prentice hall © 2003chapter 24 chapter 24 chemistry of coordination compounds chemistry the central...

Prentice Hall © 2003 Chapter 24

Chapter 24Chapter 24Chemistry of Coordination Chemistry of Coordination

CompoundsCompounds

CHEMISTRY The Central Science

9th Edition

David P. White


• We know Lewis acids are electron pair acceptors.• Coordination complexes: 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.

Metal ComplexesMetal Complexes


The Development of Coordination Chemistry: Werner’s Theory

• Werner discovered that CoCl3·nNH3 (n = 1 – 4) can exist as four different compounds with different numbers of “free” Cl- ions per formula unit.

• He deduced that the NH3 ligands were covalently bonded to the central Co3+ ion.

• Werner found that a total of six ligands were attached to the central Co.

• In the case of CoCl3·4NH3, there are two isomers for the Cl ligands attached to Co.



The Development of Coordination Chemistry: Werner’s Theory



The Metal-Ligand Bond• All ligands have lone pairs that are donated to the metal

ion.• The bond between metal and ligand is a 2-electron bond,

but both electrons come from the ligand and none from the metal.

• The metal-ligand bond alters the physical 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



The Metal-Ligand Bond

• 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.– Most common coordination numbers are 4 and 6.



Charges, Coordination Numbers, and Geometries

– 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.



• 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.

Ligands with More than Ligands with More than One Donor AtomOne Donor Atom


[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-).

• 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.



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.



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.



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.



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).



Metals and Chelates in Living Systems• The reaction

6CO2 + 6H2O C6H12O6 + 6O2

is highly endothermic. • Plant photosynthesis sustains life on Earth.



• 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 of Nomenclature of Coordination CompoundsCoordination Compounds


• 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.

Nomenclature of Nomenclature of Coordination CompoundsCoordination Compounds


• 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.

IsomerismIsomerism


• Stereoisomers have the same connectivity but different spatial arrangements of atoms.

IsomerismIsomerism


IsomerismIsomerism


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.

IsomerismIsomerism


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).

Structural Isomerism

• Consider square planar [Pt(NH3)2Cl2].

• The two NH3 ligands can either be 90 apart or 180 apart.

IsomerismIsomerism


Stereoisomers

IsomerismIsomerism


Structural Isomerism• The spatial arrangement of the atoms in the cis and trans

isomers is different.• 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.

IsomerismIsomerism



• For example, cis-[Co(NH3)4Cl2]+ is violet and trans-[Co(NH3)4Cl2]+ is green.

• The two isomers have different solubilities.• 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.)

IsomerismIsomerism


Structural Isomerism• 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).

IsomerismIsomerism



IsomerismIsomerism


Structural Isomerism• 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.

IsomerismIsomerism


Structural Isomerism• 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 in the opposite direction.

IsomerismIsomerism


Structural Isomerism• 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.

IsomerismIsomerism


Structural Isomerism• 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.

IsomerismIsomerism


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.

Color and MagnetismColor and Magnetism


Color• The color perceived is the sum of the light not absorbed

by the complex.• 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,



Color– the monochromatic light (i.e. a single wavelength) is passed

through the sample,

– the unabsorbed light is detected.



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.



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 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.

Crystal-Field TheoryCrystal-Field Theory


• The more directly the ligand attacks the metal orbital, the higher the energy of the d orbital.



• 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.



• 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 and metal for these orbitals.

– These orbitals form the degenerate low energy set of energy levels.

• 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.



• Spectrochemical series is a listing of ligands in order of increasing :

Cl- < F- < H2O < NH3 < en < NO2- (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 Cr3+ goes from being attached to a weak field ligand to a strong field ligand, increases.



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, Ti3+ is a d1 ion, V3+ is a d2 ion and Cr3+ is a d3 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 ();



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.



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.



Tetrahedral and Square-Planar Complexes

• Most d8 metal ions for square planar complexes.– the majority of complexes are low spin (i.e. diamagnetic).

– Examples: Pd2+, Pt2+, Ir+, and Au3+.


Tetrahedral and Square-Planar

Complexes• Most d8 metal ions for

square planar complexes.– the majority of

complexes are low spin (i.e. diamagnetic).

– Examples: Pd2+, Pt2+, Ir+, and Au3+.


End of Chapter 24End of Chapter 24Chemistry of Coordination Chemistry of Coordination

CompoundsCompounds

prentice hall © 2003chapter 24 chapter 24 chemistry of coordination compounds chemistry the central...

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