Chemistry of Transition Metals Bonding in transition metal compounds - PowerPoint PPT Presentation

Chemistry of Transition Metals

Bonding in transition metal compounds Theories : (i) Werner Coordination Theory (ii) 18 electron rule/ EAN (iii) Valence Bond Theory (iv) Crystal field theory (v) Molecular orbital approach Consequences : (i) High spin - low spin complexes (ii) Spectrochemical series (iii) Crystal Field Stabilization Energy (CFSE) (iv) Jahn-Teller distortions (v) Spinels

Werner Coordination Theory • Cobalt(III) complexes: CoCl 2 + NH 3 (aq) and then oxidized by air. CoCl 3 6 NH 3 orange-yellow CoCl 3 5 NH 3 H 2 O red CoCl 3 5 NH 3 purple CoCl 3 4 NH 3 green •  NH 4+(aq) + Cl – (aq) NH 3 (aq) + HCl (aq) However, CoCl 3 6 NH 3( aq ) + HCl ( aq )  no reaction • Ag + (aq) + Cl – (aq)  AgCl (s) ( white )  + Ag + Expected: CoCl 3 6 NH 3 3 AgCl (s)  + Ag + Observed: CoCl 3 6 NH 3 3 AgCl (s)  + Ag + Expected: CoCl 3 5 NH 3 3 AgCl (s)  + Ag + Observed: CoCl 3 5 NH 3 2 AgCl (s)  + Ag + Expected: CoCl 3 4 NH 3 3 AgCl (s)  + Ag + Observed: CoCl 3 4 NH 3 1 AgCl (s)

• Measurements of the conductivity : CoCl 3 6 NH 3 four ions CoCl 3 5 NH 3 H 2 O four ions CoCl 3 5 NH 3 three ions CoCl 3 4 NH 3 two ions • The secondary valence is the number of ions of molecules that are coordinated to the metal ion. Werner assumed that the secondary valence of the transition metal in these cobalt(III) complexes is six. The formulas of these compounds can therefore be written as follows. [Co(NH 3 ) 6 ] 3+ 3Cl – orange-yellow [Co(NH 3 ) 5 (H 2 O)] 3+ 3Cl – red [Co(NH 3 ) 5 Cl] 2+ 2Cl – purple [Co(NH 3 ) 4 Cl 2 ] + Cl – green Werner : Nobel prize in 1913.

An example of a metal complex Metal ion: Central & is a Lewis Acid Ligand: Is Attached & is a Lewis Base

How strong are the Co-ordinate bond? Prussian blue CH 3 CN Fe 4 [Fe(CN) 6 ] 3 · x H 2 O. Prussian blue is used for certain heavy metal poisons. Modest toxicity. Can be metabolised to produce HCN, which is very toxic KCN is a poison by itself....

[Co(NH 3 ) 6 ]Cl 3 and [CoCl(NH 3 ) 5 ]Cl 2 Two compounds made of the same chemicals , yet they look different , and react differently with silver(I) salt .

Colors & How We perceive it Black and White When a sample absorbs light, what we see is the sum of the remaining colors that strikes our eyes. If a sample absorbs all wavelength If the sample absorbs no of visible light, none reaches our visible light, it is white eyes from that sample. or colorless. Consequently, it appears black.

Absorption and reflection If the sample absorbs all but orange, the sample appears orange. Further, we also perceive orange color when visible light of all colors except blue strikes our eyes. In a complementary fashion, if the sample absorbed only orange, it would appear blue; blue and orange are said to be complementary colors.

Examples: Number Geometry Polyhedron Linear (2) [Ag(NH 3 ) 2 ] + Trigonal plane (3) [HgI 3 ] - Square planar (4) [Ni(CN) 4 ] 2- Tetrahedral (4) [Zn(NH 3 ) 4 ] 2+ , [FeCl 4 ] -

Coordination No. 5 Geometry Polyhedron Number Square pyramid (5) [Ni(CN) 5 ] 3- Example: Trigonal bipyramid (5) Example: [Cu(Cl) 5 ] 3-

Coordination No. 6 Octahedral (6) Triagonal prism

Coordination No. 7 Geometry Polyhedron Number Singly capped octahedron (7) [Mo(CN) 7 ] 2-

Pentagonal bipyramidal (7) [Re(CN) 7 ] 2-

Coordination No. 8 Doubly capped octahedral (8) [W(CN) 8 ] 3-

Ligands Ligands are species (neutral or anionic) bonded to the central metal ion They may be attached to the metal center through a single atom (monodentate) or two or three atoms or higher (bidentate, tridentate, etc.) Such polydentate (bidentate or higher) ligands are called chelating ligands

Some very common & simple ligands

Some very common but chelating ligands

Self Study Isomerism in metal complexes Isomers are compounds with the same chemical formula but different structures ---Note that as they have different structures, they will have different physical and chemical properties.

Self Study

Variable oxidation state Why does Scandium only have one oxidation state? 1. Increase in the number Because if it had two - it would be Scandalous! of oxidation states from Sc to Mn. All are possible only in case of Mn. 2. Decrease in the number of oxidation states from Mn to Zn, due to the pairing of d-e ’ s after Mn 3. Stability of higher oxidn states decreases along Sc to Zn. Mn(VII) and Fe(VI) are powerful oxidizers. 4. Down the group, the stability of high oxidation states increases (easier availability of both d and s electrons for ionization).

18 electron rule (earlier EAN: Sidgwick) Stable metal complex at low oxidation state: metal electrons + lone pairs from ligands = 18 Ni(CO) 4 - 4s 2 3d 8 and 4 lone pairs = 18 Fe(CO) 5 - 4s 2 3d 6 and 5 lone pairs = 18 Cr(CO) 6 - 4s 2 3d 4 and 6 lone pairs = 18 18 electron rule: explained by MO theory . ( Filling of all the molecular bonding orbitals and none of the antibonding orbitals ). Drawback: not applicable for all complexes.

EAN: Uses and Limitations

Valence Bond Theory (VBT) G. N. Lewis (1902): atoms form covalent bonds by sharing electron pair. W. Heitler and F. London (1927): showed how the sharing of pairs of electrons holds a covalent molecule together. The Heitler-London model of covalent bonds was the basis of the VBT. L. Pauling : atomic orbitals are mixed to form hybrid orbitals , such as sp, sp 2 , sp 3 , dsp 2 , dsp 3 , and d 2 sp 3 orbitals.

VBT – Assumptions / Features (i) Ligands form covalent-coordinate bonds to the metal. (ii) Ligands must have lone pair of electrons. (iii) Available empty orbital of suitable energy for metal for bonding. (iv) Atomic (hybrid) orbitals are used for bonding (rather than molecular orbitals) Can explain : shape and stability of the metal complex. Can not explain : (i) Color (ii) Temperature dependence of magnetic properties

VBT – Examples Outer sphere complex Reactive/Labile High spin Paramagnetic Inner sphere complex Less reactive Low Spin Diamagnetic

Crystal Field Theory (CFT) CFT largely replaces VB Theory for interpreting the chemistry of coordination compounds. Ref: C. J. Ballhausen, J. Chem. Ed. 1979 56 194-197, 215-218, 357-361. J D Lee; pp. 204-222.

CFT: Assumptions  Pure electrostatic interaction  Ligands: point charges  Negative ligand: ion-ion interaction, neutral ligand: ion dipole interaction  The electrons on the metal are under repulsive from those on the ligands  The electrons on metal occupy those d-orbitals farthest away from the direction of approach of ligands

Symmetric field Degenerate 5 – d orbitals in an isolated gaseous metal atom. Spherically symmetric field of negative charges : All degenerate 5 – d orbitals are raised in energy due to repulsion from ligand.

d xy d xz d yz d z2 d x2-y2

d xy d xz d yz d z2 d x2-y2

d xy d xz d yz d z2 d x2-y2

d xy d xz d yz d z2 d x2-y2

d xy d xz d yz d z 2 d x 2 -y 2

Octahedral Field • Interaction of d-orbitals with six point charges at +x, -x, +y, -y, +z and -z axes are not same . • The orbitals lying along the axes (i.e. x 2 -y 2 , z 2 ) will be destabilized more than the orbitals lying in-between the axes (i.e. xy, xz, yz).

CFT- Octahedral Field

CFT-Octahedral Complexes For O h point group x 2 -y 2 , z 2 orbitals: e g xy, xz, yz orbitals: t 2g  Difference between t 2g and e g = Δ 0 or 10 Dq.  Conservation of barycenter from a spherical field to octahedral field indicates t 2g set must be stabilized as much as the e g set is destabilized .

Illustration of CFSE [Ti(H 2 O) 6 ] 3+ : 3d 1 complex e − in lowest energy t 2g orbitals. 1 e g 0 0 e g 1 : Purple color t 2g t 2g Maxima (UV-vis absorption spectrum): 20300 cm -1 So, Δ o = 243 kJ/mol . (1000 cm -1 = 11.96 kJ/mol or 2.86 kcal/mol or 0.124 eV) Typical Δ 0 values ~ energy of a chemical bond.

More than one d electrons • For d 1 -d 3 systems: t 2g 3 ( Hund's rule of maximum multiplicity) • For d 4 -d 7 systems: 4 e g 0 (low spin case or strong field situation) (i) t 2g 1 (high spin case or weak field situation) 3 e g (ii) t 2g Parameters (for HS and LS case): (i) CFSE (value of Δ 0 , 10Dq) (ii) Pairing energy of e - (repulsive)  CFSE > P.E.: LS complex  CFSE < P.E.: HS complex

(d 1 ) (d 2 ) (d 3 ) d 4 : Option I d 4 : Option II So which one? Decided by (i) Δ 0 High spin Low spin (ii)Pairing E. complexes complexes

Therefore, there are two important parameters to consider: • The Pairing energy (P) [is a repulsive energy] • e g to t 2g Splitting (referred to as Δ 0 , 10Dq) • For both the high spin (H.S.) and low spin (L.S.) situations, it is possible to compute the CFSE.

Δ o vs. P (pairing energy repulsive energy)

Δ o is dependent on L & M M 2+ < M 3+ < M 4+ 3d < 4d < 5d