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Chemistry of Coordination Compounds

Chemistry of Coordination Compounds. Chapter 24. General Remarks. A coordination complex or metal complex, consists of an atom or ion (usually metallic), and a surrounding array of bound molecules or anions, that are in turn known as ligands or complex agents .

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Chemistry of Coordination Compounds

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  1. Chemistry of Coordination Compounds Chapter 24 Chapter 24

  2. General Remarks • A coordination complex or metal complex, consists of an atom or ion (usually metallic), and a surrounding array of bound molecules or anions, that are in turn known as ligands or complex agents. • Many metal-containing compounds consist of coordination complexes. • Most of them are colorful, some are magnetic, most of them are reactive. Chapter 24

  3. The Coordination Compounds • Coordination compounds: metal compounds formed by Lewis acid-base interactions. • Metals are Lewis acid – accepting electron pairs • Ligands must be Lewis base – donating electron pairs • A general form is M (L)n, n=6, 4, 2 • We use square brackets enclose the metal ion and ligands. • Coordination Sphere: The area of space encompassing the ligands and metal ion. Chapter 24

  4. The Structure of Complexes • Charges, Coordination Numbers and Geometries • Complex ions can be charged. Example, [Ag(NH3)2]+. • Charge on complex ion = charge on metal + charges on ligands. • Donor atom: the atom bonded directly to the metal. • Coordination number: the number of ligands attached to the metal. Chapter 24

  5. The Structure of Complexes • The ligand donate unshared electron pairs to empty orbitals of metal to form the coordination bonds • Ligands can alter properties of the metal: Ag+(aq) + e- Ag(s), E = +0.799 V [Ag(CN)2]-(aq) + e- Ag(s) + 2CN-(aq), E = -0.031 V • Most metal ions in water exist as [M(H2O)6]n+. Chapter 24

  6. The Structure of Complexes • Charges, Coordination Numbers and Geometries • Most common coordination numbers are 4 and 6. • Some metal ions have constant coordination number (e.g. Cr3+ and Co3+ have coordination numbers of 6). • The size of the ligand affects the coordination number (e.g. [FeF6]3- forms but only [FeCl4]- is stable). • The amount of charge transferred from ligand to metal affects coordination number (e.g. [Ni(NH3)6]2+ is stable but only [Ni(CN)4]2- is stable). Chapter 24

  7. The Structure of Complexes • Charges, Coordination Numbers and Geometries • Four coordinate complexes are either tetrahedral or square planar (commonly seen for d8 metal ions). • Majority of six coordinate complexes are octahedral. Structures of (a) [Zn(NH3)4]2+ and (b) [Pt(NH3)4]2+, the tetrahedral and square-planar geometries, respectively. CN=6, octahedral coordination sphere Chapter 24

  8. Chelates • Monodentate (one-toothed) ligands bind through one donor atom only. • Therefore they occupy only one coordination site. • Polydentate (many toothed) ligands (or chelating agents) bind through more than one donor atom per ligand. Chapter 24

  9. Chelates For Example, ethylenediamine (en), H2NCH2CH2NH2 is a typical chelate ligand The octahedral [Co(en)3]3+ is a typical chelate complex. Chapter 24

  10. Common Chelate ligands Chapter 24

  11. Chelates Effect • [Ni(H2O)6]2+(aq) + 6NH3 [Ni(NH3)6]2+(aq) + 6H2O(l) Kf = 1.2  109 • [Ni(H2O)6]2+(aq) + 3en [Ni(en)3]2+(aq) + 6H2O(l) Kf = 6.8  1017 • Although in both cases the donor atoms are nitrogen, the free energy charges are very different. • The bind energies are approximately similar • The entropy changes are very different, make the big difference in free energy changes. Chapter 24

  12. Chelates Effect • In medicine sequestering agents are used to selectively remove toxic metal ions (e.g. Hg2+ and Pb2+) while leaving biologically important metals. • One very important chelating agent is ethylenediaminetetraacetate (EDTA4-). Chapter 24

  13. Chelates • Metals and Chelates in Living Systems • Many natural chelates are designed around the porphyrin molecule. • After the two H atoms bound to N are lost, porphyrin is a tetradentate ligand. • Porphyrins: Metal complexes derived from porphyrin. • Two important porphyrins are heme (Fe2+) (血红素) and chlorophyll (Mg2+) (叶绿素) • Myoglobin (肌红蛋白)is protein containing a heme unit, which stores oxygen in cells. Chapter 24

  14. Chelates Metals and Chelates in Living Systems The porphine molecules Myoglobin Chapter 24

  15. Chelates • Metals and Chelates in Living Systems • A five membered nitrogen containing ring binds the heme unit to the protein. • When oxygen is attached to the iron(II) in heme, oxymyoglobin is formed. • The Fe2+ ion in oxymyoglobin • is octahedral. • CO binds stronger than O2 Chapter 24

  16. Chelates Chlorophyll Photosynthesis is the conversion of CO2 and water to glucose and oxygen in plants in the presence of light. • Chlorophyll absorbs red and blue light and is green in color. • The pigments that absorb this light are chlorophylls. Chapter 24

  17. Chelates • Metals and Chelates in Living Systems • Mg2+ is in the center of the porphyrin-like ring. • The alternating double bonds give chlorophyll its green color (it absorbs red light). • Chlorophyll absorbs red light (655 nm) and blue light (430 nm). Adsorption spectrum of chlorophy11 and ground solar radiation. 655 nm - red, 439 nm - blue Chapter 24

  18. Chelates • Metals and Chelates in Living Systems • The reaction • 6CO2 + 6H2O  C6H12O6 + 6O2 • is highly endothermic, and ∆G=2897 kJ/mol • Plant photosynthesis sustains life on Earth. Chapter 24

  19. Nomenclature • Rules: • For salts, name the cation before the anion. Example in [Co(NH3)5Cl]Cl2 we name [Co(NH3)5Cl]2+ before Cl-. • Within a complex ion, the ligands are named (in alphabetical order) before the metal. Example [Co(NH3)5Cl]2+ is pentaamminechlorocobalt(II). Note the “penta” portion is an indication of the number of NH3 groups and is therefore not considered in the alphabetizing of the ligands. • Anionic ligands end in o and neutral ligands are simply the name of the molecule. Exceptions: H2O (aqua) and NH3 (ammine). For example: OH- hydroxo, C5H5N pyridine Chapter 24

  20. Nomenclature Chapter 24

  21. Nomenclature • Rules: • Greek prefixes are used to indicate number of ligands (di-, tri-, tetra-, penta-, and hexa-). • Exception: if the ligand name has a Greek prefix already. Then enclose the ligand name in parentheses and use bis-, tris-, tetrakis-, pentakis-, and hexakis. • Example [Co(en)3]Cl3 is tris(ethylenediamine)cobalt(III) chloride. • If the complex is an anion, the name ends in -ate. • [CoCl4]2- is called tetrachlorocobaltate(II) ion • Oxidation state of the metal is given in Roman numerals in parenthesis at the end of the complex name. • [Co(NH3)4Cl2]Cl tetraamminedichlorocobalt(III) chloride • [Pt(NH3)2Cl2] diamminedichloroplatinum(II) • [Cr(NH3)5H2O](NO3)3 pentaammineaquachromium(III) nitrate Chapter 24

  22. Isomerism • Isomers: two compounds with the same formulas but different arrangements of atoms. • Coordination-sphere isomers and linkage isomers: have different structures (i.e. different bonds). • Geometrical isomers and optical isomers are stereoisomers (i.e. have the same bonds, but different spatial arrangements of atoms). • Structural isomers have different connectivity of atoms. • Stereoisomers have the same connectivity but different spatial arrangements of atoms. Chapter 24

  23. Isomerism • Structural Isomerism • Some ligands can coordinate in different ways. That is, the ligand can link to the metal in different ways. • These ligands give rise to linkage isomerism. • Example: NO2- can coordinate through N or O (e.g. in two possible [Co(NH3)5(NO2)]2+ complexes). • When nitrate coordinates through N it is called nitro. • Pentaamminenitrocobalt(III) is yellow. • When ONO- coordinates through O it is called nitrito. • Pentaamminenitritocobalt(III) is red. Chapter 24

  24. Isomerism • Structural Isomerism • Example: NO2- can coordinate through N or O (e.g. in two possible [Co(NH3)5(NO2)]2+ complexes). • When nitrate coordinates through N it is called nitro. • Pentaamminenitrocobalt(III) is yellow. • When ONO- coordinates through O it is called nitrito. • Pentaamminenitritocobalt(III) is red. Chapter 24

  25. Isomerism • Structural Isomerism • Similarly, SCN- can coordinate through S or N. • Coordination sphere isomerism occurs when ligands from outside the coordination sphere move inside. • Example: CrCl3(H2O)6 has three coordination sphere isomers: [Cr(H2O)6]Cl3 (violet), [Cr(H2O)5Cl]Cl2.H2O (green), and [Cr(H2O)4Cl2]Cl.2H2O (green). Chapter 24

  26. Isomerism • Stereoisomerism • Consider square planar [Pt(NH3)2Cl2]. • The two NH3 ligands can either be 90 apart or 180 apart. • In the cis isomer, the two NH3 groups are adjacent. The cis isomer (cisplatin) is used in chemotherapy. • The trans isomer has the two NH3 groups across from each other. Chapter 24

  27. Isomerism • Stereoisomerism • It is possible to find cis and trans isomers in octahedral complexes. • For example, cis-[Co(NH3)4Cl2]+ is violet and trans-[Co(NH3)4Cl2]+ is green. • The two isomers have different solubilities. Chapter 24

  28. Isomerism • Stereoisomerism • In general, geometrical isomers have different physical and chemical properties. • It is not possible to form geometrical isomers with tetrahedral complexes. (All corners of a tetrahedron are identical.) Chapter 24

  29. Isomerism • Stereoisomerism – Optical isomers • Optical isomers are mirror images which cannot be superimposed on each other. • Optical isomers are called enantiomers. • Complexes which can form enantiomers are chiral. • Most of the human body is chiral (the hands, for example). Chapter 24

  30. Isomerism • Stereoisomerism • Most physical and chemical properties of enantiomers are identical. • Enantiomers are different only in chiral environment. Very important in pharmaceuticals • Enantiomers are capable of rotating the plane of polarized light, called optical isomers. Chapter 24

  31. Isomerism • Stereoisomerism • Chiral molecules are optically active because of their effect on light. • Dextrorotatory solutions rotate the plane of polarized light to the right. This isomer is called the d-isomer. • Levorotatory solutions rotate the plane of polarized light to the left. This isomer is called the l-isomer. • Racemic mixtures contain equal amounts of l- and d-isomers. They have no overall effect on the plane of polarized light. Chapter 24

  32. Isomerism • Separation of optical isomers • Enzymes are the most highly chiral substances known. One enantiomer can produce a specific physiological effect whereas its mirror image produces a different effect. • Chemical method - reach a racemic mixture with a chiral acid/base to form a salt. Since the salt exhibits different chemical and physical properties, it is possible to separate them. • Chiral Chromatography: Pass racemic mixture through a column containing a chiral solid matrix (one isomer of a chiral solid) and the different enantiomers will have different elution times Chapter 24

  33. Color and Magnetism • Color • Color of a transition metal complex depends on: (i) the metal, (ii) its oxidation state, and (iii) the ligands • Pale blue [Cu(H2O)6]2+ can be converted into dark blue [Cu(NH3)6]2+ by adding NH3(aq). • A partially filled d orbital is usually required for a complex to be colored. • Colored compounds absorb visible light. The color perceived is the sum of the light not absorbed by the complex. Chapter 24

  34. Color and Magnetism • Color • To determine the absorption spectrum of a complex: • a narrow beam of light is passed through a prism (which separates the light into different wavelengths), • the prism is rotated so that different wavelengths of light are produced as a function of time, • the monochromatic light (i.e. a single wavelength) is passed through the sample, • the unabsorbed light is detected. Chapter 24

  35. Color and Magnetism • Color • The plot of absorbance versus wavelength is the absorption spectrum. • For example, the absorption spectrum for [Ti(H2O)6]3+ has a maximum absorption occurs at 510 nm (green and yellow). • So, the complex transmits all light except green and yellow. • Therefore, the complex is purple. • From the energy adsorbed • is known. Chapter 24

  36. Color and Magnetism • Magnetism • Many transition metal complexes are paramagnetic (i.e. they have unpaired electrons). • Magnetism provides important bonding information. • There are three types of magnetic behavior • Diamagnetic (no response to external magnetic field); • Paramagnetic (stabilized in magnetic field); • Ferromagnetic (permanent magnetic). Chapter 23

  37. Transition Metals Magnetism • Electron spin generates a magnetic field with a magnetic moment. Single electron behaves like a “tiny magnet”. • Paired electrons do not have the net magnet. –Diamagnetic. • The tiny magnet interact differently with each other depending on the electron configuration • If the interactions are not strong enough, the spins are randomly oriented, unless external field is applied. - paramagnetic • If the interactions are strong enough, the spins of all atoms are aligned - ferromagnetic Chapter 23

  38. Bonding Theories Chapter 24

  39. Valence Bond Theory Co(NH3)63+ ion has octahedral structure. The electron configuration of the transition-metal ion Co3+: [Ar] 3d6 4s0 4p0 It may adopt hybrid orbitals (takes energy): Then form 6 coordination bonds using d2sp3 (release energy) "inner-shell" complexes, which use 3d, 4s and 4p orbitals to form a set of d2sp3hybrids,

  40. Valence Bond Theory Ni(CO)4has tetrahedral structure Consider electron configuration of Ni Ni: [Ar] 3d8 4s2 4p0 Rearranged, takes energy Ni: [Ar] 3d10 4s0 4p0 Then the SP3 hydride orbitals accept electrons from ligands, release energy Chapter 24

  41. Valence Bond Theory Ni(NH3)62+ has octahedral structure The center atom Ni2+: [Ar] 3d8, even rearrange electrons to 4 3d orbitals, Cannot form d2sp3 But, use 4d, we may use 1 4s, 3 4pand 2 4dorbitals to form sp3d2 hybrid orbitals (does not cost much energy) "outer-shell" complexes, which use 4s, 4p and 4d orbitals to form sp3d2hybrid orbitals. Chapter 24

  42. More Theories • Valence Theory can be used to understand structures, but not spectral and magnetic properties. • Crystal field theory (CFT) describes the energy splitting of d orbitals in transition metal complexes. CFT accounts for some magnetic properties, colors of transition metal complexes, but it does not attempt to describe the chemical bonding. • The Ligand Field Theory (a branch of the molecular orbital theory) treats bonding mechanisms. Chapter 24

  43. Crystal-Field Theory • CFT supposes that each ligand can be represented by a point negative charge. • The interaction between ligand and metal is due to electrostatic interactions between negative charge and electrons. • The energy splitting of d orbital is due to the electron repulsion between the d orbital and the ligand • The more directly the ligand attacks the metal orbital, the higher the energy of the d orbital. Chapter 24

  44. Crystal-Field Theory • The complex has a lower energy than the separated metal and ligands due to coordination bonding • However, the ligand-d-electron repulsions raise the energy since the metal has partially filled d-orbitals. • The differences in the repulsion cause the energy splitting Chapter 24

  45. Crystal-Field Theory • In an octahedral field, • the five d orbitals do not have the same energy • The dz2 and dx2-y2 lie on the same axes as negative charges - stronger • The dxy, dyz, and dxz orbitals bisect the negative charges - weaker Chapter 24

  46. Crystal-Field Theory • Due to interactions, the energies of 5 d orbitals split: • The energies of dz2 and dx2-y2orbitals are higher • The energies of dxy, dyz, and dxz orbitals are lower • The energy gap  is called crystal field splitting energy 0.6∆ -0.4∆ Chapter 24

  47. Crystal-Field Theory • The splitting is affected by the following factors: • the nature of the metal ion. • the metal's oxidation state. A higher oxidation state leads to a larger splitting. • the arrangement of the ligands around the metal ion. • the nature of the ligands surrounding the metal ion. The stronger the effect of the ligands then the greater the difference between the high and low energy d groups. Chapter 24

  48. Crystal-Field Theory • Ti3+ is a d1 metal ion. The one electron is in a low energy orbital. • As the [Ti(H2O)6]3+ complex absorbs visible light, the electron is promoted to a higher energy level, . • Since there is only one d electron, there is only one possible absorption line for this molecule. • For Ti3+, the gap between energy levels,  is of the order of the wavelength of visible light. Chapter 24

  49. Crystal-Field Theory • Splitting energy depends on the strength of ligands • Spectrochemical series is a listing of ligands in order of increasing  • The higher slitting energy, the higher strength, and higher adsorption frequency Chapter 24

  50. Crystal-Field Theory • As Cr3+ goes from being attached to a weak field ligand to a strong field ligand,  increases and the color of the complex changes from green to yellow. Chapter 24

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