-y2 large ΔE dz2 d xy Δo d z2 dxy TM d-orbitals dxz, yz dxz,dyz,dxy d d xz, yz M M M octahedral square pyramidal square planar - From the octahedral crystal field splitting di agram, pull off one of the ligands along the z-axis. - Since this partly removes an unfavourable ( according to CF theory) interaction in the z-direction, all orbitals with a z-component (d z2 dxznd d )yzall in energy. The others go up in energy. - Removing the other ligand on the z-axis has the same effect, leading to the energy level diagram for a square planar complex. Note the large energy gap between the d xy and d x2-y2 orbitals. As a result, square planar comp lexes are almost excl usively those of d metals (all metal orbitals are filled except the x2-y2 and so have 16 electrons rather than 18. See later. - Crystal field theory is therefore quite useful for quickly predicting the order of the energy levels in transition metal complexes. Howeve r, the assumption that all bonding is ionic and no covalent interactions occur is unrealistic. According to th is assumption, CO (an excellent ligand for reasonably electron rich metals) would be predicted to be a terrible ligand (no negative charge and v. small dipole), ligands would not be expected to bond to metals that do not have a formal positive charge (i.e. Pd(PPh 3 4and Fe(CO) 42-would definitely not exist), ligands cannot be σ-donors, π-donors or π-acceptors, and it is therefore impossible to assess why Δ vories with the ligands as it does. Crystal Field Splitting Diagram for Tetrahedral Geometry OCTAHEDRAL TETRAHEDRAL Inter-axial orbitals point more Axial orbitals point less directly at the 4 ligands Æ more of an directly at the 4 ligands Æ less of an unfavourable interaction between the unfavourable interaction between the ligand electrons and those of the metal ligand electrons and those of the metal Æ higher in energy Æ lower in energy t2 d orbitals Δ ~ 4/9 Δ Δtis much smaller than Δ paotly t o because none of the metal orbitals point directly at the ligands, so overlap is less efficient. e M [ML 4] The same type of approach can be extended in order to figure out the splitting pattern of the d-orbitals for a variety of other common geomet ries. Starting from octahedral, one can derive Jahn-Teller distorted octahedral geometry, square pyramidal and square planar geometry. Byplacing the ligands either in the 8 corners or on the 12 edges of a cube, one can derive the crystal field splitting diagram for 8-coordinate c ubic or 12-coordinate cube-octahedral geometry (similar approach to that used for tetrahedral ge ometry). The splitting pattern for trigonal planar geometry can be derived by considering which orbitals or groups of orbitals point most directly towards the ligands. A similar approach can be us ed for trigonal bipyramidal geometry, although the actual ordering of the orbitals is more difficult to rationalise using such simple arguments. MO theory ~ Ligand Field Theory Constructing an MO diagram for an octahedral complex Symmetry Adapted combinations of Ligand σ-orbitals in an Octahedral Complex st - In an octahedral environment, the metal orbitals (3d, 4s, 4p for a 1 row TM) divide by symmetry into 4 sets: s = a 1g, p = 1u, axial d = g , inter-axial d =2g . - The orbitals of the six ligands can be comb ined to give six symmetry-adapted linear combinations which are of the correct symmetry to interact with the s, 3 x p and 2 x axial-d orbitals, but not the inter-axial d orbitals. - The result is that 3 orbitals (the inter-axial d-orbitals) are non-bonding, while the rest (6 metal orbitals and 6 ligand orbitals ) combine to form six bonding and six anti-bonding MOs. See next page. - This is a much more correct approach than crystal field theory, but is not as easy to use. MO energy levels for an octahedral complex (only σ-bonding considered) 1u* 4p (1u) a * 1g 4s (a ) 1g e g ΔO 3d (eg+ t2g t2g (a1g + 1u + eg) eg TM 6 L t1u a 1g • The six bonding orbitals are filled with 12 electrons from the six ligands • Orbitals shown in red (t 2g and eg*) are the frontier orbitals where d-electrons reside. π-bonding ligands (the above MO diagram does not take into account π-bonding) filled orbital empty orbital M L M L M L σ-donor π-donor π-acceptor - - - - - + - NH 3, CH3, H Cl , OH , N2 , OH2 CO, NO , CN • For π-acceptor ligands, the bonding is SYNERGIC: σ-donation to the metal strengthens π-backbonding to the ligand, and π-donation from the metal to the ligand strengthens the σ-donor component of bonding. • This is because σ-donation leads to increased electron density on the metal, which allows increased π-backdonation. Conversely, π-backdonation reduces the amount of electron density on the metal, which allows more σ-donation from the ligand to the metal. Cr(CO) : Oc6ahedral complex with good π-acceptor ligands σ-interaction π-interaction C M O M C O lobe of slightly antibonding acceptor orbital HOMO of CO (whole orbital not shown) M C O axial d-orbitals strongly antibonding π* orbitals of CO Note: The 12 empty π*-orbitals of the six CO ligands in a molecule like [Cr(CO) 6] can be combined to form 12 linear combinations of orbitals (3 x T , 3 x T , 3 x T , 3 x T ). Only the 1u 2g 1g 2u three linear combinations with T2g symmetry are of the correct sy mmetry to interact with the 2g orbitals (xy, xzand dyz on the metal. MO energy levels for an octahedral complex with π-acceptor ligands (e.g. [Cr(CO) ]) 6 t1u 4p (1u) a * 1g 4s (a1g t2g π* orbitals of CO 2g) eg* large Δ 3d (eg+ t2g O 2g (a1g+ t1u+ e g e g Cr 6 CO 1u a1g • π-backdonation to CO from the t 2g orbitals (which are non -bonding in the absence of π-interactions between the metal and the ligands). • The 3 t2g orbitals and 3 high lying π* orbitals of the CO ligands form 3 bonding MOs and 3 antibonding MOs. • Since the CO π* orbitals are empty, the d-electrons occupy the bonding MO from this interaction. • The result is (1) a very large o , so the g orbital is likely to remain empty. (2) the t orbital is strongly bonding (wants to be filled with 6 electrons) 2g Æ complexes of strong π-acceptor ligands are the most likely to obey the 18 electron rule Simplified picture of how π-acceptor and π-donor interactions affect the MO diagram - Only the frontier orbitals are shown 2g* π* (2g) e g eg* L large Δ O π-acceptor ligands 2g increase Δ O TM complex with σ-bo nding oly t2g [Cr(CO)6] with π-backdonation from Cr to CO π-donor ligands eg* eg* decrease Δ small O Δ O t2g* t2g TM complex with t 2g L σ-bonding only TM complex withπ- donation from LigandtoMetal π-donor ligands • π-donation from the ligands to the t 2gbitals Æ the 3 t m2gal orbitals and 3 low lying, filled ligand orbitals of π-symmetry form 3 bonding MOs and 3 antibonding MOs. • Since the interacting ligand orbitals are full, these electrons occupy the bonding MO from this interaction, and the d-electrons occupy the antibonding MO. • The result is (1) a small Δo (2) the 2g orbital is weakly antibonding The Spectrochemical Series of Ligands Using MO theory (Ligand Field Theory) instead of Crystal Field Theory, we have seen that ligands can be σ-donors, π-donors or π -acceptors, and that the nature of the ligands strongly affects Δ in octahedral complexes. o This effect can often be seen visually: e.g. for [Co IIX(NH ) 3 5a d , 18-electron complex). Colour of eg Colour X n the absorbed complex Δ - o I 2 Purple Yellow Cl 2 Pink Green d t2g NH 33Yellow Blue • The colo ur of the complexes above results from promotion of an electron from the t 2g orbital to the e gorbital. The energy of light absorbed therefore correspond s to the size of Δo. 1 9 [Note: this type of simple treatmen t can only be applied in certain cases ( e.g. d or d complexes and octahedral 3d complexes with a d 4or d 6 configuration). For other configurations, it is necessary to consult a Tanabe-Sugano diagram and apply a more rigorous treatment]. • For complexes with a single absorption in the vi sible region of the spectrum, the colour of light absorbed can be determined from the colour wheel (thecolour of the light absorbed is found opposite the colour of the complex) • Since the order of energy is blue > green > yellow, then the NH 3 complex can be seen to have a larger Δ ohan the Cl complex, which has a larger Δ than tho I complex.- Since the effect of the ligands on the size of Δ is o visible change, the resulting series of ligands is called the spectrochemical series. The order of ligands in the spectrochemical series follows their behaviour as π-donors, σ-donors or π-acceptors very nicely. For Δ : o I-< Br < S < SCN < Cl < F < OH < OH < MeCN < NH 2 PR < CN < CO < NO3 3 - + good π-donors │ OK π-donors │ σ-donors │ good π-acceptors However, these π-affects are not the whole story, because organometallic ligands such as CH 3 - - - - or H are located high in the spectrochemical series (H is similar to CO). Both H and CH 3 (in the absence of α-agostic interactions) are purely σ-donor ligands, so their ability to act as high field ligands is a result of their extremely high σ-donor ability. The Spectrochemical Series of Metals Δ oepends not only on the nature of the ligands, but also on the metal and its oxidation state. The spectrochemical series for metal ions (approximate) is shown below: 2+ 2+ 2+ 2+ 2+ 3+ 4+ 3+ 3+ 3+ 4+ 3+ 4+ Mn < Ni < Co < Fe < V < Co < Mn < Mo < Rh < Ru < Pd < Ir < Pt This series is not quite as regular as the spec trochemical series of ligands, but there are obvious trends: 2+ 3+ (1) Δoincreases with increasing oxidation number (Co < Co etc.) Reason- Metal ions in a higher oxidati on state have greater polarising power Æ bonding becomes more covalent and less ionic (the nephelauxetic effect = Latin for cloud expanding). 3+ 3+ 3+ (2) Δo increases down a group (Co << Rh < Ir ) Reasons - Improved M-L bonding for larger 4d and 5d orbitals relative to the 3d orbitals. - Higher effective nuclear charge ( i.e. the actual charge on the metal in a Rh(III) complex will be higher than in an an alogous Co(III) complex because heavier congeners are easier to oxidise).


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