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Chem 261: Inorganic Chemistry

Chem 261: Inorganic Chemistry. The elements in the periodic table are often divided into four categories: (1) main group elements, (2) transition metals, (3) lanthanides, and (4) actinides. How do we determine the electronic configuration of the central metal ion in any complex?.

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Chem 261: Inorganic Chemistry

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  1. Chem 261: Inorganic Chemistry

  2. The elements in the periodic table are often divided into four categories: (1) main group elements, (2) transition metals, (3) lanthanides, and (4) actinides.

  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 and knowing whether the ligands are neutral or anionic, so that you can determine the oxidation state of the metal ion. A simple procedure exists for the M(II) case. Cross off the first 2,

  4. Evaluating the oxidation state

  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. The most prevalent oxidation numbers are shown in green.

  6. An increase in the number of oxidation states from Sc to Mn. All seven oxidation states are exhibited by Mn. There is a decrease in the number 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. Mn(VII) and Fe(VI) are powerful oxidizing agents and the higher oxidation states of Co, Ni and Zn are unknown. The relative stability of +2 state with respect to higher oxidation states, particularly +3 state increases in moving from left to right. This is justifiable since it will be increasingly difficult to remove the third electron from the d-orbital.

  7. Summary of Physical Properties 1. have large charge/radius ratio; 2. are hard and have high densities; 3. have high melting and boiling points; 4. form compounds which are often paramagnetic; 5. show variable oxidation states; 6. form coloured ions and compounds; 7. form compounds with profound catalytic activity 8. form stable complexes.

  8. Coordination Chemistry A coordination compound, sometimes called a coordination 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.

  9. Ligands: [Zn(CN)4]2- CN- Zn2+ C 4 [PtCl6]2- Cl- Pt4+ Cl 6 [Ni(NH3)6]2+ NH3 Ni2+ N 6

  10. Mono-dentate Multidentate Ligands

  11. Chelating ligands bonded to metal – rings – chelate rings -any number of atoms in the ring. most common – five or six atoms, including metal. Coordination Numbers and Geometries

  12. Isomers Primarily in coordination numbers 4 and 6. Arrangement of ligands in space, but also the ligands themselves. Ionization isomers Isomers can produce different ions in solution e.g. [PtCl2(NH3)4]Br2 [PtBr2(NH3)4]Cl2 Polymerization isomers Same stoichiometry, different arrangement in space. Eight compounds with formula Co(NH3)3(NO2)3.

  13. Coordination isomers In compounds, both cation and anion are complex, the distribution of ligands can vary, giving rise to isomers. [Co(NH3)6]3+ [Cr(CN)6]-3 and [Cr(NH3)6]+3 [Co(CN)6]-3 Linkage isomers e.g. Nitro and nitito, N or O coordination possible

  14. 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.

  15. For hexacoordinate systems other species can also occur.

  16. For M(X)3(Y)3 systems there is facial and meridian

  17. Are “stereo” isomers also possible? An analogy to organic chirality. Molecules which can rotate light. Enantiomers (non-superimposable mirror images)

  18. Complex Stabilities In aqueous solution a comparison of metal complexes and their affinity for the H2O molecule as a competing ligand has been studied. Here are some general observations: • For a given metal and ligand, complexes where the metal oxidation state is +3 are more stable than +2. • Stabilities of complexes of the first row of transition metals vary in reverse of their cationic radii (in general) MnII < FeII < CoII < NiII > CuII > ZnII • Hard and soft Lewis acid-base theory • Chelate Effect - 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. 

  19. Mainly an entropy effect.

  20. 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 complexligand retained by end still attached so the nitrogen atom cannot move away – swings back and attach to metal again. Therefore the complex has a smaller probability of dissociating.

  21. The origin of colour - absorption

  22. The colour can change depending on a number of factors e.g. • Metal charge • Ligand

  23. Physical phenomenon

  24. 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

  25. Bonding in Transition Metal Complexes Covalent bonds by sharing pairs of electrons was first proposed by G. N. Lewis in 1902. It was not until 1927, however, that Walter Heitler and Fritz London 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 valence-bond theory. The last major step in the evolution of this theory was the suggestion by Linus Pauling that atomic orbitals mix to form hybrid orbitals, such as the sp, sp2, sp3, dsp3, and d2sp3 orbitals.

  26. How do we view this and do we really need to ?

  27. Valence-Bond Theory It is easy to apply the valence-bond theory to some coordination complexes, such as the Co3+ complexes below.

  28. d2sp3- inner sphere complex  low spin complex sp3d2- outer sphere complex  high spin complex Note: Such a situation will not arise for d1, d2 and d3 ion configuration.

  29. Deficiencies of VB approach to bonding • Assumes that all d orbitals in a complex are equal in energy. • The arbitrary use of 3d and 4d orbitals for bonding  energy differential ignored. • The theory is unable to adequately explain electronic and magnetic properties of complexes. • VBT is widely used in organic and main group element chemistry. • In TM metal chemistry VBT is superseded by the Crystal Field Theory (CFT).

  30. The Crystal-Field Theory Crystal Field Theory is based on the idea that a purely electrostatic interaction exists between the central metal ion and the ligands. Covalent bonding is ignored. Crystal field theory was developed by considering two compounds: manganese(II) oxide, MnO  octahedral geometry, copper(I) chloride, CuCl  tetrahedral geometry. We will start with octahedral and then expand to tetrahedral and square planar complexes

  31. The five d-orbitals in an octahedral field of ligands

  32. Tsuchida noticed a trend in while looking at a series of Cobalt(III) Complexes. With the general formula : [Co(NH3)5X] He arrived a series which illustrates the effect of ligands on Do (10Dq) He called it: The Spectrochemical Series Tsuchida, R. Bull. Chem. Soc. Jpn. 1938, 13, 388

  33. The magnitude of the splitting (ligand effect) Weak field Strong field The energy gap between t2gand eglevels is designated Do or 10Dq The spectrochemical series CO, CN- > phen > NO2- > en > NH3 > NCS- > H2O > F- > RCO2- > OH- > Cl- > Br- > I-

  34. The magnitude of the splitting (metal ion effect) Weak field Strong field • increases with increasing formal charge on the metal ion • increases on going down the periodic table

  35. Splitting of d orbitals in an octahedral field eg 3/5 Do Do 2/5 Do t2g E(t2g) = -0.4Do x 3 = -1.2Do E(eg) = +0.6Do x 2 = +1.2Do CFSE(Oh) = (–0.4x + 0.6y )Do + nP

  36. d5 d6 d7 1 u.e. 5 u.e. 0 u.e. 4 u.e. 1 u.e. 3 u.e. d8 d9 d10 1 u.e. 1 u.e. 2 u.e. 0 u.e. 0 u.e. 2 u.e. Placing electrons in d orbitals

  37. Crystal–Field Stabilisation Energy (CFSE) Q. Determine which of the following are more likely to be high spin complexes: [Fe(CN)6]3- [FeF6]3- [Co(H2O)6]+3 [Co(CN)6]-3 [Co(NH3)6]+3 [Co(en)3]+3 Solution: Compare the ligands on the spectrochemical series. Since we want a high spin complex, we want weak field ligands. The weaker field ligands in the above are H2O and F-, so complexes 2 and 3 are more likely to be high spin. (The cyanide complexes are least likely)

  38. When the 4th electron is assigned it will either go into the higher energy eg orbital at an energy cost of D0 or be paired at an energy cost of P, the pairing energy. d4 Strong field = Low spin (2 unpaired) Weak field = High spin (4 unpaired) P < Do P > Do Notes: the pairing energy, P, is made up of two parts. 1) Coulombic repulsion energy caused by having two electrons in same orbital

  39. Pairing Energy, P The pairing energy, P, is made up of two parts. • Coulombic repulsion energy caused by having two electrons in same orbital. Destabilizing energy contribution of Pc for each doubly occupied orbital. • Exchange stabilizing energy for each pair of electrons having the same spin and same energy. Stabilizing contribution of Pe for each pair having same spin and same energy P = sum of all Pc and Pe interactions

  40. Another way to view the energy in textbooks The energy increase of the eg orbitals and the energy decrease of the t2g orbitals must be balanced relative to the energy of the hypothetical spherical field (sometimes called the barycenter) The energy of the egset rises by +3/5Do = +6Dq while the energy of the t2g set falls by –2/5Do = –4Dq, resulting in no net energy change for the system. DE = E(eg)8 + E(t2g)9 = (2)(+3/5Do) + (3)(–2/5Do) = (2)(+6Dq) + (3)(–4Dq) = 0

  41. The magnitude of Do depends upon both the metal ion and the attaching ligands Magnitudes of Do are typically ~100 – 400 kJ/mol (~8,375 – 33,500 cm–1) 11 kJ/mol = 83.7 cm–1 Most aquo complexes are high spin, because H2O is a weak field ligand. Almost all Co3+ (d6) complexes are low spin, including [Co(H2O)6]3+, except [CoF6]3–, which is high spin. Second and third row transition metal ions tend to have low spin states - These ions tend to have larger Do values Larger 4d and 5d orbitals result in smaller P values, owing to lesser electronic repulsions 4d and 5d orbitals overlap with ligand orbitals, delocalizing electron density onto the ligands - Can we calculate or guestimate Do?

  42. In comparing groups of similar ligands - rationalize the order. These effects have been placed on a semi-quantitative basis by Jorgensen who assigned a factor g to a sampling of metal ions and a factor f: Δo ≈ g x f x 1000 cm-1

  43. High- spin d 4 Low- spin d 4 t2g3 eg1 t2g4eg0 x = 3 , y = 1 x = 4 , y = 0 E = (0.4x – 0.6y)Δo = 0.6 Δo E = (0.4x – 0.6y)Δo = 1.6 Δo + P

  44. Results and Observations 1. Doctahedral gets larger for increasing oxidation state 2. It increases down a group e.g. Co < Rh < Ir 3. With a given ligand and oxidation state Doctahedral varies irregularly across the first row transition metals The spectrochemical series The splitting of d orbitals in the crystal field model not only depends on the geometry of the complex, it also depends on the nature of the metal ion, the charge on this ion, and the ligands that surround the metal. When the geometry and the ligands are held constant, this splitting decreases in the following order.

  45. For metals the series is: Pt4+ > Ir3+ > Rh3+ > Co3+ > Cr3+ > Fe3+ > Fe2+ > Co2+ > Ni2+ > Mn2+ When the geometry and the metal are held constant, the splitting of the d-orbitals increases in the following order For ligands the series is: I - < Br - < [NCS] - < Cl - < F - < - OH < NH3 < en < CN - Weak field < Increasing DO < Strong field

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