2 ComplexesA central metal atom bonded to a group of molecules or ions is a metal complex.If it’s charged, it’s a complex ion.Compounds containing complexes are coordination compounds.
3 ComplexesThe molecules or ions coordinating to the metal are the ligands.They are usually anions or polar molecules.The must have lone pairs to interact with metal
4 A chemical mystery: Same metal, same ligands, different number of ions when dissolved Many coordination compounds are brightly colored, but again, same metal, same ligands, different colors.
5 Werner’s TheoryCo(III) oxidation stateCoordination # is 6suggested in 1893 that metal ions have primary and secondary valences.Primary valence equal the metal’s oxidation numberSecondary valence is the number of atoms directly bonded to the metal (coordination number)
6 Werner’s TheoryThe central metal and the ligands directly bonded to it make up the coordination sphere of the complex.In CoCl3 ∙ 6 NH3, all six of the ligands are NH3 and the 3 chloride ions are outside the coordination sphere.
7 Werner’s TheoryIn CoCl3 ∙ 5 NH3 the five NH3 groups and one chlorine are bonded to the cobalt, and the other two chloride ions are outside the sphere.
8 Werner’s TheoryWerner proposed putting all molecules and ions within the sphere in brackets and those “free” anions (that dissociate from the complex ion when dissolved in water) outside the brackets.
9 Werner’s TheoryThis approach correctly predicts there would be two forms of CoCl3 ∙ 4 NH3.The formula would be written [Co(NH3)4Cl2]Cl.One of the two forms has the two chlorines next to each other.The other has the chlorines opposite each other.
10 What is Coordination?When an orbital from a ligand with lone pairs in it overlaps with an empty orbital from a metalSometimes called a coordinate covalent bondMLSo ligands must have lone pairs of electrons.
11 Metal-Ligand BondThis bond is formed between a Lewis acid and a Lewis base.The ligands (Lewis bases) have nonbonding electrons.The metal (Lewis acid) has empty orbitals.
12 Metal-Ligand BondThe metal’s coordination ligands and geometry can greatly alter its properties, such as color, or ease of oxidation.
13 Oxidation NumbersKnowing the charge on a complex ion and the charge on each ligand, one can determine the oxidation number for the metal.
14 Oxidation NumbersOr, knowing the oxidation number on the metal and the charges on the ligands, one can calculate the charge on the complex ion.Example: Cr(III)(H2O)4Cl2
15 Coordination NumberThe atom that supplies the lone pairs of electrons for the metal-ligand bond is the donor atom.The number of these atoms is the coordination number.
16 Coordination NumberSome metals, such as chromium(III) and cobalt(III), consistently have the same coordination number (6 in the case of these two metals).The most commonly encountered numbers are 4 and 6.
17 GeometriesThere are two common geometries for metals with a coordination number of four:TetrahedralSquare planarTetrahedralSquare planarWhy square planar? We’ll get to that
18 GeometriesBy far the most-encountered geometry, when the coordination number is six, is octahedral.
19 Polydentate Ligands Some ligands have two or more donor atoms. These are called polydentate ligands or chelating agents.In ethylenediamine, NH2CH2CH2NH2, represented here as en, each N is a donor atom.Therefore, en is bidentate.
20 Polydentate LigandsEthylenediaminetetraacetate, mercifully abbreviated EDTA, has six donor atoms.Wraps around the central atom like an octopus
21 Polydentate LigandsChelating agents generally form more stable complexes than do monodentate ligands.
22 Chelating Agents Bind to metal ions removing them from solution. 5--....-::....::::::......---Bind to metal ions removing them from solution.Phosphates are used to tie up Ca2+ and Mg2+ in hard water to prevent them from interfering with detergents.
23 Chelating AgentsPorphyrins are complexes containing a form of the porphine molecule shown at right.Important biomolecules like heme and chlorophyll are porphyrins.
24 Chelating AgentsPorphines (like chlorophyll a) are tetradentate ligands.
25 Nomenclature of Coordination Compounds The basic protocol in coordination nomenclature is to name the ligands attached to the metal as prefixes before the metal name.Some common ligands and their names are listed above.
26 Nomenclature of Coordination Compounds As always 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.
27 Nomenclature of Coordination Compounds 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.
28 Nomenclature of Coordination Compounds 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.
29 IsomersIsomers have the same molecular formula, but their atoms are arranged either in a different order (structural isomers) or spatial arrangement (stereoisomers).
30 Structural IsomersIf a ligand (like the NO2 group at the bottom of the complex) can bind to the metal with one or another atom as the donor atom, linkage isomers are formed.
31 Structural IsomersSome isomers differ in what ligands are bonded to the metal and what is outside the coordination sphere; these are coordination-sphere isomers.Three isomers of CrCl3(H2O)6 areThe violet [Cr(H2O)6]Cl3,The green [Cr(H2O)5Cl]Cl2 ∙ H2O, andThe (also) green [Cr(H2O)4Cl2]Cl ∙ 2 H2O.
32 Geometric isomersWith these geometric isomers, two chlorines and two NH3 groups are bonded to the platinum metal, but are clearly different.cis-Isomers have like groups on the same side.trans-Isomers have like groups on opposite sides.# of each atom the sameBonding the sameArrangement in space different
33 StereoisomersOther stereoisomers, called optical isomers or enantiomers, are mirror images of each other.Just as a right hand will not fit into a left glove, two enantiomers cannot be superimposed on each other.
34 EnantiomersA molecule or ion that exists as a pair of enantiomers is said to be chiral.
35 EnantiomersMost of the physical properties of chiral molecules are the same, boiling point, freezing point, density, etc.One exception is the interaction of a chiral molecule with plane-polarized light.
36 EnantiomersIf one enantiomer of a chiral compound is placed in a polarimeter and polarized light is shone through it, the plane of polarization of the light will rotate.If one enantiomer rotates the light 32° to the right, the other will rotate it 32° to the left.
39 Figure: 24-21Title:Optical activity.Caption:Effect of an optically active solution on the plane of polarization of plane-polarized light. The unpolarized light is passed through a polarizer. The resultant polarized light thereafter passes through a solution containing a dextrorotatory optical isomer. As a result, the plane of polarization of the light is rotated to the right relative to an observer looking toward the light source.
40 Explaining the properties of transition metal coordination complexes Magnetismcolor
41 Metal complexes and color The ligands of a metal complex effect its colorFigure: 24-22Title:Effect of ligands on color.Caption:An aqueous solution of CuSO4 is pale blue because of [Cu(H2O)4]2+ (left). When NH3(aq) is added (middle and right), the deep blue [Cu(NH3)4]2+ ion forms.Addition of NH3 ligand to Cu(H2O)4 changes its color
42 Why does anything have color? Light of different frequencies give different colorsWe learned that elements can emit light of different frequency or color.But these coordination complexes are not emitting lightThey absorb light.How does that give color?Figure: 24-23Title:Visible spectrum.Caption:The relationship between color and wavelength for visible light.
43 No light absorbed, all reflected get white color Light can bounce off an object or get absorbed by objectNo light absorbed, all reflected get white colorAll light absorbed, none reflected get Black colorWhat if only one color is absorbed?
44 Complimentary color wheel If one color absorbed, the color opposite is perceived.Absorb OrangeSee BlueAbsorb RedSee GreenFigure: 24-24Title:Artist's color wheel.Caption:Colors with approximate wavelength ranges are shown as wedges. The colors that are complementary to each other lie opposite each other.
45 Absorbs in green yellow. Looks purple. [Ti(H2O)6]3+Absorbs in green yellow.Looks purple.Figure: 24-26Title:The color of [Ti(H2O)6]3+.Caption:(a) A solution containing the [Ti(H2O)6]3+ ion. (b) The visible absorption spectrum of the [Ti(H2O)6]3+ ion.
46 A precise measurement of the absorption spectrum of Compounds is critical Figure: 24-25Title:Experimental determination of the absorption spectrum of a solution.Caption:The prism is rotated so that different wavelengths of light pass through the sample. The detector measures the amount of light reaching it, and this information can be displayed as the absorption at each wavelength. The absorbance is a measure of the light absorbed.
47 Metal complexes and color But why do different ligands on same metal giveDifferent colors?Why do different ligands change absorption?Figure: 24-22Title:Effect of ligands on color.Caption:An aqueous solution of CuSO4 is pale blue because of [Cu(H2O)4]2+ (left). When NH3(aq) is added (middle and right), the deep blue [Cu(NH3)4]2+ ion forms.Addition of NH3 ligand to Cu(H2O)4 changes its color
48 Assumption: interaction pure electrostatic. Model of ligand/metal bonding.Electron pair comes from ligandBond very polarized.Assumption: interaction pure electrostatic.Figure: 24-27Title:Metal-ligand bond formation.Caption:The ligand, which acts as a Lewis base, donates charge to the metal via a metal hybrid orbital. The bond that results is strongly polar, with some covalent character. It is often sufficient to assume that the metal-ligand interaction is entirely electrostatic in character, as is done in the crystal-field model.
49 Now, think of point charges being attracted to metal nucleus Positive charge. What about electrons in d orbitals?Ligand negative chargeIs repelled by d electrons, d orbital energy goes upFigure: 24-28Title:Energies of d orbitals in an octahedral crystal field.Caption:On the left the energies of the d orbitals of a free ion are shown. When negative charges are brought up to the ion, the average energy of the d orbitals increases (center). On the right the splitting of the d orbitals due to the octahedral field is shown. Because the repulsion felt by the dz2 and dx2–y2 orbitals is greater than that felt by the dxy, dxz, and dyz orbitals, the five d orbitals split into a lower-energy set of three (the t2 set) and a higher-energy set of two (the e set).
50 Ligands will interact with some d orbitals more than others Depends on relative orientation of orbital and ligandLigands point right at lobesFigure: 24-29abcTitle:The five d orbitals in an octahedral crystal field.Caption:(a) An octahedral array of negative charges approaching a metal ion. (b) The orientations of the d orbitals relative to the negative charges. Notice that the lobes of the dz2 and dx2–y2 orbitals (b and c) point toward the charges, whereas the lobes of the dxy, dyz, and dxz orbitals (d–f) point between the charges.
51 In these orbitals, the ligands are between the lobes Interact less stronglyFigure: 24-29defTitle:The five d orbitals in an octahedral crystal field.Caption:(a) An octahedral array of negative charges approaching a metal ion. (b) The orientations of the d orbitals relative to the negative charges. Notice that the lobes of the dz2 and dx2–y2 orbitals (b and c) point toward the charges, whereas the lobes of the dxy, dyz, and dxz orbitals (d–f) point between the charges.
53 Figure: 24-30Title:Electronic transition accompanying absorption of light.Caption:The 3d electron of [Ti(H2O)6]3+ is excited from the lower-energy d orbitals to the higher-energy ones when irradiated with light of 495-nm wavelength.= 495 nm
54 Different ligands interact more or less, change E spacing Of D orbitals.Figure: 24-31Title:Effect of ligand on crystal-field splitting.Caption:This series of octahedral chromium(III) complexes illustrates how , the energy gap between the t2 and e orbitals, increases as the field strength of the ligand increases.
55 Cl- < F- < H2O < NH3 < en < NO2- < CN- Spectrochemical series (strength of ligand interaction)Increasing Cl- < F- < H2O < NH3 < en < NO2- < CN-Increasing
56 Electron configurations of some octahedral complexes eleFigure: 24-32Title:Electron configurations in octahedral complexes.Caption:Orbital occupancies are shown for complexes with one, two, and three d electrons in an octahedral crystal field.
57 As Energy difference increases, electron configuration changes “Low spin”“High spin”Figure: 24-33Title:High-spin and low-spin complexes.Caption:Population of d orbitals in the high-spin [CoF6]3– ion (small )and low-spin [Co(CN)6]3– ion (large ). Both complexes contain cobalt(III), which has six 3d electrons, 3d6.Co(III) is d6
59 Tetrahedral Complexes In tetrahedral complexes, orbitals are inverted.Again because of orientation of orbitals and ligands is always small, always low spin (less ligands)Figure: 24-34Title:Tetrahedral crystal field.Caption:Energies of the d orbitals in a tetrahedral crystal field
60 Square planar complexes are different still Figure: 24-35Title:Energies of the d orbitals in a square-planar crystal field.Caption:The boxes on the left show the energies of the d orbitals in an octahedral crystal field. As the two negative charges along the z-axis are removed, the energies of the d orbitals change, as indicated by the lines between the boxes. When the charges are completely removed, the square-planar geometry results, giving the splitting of d-orbital energies represented by the boxes on the right.
62 Figure: UNEx10Title:Sample Exercise 24.10Caption:Orbital diagrams of tetrahedral and square planar–complexes.
63 Intense color can come from “charge transfer” Ligand electrons jump to metal orbitalsKMnO4KCrO4KClO4Figure: 24-36Title:Effect of charge-transfer transitions.Caption:The compounds KMnO4, K2CrO4, and KClO4 are shown from left to right. The KMnO4 and K2CrO4 are intensely colored because of ligand-to-metal charge transfer (LMCT) transitions in the MnO4– and CrO42– anions. There are no valence d orbitals on Cl, so the charge-transfer transition for ClO4– requires ultraviolet light and KClO4 is white.No d orbitals inCl, orbitals higherIn energy
64 Figure: 24-37Title:Ligand-to-metal charge transfer (LMCT) transition in MnO4–.Caption:As shown by the blue arrow, an electron is excited from a nonbonding pair on O into one of the empty d orbitals on Mn.
65 Exam 4, MO theory and coordination compounds Chapter 9, end and Chapter 24.MO theory: Rules:1. The number of MO’s equals the # of Atomic orbitals2. The overlap of two atomic orbitals gives two molecular orbitals, 1 bonding, one antibonding3. Atomic orbitals combine with other atomic orbitals of similar energy.4. Degree of overlap matters. More overlap means bonding orbital goes lower in E, antibonding orbital goes higher in E.5. Each MO gets two electrons6. Orbitals of the same energy get filled 1 electron at a time until they are filled.
66 Difference between pi and sigma orbitals End onSide to side.
67 A typical MO diagram, like the one below A typical MO diagram, like the one below. For 2p and 2s atomic orbital mixing.
70 Exam 4 Chapter 24. Concentrate on the homeworks and the quiz! Terms: Coordination sphereLigandCoordination compoundMetal complexComplex ionCoordinationCoordination numberSame ligands different properties?Figuring oxidation number on metal
71 Polydentate ligands (what are they)? Isomers.structural isomers (formula same, bonds differ)geometric isomers (formula AND bonds same, structure differs)Stereoisomers:Chirality, handedness,