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Transition Metales and Coordination Compound coordination bond [ Cu ionic bond ( NH3 )4 ]2+ Complex ion (inner sphere) SO42Counter ion (outer sphere) coordination compounds ★ coordination molecules have no outer sphere but inner sphere The molecules or ions that surround a metal ion in a complex are called ligands. Ligands may be anions or polar molecules. Some common ligands are water, ammonia, hydroxide, cyanide, and chloride. Each of these has at least one unshared pair of electrons that can be used to coordinate (form a coordinate covalent bond to the metal ion). Metal ions can act as Lewis acids, accepting electrons from ligands, which act as Lewis bases. The atom that is bound directly to the metal ion is called the donor atom. It's the atom that donates a lone pair of electrons to the metal ion. [ Cu(NH3)4 ] 2+ [ Fe(H2O) 6]3+ [ SiF6 ] 2Cu(Ⅱ) Fe(Ⅲ) NH3 H2O F O F N Si(Ⅳ) - : cations or atoms offering (central atom) empty orbitals : (Ligand) the molecules or ions that surround the metal ion in a complex : the atoms of ligand bound (donor atom) directly to the metal ★ Ligand is from the Latin word ligare, meaning “to bind” Coordination sphere – Metal and ligands bound to it Coordination number – number of donor atoms bonded to the central metal atom or ion in the complex • Most common = 4, 6 • Determined by ligands – Larger ligands and those that transfer substantial negative charge to metal favor lower coordination numbers Ligands – classified according to the number of donor atoms – Examples • monodentate = 1 • bidentate = 2 • tetradentate = 4 • hexadentate = 6 • polydentate = 2 or more donor atoms Coordination Compounds Common ligands: – water – ammonia H-O-H NH3 – chloride ion Cl - – cyanide ion C N - (monodentate ligand ): — the ligands possess a single donor atom and are able to occupy only one site in a coordination sphere • Monodentate – Examples: • H2O, CN-, NH3, NO2-, SCN-, OH-, X- (halides), CO, O2- – Example Complexes • [Co(NH3)6]3+ • [Fe(SCN)6]3- (bidentate ligand ): — the ligands possess two donor atoms – Examples • oxalate ion = C2O42• ethylenediamine (en) = NH2CH2CH2NH2 • ortho-phenanthroline (o-phen) – Example Complexes • [Co(en)3]3+ • [Cr(C2O4)3]3• [Fe(NH3)4(o-phen)]3+ H H H N - CH2- CH2 - N H Ethylenediamine (en) oxalate ion Orthophenanthroline carbonate ion bipyridine Structures of other bidentate ligands. The coordinating atoms are shown in blue. H H Ethylenediamine (en) H N - CH2- CH2 - N H (hexadentate ligand ): — the ligands possess six donor atoms ethylenediaminetetraacetate (EDTA) = (O2CCH2)2N(CH2)2N(CH2CO2)24– Example Complexes • [Fe(EDTA)]-1 • [Co(EDTA)]-1 Coordination Compounds Uses for EDTA – Used in food products to complex metals that promote decomposition reactions – Used to remove toxic heavy metals from the body – Used to complex metal ions that inhibit or promote polymerization reactions EDTA Porphine, an important chelating agent found in nature Metalloporphyrin N N 2+ NH NH N N Fe N N Coordination Compounds • Porphyrins are naturally occurring complexes derived from porphin. – Heme (Fe2+) – Chlorophyll (Mg2+) Hymoglobin, a protein that stores O2 in cells Coordination Environment of Fe2+ in Oxymyoglobin, and Oxyhemoglobin Ferrichrome (Involved in Fe transport in bacteria) Complex charge ☛ The charge on a central metal ion can be determined in much the same way that oxidation numbers are determined in molecular compounds. It is important to recognize common molecules, to remember that they are neutral, and to know the charges on both monoatomic and polyatomic ions. The sum of charges on the metal ion and the ligands must equal the net charge on the complex. Complex charge = sum of charges on the metal and the ligands [Fe(CN)6]3+3 6(-1) Neutral charge of coordination compound = sum of charges on metal, ligands, and counterbalancing ions neutral compound [Co(NH3)6]Cl2 +2 6(0) 2(-1) ☛ Coordination numbers of 2, 4, and 6 are common. The geometry of a complex is determined in part by its coordination number. (Size of the ligands can also impact the geometry.) Coordination Geometry Number 2 Linear 4 Tetrahedral(more common) or square planar 6 Octahedral Structures of geometries, (a) [Zn(NH3)4]2+ (b) [Pt(NH3)4]2+, tetrahedral square-planar These are the two common geometries for complexes in which the metal ion has a coordination number of 4. Attentions: In mono-core coordination compounds, each donor atom can only bind directly to one central atom to form coordination bond. e.g:H2O、 F Though some ligands have two or more donor atoms, they are still monodentate ligand because only one donor atom can form bond with central atom. e.g:SCN-、 NCS-、 NO2- 、ONO- 1. Simple complex — The one formed by central metal atom and monodentate ligand. e.g: [Cu(NH3)4]SO4、K[Ag(CN)2]、 K2[PtCl4]、etc. 2. (Chelate compound) — the ringed coordination compounds formed by metal atom and polydentate Polydentate ligands tend to coordinate more readily and form more stable complexes than monodentate ligands. This phenomenon is known as the Chelate effect . (chelating agents) — polydentate ligands are also known as chelating agents which are usually organic ligands possessing donor atoms like N、P、O、S. Cd2+ The [CoEDTA]– ion, showing how the Ethylene diamine tetraacetate ion is able to wrap around a metal ion, occupying six positions in the coordination sphere. Nomenclature of CC (IUPAC Rules) The cation is named before the anion When naming a complex: – Ligands are named first • alphabetical order – Metal atom/ion is named last • oxidation state given in Roman numerals follows in parentheses – Use no spaces in complex name [Co(NH3)5Cl]2+ Pentaamminechlorocobalt (III) Nomenclature of CC (IUPAC Rules) The names of anionic ligands end with the suffix -o – -ide suffix changed to -o – -ite suffix changed to -ito – -ate suffix changed to -ato Ligand Name bromide, Br- bromo chloride, Cl- chloro cyanide, CN- cyano hydroxide, OH- hydroxo oxide, O2- oxo fluoride, F- fluoro Nomenclature of CC (IUPAC Rules) Ligand Name carbonate, CO32- carbonato oxalate, C2O42- oxalato sulfate, SO42- sulfato thiocyanate, SCN- thiocyanato thiosulfate, S2O32- thiosulfato Sulfite, SO32- sulfito Nomenclature of CC (IUPAC Rules) Neutral ligands are referred to by the usual name for the molecule – Example • ethylenediamine – Exceptions • water, H2O = aqua • ammonia, NH3 = ammine • carbon monoxide, CO = carbonyl Nomenclature of CC (IUPAC Rules) Greek prefixes are used to indicate the number of each type of ligand when more than one is present in the complex – di-, 2; tri-, 3; tetra-, 4; penta-, 5; hexa-, 6 If the ligand name already contains a Greek prefix, use alternate prefixes: – bis-, 2; tris-, 3; tetrakis-,4; pentakis-, 5; hexakis-, 6 – The name of the ligand is placed in parentheses e.g: [Co(en)3]Cl3-tris(ethylenediamine)cobalt(III) chloride Nomenclature of CC (IUPAC Rules) If a complex is an anion, its name ends with the -ate – appended to name of the metal e.g: K4[Fe(CN)6] potassium hexacyanoferrate(II) [CoCl4]2- tetrachlorocobaltate(II) ion Nomenclature of CC (IUPAC Rules) Transition Metal Name if in Cationic Complex Name if in Anionic Complex Sc Scandium Scandate Ti titanium titanate V vanadium vanadate Cr chromium chromate Mn manganese manganate Fe iron ferrate Co cobalt cobaltate Ni nickel nickelate Cu Copper cuprate Zn Zinc zincate ★ Some names of coordination compounds [Co(NH3)4Cl2]Cl: tetraamminedichlorocobalt(Ⅲ) chloride [Pt(NH3)2Cl2]: diamminedichloroplatinum(Ⅱ) K[Ni(C2O4)2]: potassium bisoxalatonicketate(Ⅱ) oxalic acid [Cr (NH3)5H2O](NO3)3: pentaammineaquachromium (Ⅲ) nitrate Questions 1. Which name is correct for the coordination compound [Fe(en)2Cl2]Cl? A dichlorobisethylenediamineiron(III) chloride B diethylenediaminechlorineiron(I) chloride C iron(III) bisethylenediaminedichloro chloride D iron(III) trichloridebisethylenediamine Classes of isomers Dr.Monther F.Salem 42 ✽ Isomerism of CC isomer :Two or more compounds with the same elemental composition but different arrangements of atoms. — Major Types structural isomers stereoisomers Structural Isomers structural isomers : different bonds. — isomers that have 1. Coordination-sphere isomers – differ in a ligand bonded to the metal in the complex, as opposed to being outside the coordination-sphere • Example [Co(NH3)5Cl]Br vs. [Co(NH3)5Br]Cl [Co(NH3)5Cl]Br vs. [Co(NH3)5Br]Cl — Consider ionization in water [Co(NH3)5Cl]Br [Co(NH3)5Cl]+ + Br[Co(NH3)5Br]Cl [Co(NH3)5Br]+ + Cl- — Consider precipitation [Co(NH3)5Cl]Br(aq) + AgNO3(aq) [Co(NH3)5Cl]NO3(aq) + AgBr(s) [Co(NH3)5Br]Cl(aq) + AgNO3(aq) [Co(NH3)5Br]NO3(aq) + AgCl(aq) Linkage Isomers 2. Linkage isomers — differ in the atom of a ligand bonded to the metal in the complex Example: – [Co(NH3)5(ONO)]2+ vs. [Co(NH3)5(NO2)]2+ Example: – [Co(NH3)5(SCN)]2+ vs. [Co(NH3)5(NCS)] Co-SCN vs. Co-NCS Stereoisomers stereoisomers — isomers with the same bonds but different spatial arrangements of those bonds. 1. Geometric isomers -Differ in the spatial arrangements of the ligands -Have different chemical/physical properties different colors, melting points, polarities, solubilities, reactivities, etc. Geometric Isomers cis isomer, trans isomer, Pt(NH3)2Cl2 [PtCl2(NH3)2] (cis-platin), [PtCl2(NH3)2] (trans-platin), Geometric Isomers cis isomer trans isomer [Co(H2O)4Cl2]+ 2. Optical isomers – isomers that are nonsuperimposable mirror images • said to be “chiral” (handed) • referred to as enantiomers – A substance is “chiral” if it does not have a “plane of symmetry” Stereoisomers • Other 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. Enantiomers A molecule or ion that exists as a pair of enantiomers is said to be chiral. Example 1 mirror plane cis-[Co(en)2Cl2]+ Example 1 rotate mirror image 180° 180 ° Example 1 Nonsuperimposable, enantiomers cis-[Co(en)2Cl2]+ Example 2 mirror plane trans-[Co(en)2Cl2]+ Example 2 rotate mirror image 180° 180 ° trans-[Co(en)2Cl2]+ Example 2 Superimposable-not enantiomers trans-[Co(en)2Cl2]+ (a) Trans isomer of Co(en)2Cl2+ and its mirror image a are identical(superimposable) (b) cis isomer of Co(en)2Cl2+ No Optical activity Does have Optical activity Properties of Optical Isomers Enantiomers – possess many identical properties • solubility, melting point, boiling point, color, chemical reactivity (with nonchiral reagents) – different in: • interactions with plane polarized light • reactivity with “chiral” reagents Example Optical Isomers polarizing filter light source plane polarized light unpolarized light (random vibrations) (vibrates in one plane) Optical Isomers (+) dextrorotatory (right rotation ) (-) levorotatory (left rotation ) polarizing filter plane polarized light optically active sample in solution rotated polarized light polarizing filter plane polarized light optically active sample in solution Dextrorotatory (d) = right rotation Levorotatory (l) = left rotation Racemic mixture = equal amounts of two enantiomers; no net rotation rotated polarized light Geometry of NH3 NH3 NH3 2+ Zn NH3 NH3 2+ Geometry of [Zn(NH3)4] 2+ 2+ [Ar]18 Zn 30 4s 4p 3d10 sp3 [Zn(NH3)4] 2+ [Ar]18 3d10 Tetrahedron 4 个sp3 Crystal-Field Theory When ligands coordinate to a metal ion, they increase the energies of the metal‘s d orbitals by ligand d-orbital repulsion. Because of their spatial arrangement, the d orbitals do not all experience the same increase in energy. If we consider the formation of an octahedral complex in which the ligands approach the metal ion along the x, y, and z axes, we see that only orbitals that lie along the axes (the dz2 and dx2-y2) are approached directly. (红宝石) Crystal-Field Theory Crystal field theory describes bonding in transition metal complexes. The formation of a complex is a Lewis acid-base reaction. Both electrons in the bond come from the ligand and are donated into an empty, hybridized orbital on the metal. Charge is donated from the ligand to the metal. Assumption in crystal field theory: the interaction between ligand and metal is electrostatic. The more directly the ligand attacks the metal orbital, the higher the energy of the d orbital. Crystal Field Theory Octahedral Crystal Field - (-) Ligands attracted to (+) metal ion; provides stability - + d orbital e-’s repulsed by (–) ligands; increases d orbital potential energy - - ligands approach along x, y, z axes greater electrostatic repulsion = higher potential energy less electrostatic repulsion = lower potential energy Splitting of d orbital of metal atoms The complex metal ion has a lower energy than the separated metal and ligands. In an octahedral field, the five d orbitals do not have the same energy: three degenerate orbitals are higher energy than two degenerate orbitals. The energy gap between them is called , the crystal field splitting energy. Splitting of d-orbital Energies by a Tetrahedral Field and a Square Planar Field of Ligands Factors affecting the magnitude of ∆ 1. effect of ligands 2. Effect of metal atoms [Co(H2O)6]3+ △o=18600 cm-1 [Co(H2O)6]2+ △o= 9300 cm-1 [Co(NH3)6]3+ △o= 23000 cm-1 [Co(NH3)6]2+ △o= 10100 cm-1 [Co(NH3)6]3+ △o=23000 cm-1 [Rh(NH3)6]3+ △o= 33900 cm-1 [Ir(NH3)6]3+ △o= 40000 cm-1 Electronic Configurations of Transition Metal Complexes • Expected orbital filling tendencies for e-’s: – occupy a set of equal energy orbitals one at a time with spins parallel (Hund’s rule) • minimizes repulsions – occupy lowest energy vacant orbitals first • These are not always followed by transition metal complexes. d orbital occupancy depends on and pairing energy, P – e-’s assume the electron configuration with the lowest possible energy cost – If > P ( large; strong field ligand) • e-’s pair up in lower energy d subshell first – If < P ( small; weak field ligand) • e-’s spread out among all d orbitals before any pair up d-orbital energy level diagrams octahedral complex d1 d-orbital energy level diagrams octahedral complex d2 d-orbital energy level diagrams octahedral complex d3 d-orbital energy level diagrams octahedral complex d4 high spin <P low spin >P d-orbital energy level diagrams octahedral complex d5 high spin <P low spin >P d-orbital energy level diagrams octahedral complex d6 high spin <P low spin >P d-orbital energy level diagrams octahedral complex d7 high spin <P low spin >P d-orbital energy level diagrams octahedral complex d8 d-orbital energy level diagrams octahedral complex d9 d-orbital energy level diagrams octahedral complex d10 Electronic Configurations of Transition Metal Complexes Determining d-orbital energy level diagrams: – determine oxidation # of the metal – determine # of d e-’s – determine if ligand is weak field or strong field – draw energy level diagram High-spin and low-spin complex ions of Mn2+ Orbital Occupancy for High- and LowSpin Complexes of d4 Through d7 Metal Ions High spin weak-field ligand d4 d5 d6 d7 Low spin strong-field ligand Application of crystal field theory Crystal field theory can offer a good explanation of many properties of transition metal, such as stability、 spectrum、 and magnetism. Properties of Transition Metal Complexes Properties of transition metal complexes: – usually have color • dependent upon ligand(s) and metal ion – many are paramagnetic • due to unpaired d electrons • degree of paramagnetism dependent on ligand(s) – [Fe(CN)6]3- has 1 unpaired d electron – [FeF6]3- has 5 unpaired d electrons Crystal Field Theory – Can be used to account for • Colors of transition metal complexes – A complex must have partially filled d subshell on metal to exhibit color – A complex with 0 or 10 d e-s is colorless • Magnetic properties of transition metal complexes – Many are paramagnetic – # of unpaired electrons depends on the ligand 1. Account for the color of complexs Compounds/complexes that have color: – absorb specific wavelengths of visible light (400 –700 nm) • wavelengths not absorbed are transmitted • color observed = complementary color of color absorbed Visible Spectrum wavelength, nm (Each wavelength corresponds to a different color) 400 nm 700 nm higher energy lower energy White = all the colors (wavelengths) absorbed color observed color Relation Between Absorbed and Observed Colors Absorbed Color (nm) Violet Blue Blue-green Yellow-green Yellow Orange Red 400 450 490 570 580 600 650 Observed Color Green-yellow Yellow Red Violet Dark blue Blue Green (nm) 560 600 620 410 430 450 520 Color & Visible Spectra Color& Spectra 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 appears to be purple. Absorption of UV-visible radiation by atom, ion, or molecule: – Occurs only if radiation has the energy needed to raise an e- from its ground state to an excited state • i.e., from lower to higher energy orbital • light energy absorbed = energy difference between the ground state and excited state • “electron jumping” white light red light absorbed For transition metal complexes, corresponds to energies of visible light. green light observed Absorption raises an electron from the lower d subshell to the higher d subshell. Different complexes exhibit different colors because: – color of light absorbed depends on • larger = higher energy light (Shorter wavelengths) absorbed • smaller = lower energy light(Longer wavelengths) absorbed – magnitude of depends on: • ligand(s) • metal white light red light absorbed (lower energy light) [M(H2O)6]3+ green light observed Colors of Transition Metal Complexes white light blue light absorbed (higher energy light) [M(en)3]3+ orange light observed Colour of transition metal complexes Ruby Corundum Al2O3 with Cr3+ impurities Sapphire Corundum octahedral metal centre coordination number 6 Al2O3 with Fe2+ and Ti4+ impurities Emerald Beryl AlSiO3 containing Be with Cr3+ impurities The wavelength of light absorbed depends on the size of the energy gap, ∆ , between lower- and higher-energy d orbitals. The size of ∆ depends in part on the identity of the ligands coordinating to the metal ion. Figure below shows a series of four chromium(III) complexes, each with a different ∆ value. ∆1 3+ 24Cr ∆2 ∆3 ∆4 violet yellow yellow violet yellow violet 3d3 green red The energy gap between the three lower-energy d orbitals and the two higher-energy d orbitals is such that visible light causes this electron promotion. Thus, a particular wavelength of visible light is absorbed, giving a coordination complex its characteristic color. The gap between lower- and upper-level d orbitals in the [Ti(H2O)6]3+ ion is of the magnitude that light of 510 nm wavelength promotes an electron. Absorption of this wavelength of light (510 nm is a green wavelength) gives the ion its characteristic red color. 22 Ti 3+ 22Ti [Ar]3d24s2 [Ar]3d 4s 1 Ti: titanium 0 [Ti(H2O)6]3 • Color of a complex depends on: (i) the metal and (ii) its oxidation state. • 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. • So, d 0 metal ions are usually colorless. Exceptions: MnO4- and CrO42-. • Colored compounds absorb visible light. • The color our eye perceives is the sum of the light not absorbed by the complex.