Transition Metals and Complexes Transition Metals Complexes and Coordination Compounds Stereoisomerism of Complexes Polydentate Ligands and Chelate Complexes Constitutional Isomerism in Complexes Nomenclature of Complexes d Orbitals Bonding in Complexes

Transition metals Not as reactive as Group IA (1), IIA (2) metals and aluminum. d-block elements Except for palladium, all have either one or two s electrons in the outer shell. They only differ in the number of d electrons in n-1 energy level. Most have high melting and boiling points, high density, and are hard and strong. Multiple oxidation states are common.

Changes in properties down groups In general, as you move down a group: The outer electron configurations are the same. Reactivity decreases. Comparing the three series. The second and third series are usually more like each other than like the elements of the first transition series. Example, zirconium and hafnium always occur together in nature but titanium does not.

Changes in properties across periods Although the first period elements differ from the other two, properties vary in a similar manner as you move across a period. Lets look at several trends. Atomic radii Standard enthalpy of atomization Melting points Density Oxidation numbers.

Atomic radii Radius, pm Group number Radii go through a minimum just after the center of each series. Periods 5 and 6 are near identical. Radii go through a minimum just after the center of each series. Periods 5 and 6 are near identical.

Standard enthalpy of atomization  H  HThis is the energy required to convert an element to individual gaseous atoms. M (s) M (g) at 25 o C and one atmosphere. A maximum is observed near the middle of each period. The same trend is found for melting point, boiling point and density. oa

Standard enthalpy of atomization  H o a, kJ/mol Group number

Melting point, o C Group number

Density Density, g/cm 3 Group number

Transition metal trends Reason for trends. All show that attractions between atoms are strong. Strongest attractions are when the orbitals are half filled. Overlap of orbitals that contain one electron results in a covalent bond between atoms. Since metals at either end of a period have the fewest unpaired electrons, bonding is not as strong.

Oxidation states. The number of oxidation states also reaches a maximum near the center of each series. They can vary by only one unit whereas nonmetals typically vary by two. ns (n-1)dFor metals on the left side, both the ns and (n-1)d electrons can be involved in reactions. Because elements near the middle of each transition series have many oxidation states, much of the chemistry of these elements involves redox reactions.

Oxidation states Hg Hg Cd +2 Cd +2 Zn +2 Zn +2 Au Au Ag +1 Ag +1 Cu Cu Hf +4 Hf +4 Zr +4 Zr +4 Ti Ti Lu +3 Lu +3 Y +3 Y +3 Sc +3 Sc +3 Pt Pt Pd Pd Ni +2 Ni +2 Ir Ir Rh Rh Co Co Os Os Ru Ru Fe Fe Re Re Tc Tc Mn Mn W W Mo Mo Cr Cr Ta Ta Nb Nb V V

Magnetic properties Many of the transition metals and their compounds have magnetic properties. ferromagnetic.Iron and to a lesser extent, cobalt and nickel are ferromagnetic. –Can be permanently magnetized. paramagnetic.For most transition metals, at least one oxidation state is paramagnetic. –Attracted to a magnetic field.

Magnetic properties Paramagnetic species have unpaired electrons. The magnitude of the effect depends on the number of unpaired electrons.Example Paramagnetic Cr [Ar] Diamagnetic Pd[Kr] 3 d 4s

Formation of complexes and coordination compounds Formation of complexes is a characteristic property of transition metals. The most easily observed property of transition metal complexes is color. The color is dependent on the identity of the central atom, its oxidation state and the type of ligand. Coordination compounds. Those that include one or more complexes.

Complexes Alfred Werner studied the formation of coordination compounds of platinum. He found that one mole of platinum(IV) chloride would combine with 2, 3, 4, 5 or 6 moles of ammonia. PtCl 4. 2NH 3 PtCl 4. 3NH 3 PtCl 4. 4NH 3 PtCl 4. 5NH 3 PtCl 4. 6NH 3 This format was used to show the proportions of PtCl 4 and NH 3 that had combined.

Complexes Werner observed that the reactivities of the five coordination compounds differed. For example, addition of silver nitrate resulted in differing amounts of AgCl being formed. PtCl 4. 2NH 3 (aq) + excess Ag + no reaction PtCl 4. 3NH 3 (aq) + excess Ag + 1 AgCl PtCl 4. 6NH 3 (aq) + excess Ag + 4 AgCl This indicated that some of the chlorides must be bound to Pt.

Complexes Werner defined the coordination number as the number of atoms or groups that are firmly bound to the central atom. Empirical Number Number Number of Formulaof ions of Cl - nonionic Cl PtCl 4. 6NH PtCl 4. 5NH PtCl 4. 4NH PtCl 4. 3NH PtCl 4. 2NH

Complexes To indicate bound verses free ions, an alternate format for the formula was developed. [Pt(NH 3 ) 6 ]Cl 4 [PtCl(NH 3 ) 5 ]Cl 3 [PtCl 2 (NH 3 ) 4 ]Cl 2 [PtCl 3 (NH 3 ) 3 ]Cl [PtCl 4 (NH 3 ) 2 ] Coordinated atoms and groups are placed inside square braces. Free ions are placed on the outside. The symbol for the central atom is placed first. Ionic then neutral ligands are then listed in that order. Coordinated atoms and groups are placed inside square braces. Free ions are placed on the outside. The symbol for the central atom is placed first. Ionic then neutral ligands are then listed in that order.

Stereoisomerism in complexes Werner also evaluated the arrangement of coordinated groups around the central atom. He found that for one of the platinum complex, [PtCl 2 (NH 3 ) 2 ], two different isomers were observed. This could only be explained if the geometry was square planer.

Stereoisomerism in complexes Two geometric isomers were observed for [PtCl 2 (NH 3 ) 2 ]. cis-[PtCl 2 (NH 3 ) 2 ]trans-[PtCl 2 (NH 3 ) 2 ]

Stereoisomerism in complexes Some complexes exist in enantiomeric forms. To determine the number of forms that can exist, do the following. Draw or use a model kit to develop different geometric forms for your complex. Make a mirror image for each geometric isomer. If the model and its mirror image can be superimposed, they represent the same compound If they can’t be superimposed, they represent different compounds.

Example 1 The left and right forms can be superimposed. They represent the same compound. Original modelMirror image

Example Original modelMirror image The left and right forms can’t be superimposed. They represent the different compound.

Polydentate ligands and chelate complexes Ligands can be classified by dentate number - number of bonds/ligandMonodentate 1 bond/ligand - ammoniaBidentate 2 bonds/ligand - ethylene diamineMultidentate variable number based on need EDTA

Monodentate ligands Possess only one accessible donor group. H 2 O is a good example since all metal ions exist as aqua complexes in water. Although two e - pairs are available, only one is accessible. The other will always point the wrong way

Monodentate ligands Some aqua complexes Ag(H 2 O) 2 + Cu(H 2 O) 4 2+ Fe(H 2 O) 6 3+ The charge and coordination number are NOT related. Fe(H 2 O) 6 2+ also exists.

Monodentate ligands Common monodentate ligands anionic neutral X - OH - H 2 O SCN - RCOO - NH 3 CN - S 2- RNH 2

Bidentate ligands Form two bonds to central species. A good example is ethylene diamine. NH 2 CH 2 CH 2 NH 2 - (en) The amine groups are far enough apart to permit both to interact. Zn en Zn CN CN CN CN

Bidentate ligands Other common bidentate ligands. 8 - hydroxyquinoline O-O- N Zn O N Zn 2

Bidentate ligands Other common bidentate ligands. Dimethylglyoxime - dmg

Bidentate ligands Ni (dmg) 2

Bidentate ligands 1,10 phenanthroline

Bidentate ligands Iron(II) 1,10-phenanthroline complex

EDTA Ethylenediamine tetraacetic acid A commonly used ligand. Forms 1:1 complexes with most metals except the Group IA (1). Forms stable, water soluble complexes High formation constants. A primary standard material.

EDTA EDTA is typically used as the disodium salt to increase solubility. H 2 Y 2- H|H| H|H| |H|H |H|H

EDTA The molecule contains 6 donor groups. Regardless of the coordination number of the central species, the molecule will adapt to the number needed. Mg 2+ + H 2 Y 2- MgY H + Fe 3+ + H 2 Y 2- FeY - + 2H + H|H| H|H| |H|H |H|H

EDTA Fe 3+ - EDTA complex

EDTA Formation constants for some metal - EDTA complexes. IonlogKIonlogKIonlogK Fe Pb La Th Cd Mn Cr Zn Ca Bi Co Mg Cu Al Sr Ni Ce Ba

Constitutional isomerism Stereoisomers Differ only in the spatial arrangement of the ligands. Geometric and mirror-image isomers. Constitutional isomers LinkageLinkage - differ in how the ligands are attached to the central atom CompositionalCompositional - different quantities of solvent or ligand are present in the complex.

Linkage isomers Some ligands can attach to the central metal ion by either of two different atoms yellowred 2+ H 3 N Co H 3 N NH 3 3 H 3 N N OO 2+ H 3 N Co H 3 NNH 3 3 H 3 N O N O

Compositional isomers Different quantities of the solvent can be present in the complex. Example. Example. CrCl 3. (H 2 O) 6 Complex Color of aqueous solution [Cr(H 2 O) 6 ]Cl 3 violet [CrCl(H 2 O) 5 ]Cl 2 pale blue-green [CrCl 2 (H 2 O) 4 ]Cl dark green

Nomenclature of complexes In the name of a complex, the ligands are named first in alphabetical order. Prefixes are used to indicate the presence of multiple ligands of the same type The metal is then given followed by its oxidation state. -ateFor anionic complexes, the metal name is modified with an -ate ending. Latin names are used in some cases. There are no spaces between the various parts of the name.

Nomenclature of complexes LigandName F - Fluoro Cl - Chloro Br - Bromo I - Iodo OH - Hydroxo CN - Cyano H 2 OAqua NH 3 Ammine COCarbonyl LigandName F - Fluoro Cl - Chloro Br - Bromo I - Iodo OH - Hydroxo CN - Cyano H 2 OAqua NH 3 Ammine COCarbonyl MetalAnionic name CobaltCobaltate CopperCuprate GoldAurate IronFerrate LeadPlumbate MercuryMercurate NickelNickelate TinStannate ZincZincate MetalAnionic name CobaltCobaltate CopperCuprate GoldAurate IronFerrate LeadPlumbate MercuryMercurate NickelNickelate TinStannate ZincZincate

Nomenclature of complexes Examples [CoCl(NH 3 ) 5 ]Cl 2 pentaamminechlorocobalt(III) chloride K 2 [Cd(CN) 4 ] Potassium tetracyanocadmate(II) cis-diamminedichloroplatinum(II) H 3 N Pt Cl NH 3

d orbitals Bonding in transition-metal complexes involves the d orbitals. The are 5 d orbitals xy- lies in the xy plane. xz- lies in the xz plane. yz- lies in the yz plane. x 2 -y 2 - aligned with the x and y axes. z 2 - aligned with the z axis.

Representative d orbitals z2z2 xy

A complete set of d orbitals Red orbitals lie along the axes. Blue orbitals lie between axes. Note how spherical a complete set of d orbitals is.

Bonding in complexes Crystal field theory. Explains both the colors and magnetic properties of complexes. It was originally applied to metal ions in crystalline solids, which accounts for the name. According to the theory - Electrostatic attractions between the positive central metal ion and the ligands bonds them together, lowering the energy of the whole system.

Octahedral complexes d orbitals are degenerate for isolated gaseous metal ions - have the same energy. In complexes, this is not always the case. For octahedral, the ligands can be viewed as approaching the central species along the x-, y- and z- axes. The ligands will then interact with the d orbitals. Since the x 2 -y 2 and z 2 orbitals are closer to incoming ligands, they are affected more.

Octahedral complexes Ligands will approach along the x-, y- and z- axes. This affects the x 2 -y 2 and z 2 orbitals most.

Octahedral complexes The d orbitals are split into a high and low energy set of orbitals. The difference in energy is called the crystal field splitting and is represented by . Since this is for an octahedral complex, use  o Five degenerate d orbitals oo x 2 -y 2 z 2 xy xz yz

Octahedral complexes Crystal field splitting can be confirmed by looking at the paramagnetism of a complex. Electron pairing strong field or low spinIf the split is large, electrons will pair at the lower level before populating the upper level - strong field or low spin. weak field or high spinIf the split is small, they will fill all of the orbitals before attempting to pair up- weak field or high spin.

Octahedral complexes Example system with 4 d electrons. Strong field Low spin Weak field High spin

Tetrahedral complexes With this type of complex, the ligands approach closest to the d xy, d xz, and d zy orbitals. The splitting is exactly the opposite of that observed for octahedral complexes. Five degenerate d orbitals TT x 2 -y 2 z 2 xy xz yz

Square planar complexes These are a bit more complicated. One must imagine that the four ligands approach the central species along the x- and y- axes. The energy of the x 2 -y 2 orbital will be raised the most because it lies along the x and y axes. The xy orbital will also also be raised but not as much. The z 2 orbital can’t be predicted. Overall, we end up with 4 energy levels.

Energy levels octahedraltetrahedralsquare planar

So why are many complexes colored? For most transition metals, we have partially populated d orbitals. For a free atom, the energy required to move an electron from one electronic level to another is very large. It corresponds to the vacuum UV region. The energy of the crystal field split is much smaller and often corresponds to the energy of visible light.

UV/Vis absorption This ‘splitting’ of the d orbitals results in a d->d transition that is in the UV/Vis range. Chromium(III) examples Ligand max Cl H 2 O573 NH CN - 380

UV/Vis absorption Charge transfer complexes A complex where one species is an electron donor and the other is an electron acceptor. The resulting complex can be described as a resonance hybrid. These species tend to show very large absorbtivities (  max > 10,000) so many analytical methods are based on forming this type of complex.

UV/Vis absorption Charge transfer complexes N N 3+Fe 2+ N N 3 1,10-phenanthroline ferroin