Chemistry of Coordination Compounds (Chapter 24)

Chemistry of Coordination Compounds(Chapter 24)

The Structure of Complexes • We know Lewis acids are electron pair acceptors. • Coordination compounds: metal compounds formed by Lewis acid-base interactions. • Complexes: Have a metal ion (can be zero oxidation state) bonded to a number of ligands. Complex ions are charged. Example, [Ag(NH3)2]+. • Ligands are Lewis bases. • Square brackets enclose the metal ion and ligands. • Coordination Sphere: The area of space encompassing the ligands and metal ion.

The Structure of Complexes • 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+. • Charges, Coordination Numbers and Geometries • 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.

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). • Four coordinate complexes are either tetrahedral or square planar (commonly seen for d8 metal ions). • Six coordinate complexes are octahedral.

Chelates • Monodentate ligands bind through one donor atom only. • Therefore they occupy only one coordination site. • Polydentate ligands (or chelating agents) bind through more than one donor atom per ligand. • Example, ethylenediamine (en), H2NCH2CH2NH2. • The octahedral [Co(en)3]3+ is a typical en complex. • Chelate effect: More stable complexes are formed with chelating agents than the equivalent number of monodentate ligands.

Chelates

Chelates • [Ni(H2O)6]2+(aq) + 6NH3 [Ni(NH3)6]2+(aq) + 6H2O(l) • Kf = 4  108 • [Ni(H2O)6]2+(aq) + 3en [Ni(en)3]2+(aq) + 6H2O(l) • Kf = 2  1018 • Sequestering agents are chelating agents that are used to remove unwanted metal ions. • 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-).

Chelates

Chelates • EDTA occupies 6 coordination sites, for example [CoEDTA]- is an octahedral Co3+ complex. • Both N atoms (blue) and O atoms (red) coordinate to the metal. • EDTA is used in consumer products to complex the metal ions which catalyze decomposition reactions.

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+) and chlorophyll (Mg2+). • Myoglobin is protein containing a heme unit, which stores oxygen in cells.

Chelates Metals and Chelates in Living Systems

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 protein has a molecular weight of about 18,000 amu. • The Fe2+ ion in oxyhemoglobin or oxymyoglobin is octahedral. • Four N atoms from the porphyrin ring (red disk) are attached to the Fe2+ center.

Chelates • Metals and Chelates in Living Systems • The fifth coordination site is occupied by O2 (or H2O in deoxyhemoglobin or CO in carboxyhemoglobin). • The sixth coordination site is occupied by a base, which attaches the structure to the protein. • Photosynthesis is the conversion of CO2 and water to glucose and oxygen in plants in the presence of light. • One mole of sugar requires 48 moles of photons. • Chlorophyll absorbs red and blue light and is green in color. • The pigments that absorb this light are chlorophylls.

Chelates • Metals and Chelates in Living Systems • Chlorophyll a is the most abundant chlorophyll. • The other chlorophylls differ in the structure of the side chains. • 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).

Chelates • Metals and Chelates in Living Systems • The reaction • 6CO2 + 6H2O  C6H12O6 + 6O2 • is highly endothermic. • Plant photosynthesis sustains life on Earth.

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 tetraamminechlorocobalt(II). Note the tetra 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).

Nomenclature

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. • Oxidation state of the metal is given in Roman numerals in parenthesis at the end of the complex name.

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.

Isomerism

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.

Isomerism Structural Isomerism

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). • Stereoisomerism • Consider square planar [Pt(NH3)2Cl2]. • The two NH3 ligands can either be 90 apart or 180 apart. • Therefore, the spatial arrangement of the atoms is different.

Isomerism Stereoisomerism

Isomerism • Stereoisomerism • This is an example of geometrical isomerism. • 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. • 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.

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

Isomerism • Stereoisomerism • Enzymes are the most highly chiral substances known. • Most physical and chemical properties of enantiomers are identical. • Therefore, enantiomers are very difficult to separate. • Enzymes do a very good job of catalyzing the reaction of only one enantiomer. • Therefore, one enantiomer can produce a specific physiological effect whereas its mirror image produces a different effect.

Isomerism • Stereoisomerism • Enantiomers are capable of rotating the plane of polarized light. • Hence, they are called optical isomers. • When horizontally polarized light enters an optically active solution. • As the light emerges from the solution, the plane of polarity has changed. • The mirror image of an enantiomer will rotate the plane of polarized light by the same amount in the opposite direction.

Isomerism • Stereoisomerism • 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. • Chiral molecules are optically active because of their effect on light. • Racemic mixtures contain equal amounts of l- and d-isomers. They have no overall effect on the plane of polarized light.

Isomerism • Stereoisomerism • Pasteur was the first to separate racemic ammonium tartarate (NaNH4C4H9O6) by crystallizing the solution and physically picking out the “right-handed” crystals from the mixture using a microscope. • Optically pure tartarate can be used to separate a racemic mixture of [Co(en)3]Cl3: if d-tartarate is used, d-[Co(en)3]Cl3 precipitates leaving l-[Co(en)3]Cl3 in solution.

Color and Magnetism • Color • 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, d0 metal ions are usually colorless. Exceptions: MnO4- and CrO42-. • Colored compounds absorb visible light. • The color perceived is the sum of the light not absorbed by the complex.

Color and Magnetism Color

Color and Magnetism • Color • The amount of absorbed light versus wavelength is an absorption spectrum for a complex. • 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.

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.

Color and Magnetism • Magnetism • Many transition metal complexes are paramagnetic (i.e. they have unpaired electrons). • There are some interesting observations. Consider a d6 metal ion: • [Co(NH3)6]3+ has no unpaired electrons, but [CoF6]3- has four unpaired electrons per ion. • We need to develop a bonding theory to account for both color and magnetism in transition metal complexes.

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

Crystal-Field Theory • The complex metal ion has a lower energy than the separated metal and ligands. • However, there are some ligand-d-electron repulsions which occur since the metal has partially filled d-orbitals. • In an octahedral field, the degeneracy of the five d orbitals is lifted. • 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.

Crystal-Field Theory • We assume an octahedral array of negative charges placed around the metal ion (which is positive). • The dz2 and dx2-y2 orbitals lie on the same axes as negative charges. • Therefore, there is a large, unfavorable interaction between ligand (-) and these orbitals. • These orbitals form the degenerate high energy pair of energy levels. • The dxy, dyz, and dxz orbitals bisect the negative charges. • Therefore, there is a smaller repulsion between ligand and metal for these orbitals. • These orbitals form the degenerate low energy set of energy levels.

Crystal-Field Theory • The energy gap is the crystal field splitting energy . • Ti3+ is a d1 metal ion. • Therefore, the one electron is in a low energy orbital. • For Ti3+, the gap between energy levels,  is of the order of the wavelength of visible light. • 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. • Color of a complex depends on the magnitude of  which, in turn, depends on the metal and the types of ligands.

Chemistry of Coordination Compounds (Chapter 24)