Chemistry of Coordination Compounds David P White University

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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(NH 3)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. Copyright 1999, PRENTICE HALL Chapter 24 2

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) + 2 CN-(aq), E = -0. 031 V • Most metal ions in water exist as [M(H 2 O)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. Copyright 1999, PRENTICE HALL Chapter 24 3

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. Cr 3+ and Co 3+ have coordination numbers of 6). – The size of the ligand affects the coordination number (e. g. [Fe. F 6]3 - forms but only [Fe. Cl 4]- is stable). – The amount of charge transferred from ligand to metal affects coordination number (e. g. [Ni(NH 3)6]2+ is stable but only [Ni(CN)4]2 - is stable). • Four coordinate complexes are either tetrahedral or square planar (commonly seen for d 8 metal ions). • Six coordinate complexes are octahedral. Copyright 1999, PRENTICE HALL Chapter 24 4

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), H 2 NCH 2 NH 2. • 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. Copyright 1999, PRENTICE HALL Chapter 24 5

Chelates [Ni(H 2 O)6]2+(aq) + 6 NH 3 [Ni(NH 3)6]2+(aq) + 6 H 2 O(l) Kf = 4 108 [Ni(H 2 O)6]2+(aq) + 3 en [Ni(en)3]2+(aq) + 6 H 2 O(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. Hg 2+ and Pb 2+) while leaving biologically important metals. • One very important chelating agent is ethylenediaminetetraacetate (EDTA 4 -). Copyright 1999, PRENTICE HALL Chapter 24 7

Chelates • EDTA occupies 6 coordination sites, for example [Co. EDTA]- is an octahedral Co 3+ 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. Copyright 1999, PRENTICE HALL Chapter 24 9

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 (Fe 2+) and chlorophyll (Mg 2+). • Myoglobin is protein containing a heme unit, which stores oxygen in cells. Copyright 1999, PRENTICE HALL Chapter 24 10

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 Fe 2+ ion in oxyhemoglobin or oxymyoglobin is octahedral. • Four N atoms from the porphyrin ring (red disk) are attached to the Fe 2+ center. Copyright 1999, PRENTICE HALL Chapter 24 12

Chelates Metals and Chelates in Living Systems • The fifth coordination site is occupied by O 2 (or H 2 O 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 CO 2 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. Copyright 1999, PRENTICE HALL Chapter 24 14

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. • Mg 2+ 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). Copyright 1999, PRENTICE HALL Chapter 24 16

Chelates Metals and Chelates in Living Systems • The reaction 6 CO 2 + 6 H 2 O C 6 H 12 O 6 + 6 O 2 is highly endothermic. • Plant photosynthesis sustains life on Earth. Copyright 1999, PRENTICE HALL Chapter 24 18

Nomenclature • Rules: – For salts, name the cation before the anion. Example in [Co(NH 3)5 Cl]Cl 2 we name [Co(NH 3)5 Cl]2+ before Cl-. – Within a complex ion, the ligands are named (in alphabetical order) before the metal. Example [Co(NH 3)5 Cl]2+ is tetraamminechlorocobalt(II). Note the tetra portion is an indication of the number of NH 3 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: H 2 O (aqua) and NH 3 (ammine). Copyright 1999, PRENTICE HALL Chapter 24 19

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]Cl 3 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. Copyright 1999, PRENTICE HALL Chapter 24 21

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. Copyright 1999, PRENTICE HALL Chapter 24 22

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: NO 2 - can coordinate through N or O (e. g. in two possible [Co(NH 3)5(NO 2)]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. Copyright 1999, PRENTICE HALL Chapter 24 24

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: Cr. Cl 3(H 2 O)6 has three coordination sphere isomers: [Cr(H 2 O)6]Cl 3 (violet), [Cr(H 2 O)5 Cl]Cl 2. H 2 O (green), and [Cr(H 2 O)4 Cl 2]Cl. 2 H 2 O (green). Stereoisomerism • Consider square planar [Pt(NH 3)2 Cl 2]. • The two NH 3 ligands can either be 90 apart or 180 apart. • Therefore, the spatial arrangement of the atoms is different. Copyright 1999, PRENTICE HALL Chapter 24 26

Isomerism Stereoisomerism • This is an example of geometrical isomerism. • In the cis isomer, the two NH 3 groups are adjacent. The cis isomer (cisplatin) is used in chemotherapy. • The trans isomer has the two NH 3 groups across from each other. • It is possible to find cis and trans isomers in octahedral complexes. • For example, cis-[Co(NH 3)4 Cl 2]+ is violet and trans[Co(NH 3)4 Cl 2]+ is green. • The two isomers have different solubilities. Copyright 1999, PRENTICE HALL Chapter 24 28

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). Copyright 1999, PRENTICE HALL Chapter 24 30

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. Copyright 1999, PRENTICE HALL Chapter 24 32

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. Copyright 1999, PRENTICE HALL Chapter 24 34

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 disomers. They have no overall effect on the plane of polarized light. Copyright 1999, PRENTICE HALL Chapter 24 35

Isomerism Stereoisomerism • Pasteur was the first to separate racemic ammonium tartarate (Na. NH 4 C 4 H 9 O 6) 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]Cl 3: if d-tartarate is used, d-[Co(en)3]Cl 3 precipitates leaving l-[Co(en)3]Cl 3 in solution. Copyright 1999, PRENTICE HALL Chapter 24 36

Color and Magnetism Color • Color of a complex depends on: (i) the metal and (ii) its oxidation state. • Pale blue [Cu(H 2 O)6]2+ can be converted into dark blue [Cu(NH 3)6]2+ by adding NH 3(aq). • A partially filled d orbital is usually required for a complex to be colored. • So, d 0 metal ions are usually colorless. Exceptions: Mn. O 4 - and Cr. O 42 -. • Colored compounds absorb visible light. • The color perceived is the sum of the light not absorbed by the complex. Copyright 1999, PRENTICE HALL Chapter 24 37

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. Copyright 1999, PRENTICE HALL Chapter 24 39

Color and Magnetism Color • The plot of absorbance versus wavelength is the absorption spectrum. • For example, the absorption spectrum for [Ti(H 2 O)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. Copyright 1999, PRENTICE HALL Chapter 24 41

Color and Magnetism • Many transition metal complexes are paramagnetic (i. e. they have unpaired electrons). • There are some interesting observations. Consider a d 6 metal ion: – [Co(NH 3)6]3+ has no unpaired electrons, but [Co. F 6]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. Copyright 1999, PRENTICE HALL Chapter 24 43

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 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 metal is electrostatic. • The more directly the ligand attacks the metal orbital, the higher the energy of the d orbital. Copyright 1999, PRENTICE HALL Chapter 24 44

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 dorbitals. • 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. Copyright 1999, PRENTICE HALL Chapter 24 46

Crystal-Field Theory • We assume an octahedral array of negative charges placed around the metal ion (which is positive). • The and 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 metal for these orbitals. – These orbitals form the degenerate low energy set of energy levels. Copyright 1999, PRENTICE HALL Chapter 24 48

Crystal-Field Theory The energy gap is the crystal field splitting energy . Ti 3+ is a d 1 metal ion. Therefore, the one electron is in a low energy orbital. For Ti 3+, the gap between energy levels, is of the order of the wavelength of visible light. • As the [Ti(H 2 O)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. • • Copyright 1999, PRENTICE HALL Chapter 24 50

Crystal-Field Theory • Spectrochemical series is a listing of ligands in order of increasing : Cl- < F- < H 2 O < NH 3 < en < NO 2 - (N-bonded) < CN • Weak field ligands lie on the low end of the spectrochemical series. • Strong field ligands lie on the high end of the spectrochemical series. • As Cr 3+ 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. Copyright 1999, PRENTICE HALL Chapter 24 54

Crystal-Field Theory Electron Configurations in Octahedral Complexes • We still apply Hund’s rule to the d-orbitals. • The first three electrons go into different d orbitals with their spins parallel. • Recall: the s electrons are lost first. • So, Ti 3+ is a d 1 ion, V 3+ is a d 2 ion and Cr 3+ is a d 3 ion. • We have a choice for the placement of the fourth electron: – if it goes into a higher energy orbital, then there is an energy cost ( ); Copyright 1999, PRENTICE HALL Chapter 24 56

Crystal-Field Theory Electron Configurations in Octahedral Complexes – if it goes into a lower energy orbital, there is a different energy cost (called the spin-pairing energy due to pairing an electron). • Weak field ligands tend to favor adding electrons to the higher energy orbitals (high spin complexes) because < pairing energy. • Strong field ligands tend to favor adding electrons to lower energy orbitals (low spin complexes) because > pairing energy. Copyright 1999, PRENTICE HALL Chapter 24 59

Crystal-Field Theory Tetrahedral and Square-Planar Complexes • By using the same arguments as for the octahedral case, we can derive the relative orbital energies for d orbitals in a tetrahedral field. • The exact opposite of an octahedral field orbital arrangement is found (i. e. the dxy, dyz, and dxz orbitals are of lower energy than the and orbitals). • Because there are only 4 ligands, for a tetrahedral field is smaller than for an octahedral field (four ninths). • This causes all tetrahedral complexes to be high spin. Copyright 1999, PRENTICE HALL Chapter 24 60

Crystal-Field Theory Tetrahedral and Square-Planar Complexes • Square planar complexes can be thought of as follows: start with an octahedral complex and remove two ligands along the z-axis. • As a consequence the four planar ligands are drawn in towards the metal. • Relative to the octahedral field, the orbital is greatly lowered in energy, the dyz, and dxz orbitals lowered in energy, the dxy, and orbitals are raised in energy. • Most d 8 metal ions for square planar complexes. – the majority of complexes are low spin (i. e. diamagnetic). – Examples: Pd 2+, Pt 2+, Ir+, and Au 3+. Copyright 1999, PRENTICE HALL Chapter 24 62