COORDINATION CHEMISTRY STRUCTURES AND ISOMERS

COORDINATION CHEMISTRY STRUCTURES AND ISOMERS

ELECTRONIC CONFIGURATION
Ground State: Progressive filling of the 3d, 4d, and 5d orbitals Exceptions: ns1 (n-1)d5 rather than ns2 (n-1)d4 ns1 (n-1)d10 rather than ns2 (n-1)d9 Transition metal ions: First in first out

TRENDS - IONIC Radii

COORDINATION COMPOUNDS
Coordination compounds – compounds composed of a metal atom or ion and one or more ligands. [Co(Co(NH3)4(OH2)3]Br6 Ligands usually donate electrons to the metal Includes organometallic compounds Werner’s totally inorganic optically active compound.

WERNER’S COORDINATION CHEMISTRY
Performed systematic studies to understand bonding in coordination compounds. Organic bonding theory and simple ideas of ionic charges were not sufficient. Two types of bonding Primary – positive charge of the metal ion is balanced by negative ions in the compound. Secondary – molecules or ion (ligands) are attached directly to the metal ion. Coordination sphere or complex ion. Look at complex on previous slide (primary and secondary)

WERNER’S COORDINATION CHEMISTRY
He largely studied compounds with four or six ligands. Octahedral and square-planar complexes. It was illustrated that a theory needed to account for bonds between ligands and the metal. The number of bonds was commonly more than accepted at that time. 18-electron rule. New theories arose to describe bonding. Valence bond, crystal field, and ligand field.

COORDINATION COMPOUNDS

LIGANDS

CHELATING LIGANDS Chelating ligands (chelates) – ligands that have two or more points of attachment to the metal atom or ion. Bidentate, tridentate, tetra.., penta…, hexa… (EDTA). trisoxalatochromate(III) ion or just [Cr(ox)3]3-

A HEXADENTATE LIGAND, EDTA
There are six points of attachment to the calcium metal. Octahedral-type geometry ethylene diamine tetraacetic acid (EDTA) ethylenediaminetetraacetatocalcium ion or just [Ca(EDTA)]2-

LIGANDS

NOMENCLATURE Coordination compounds that are ionic, the cation is named first and separated by a space from the anion, as is the case for all ionic compounds. The names of neutral coordination complexes are written without spaces. Na[PtCl3(NH3)] Sodium amminetrichloroplatinate(II) K2[CuBr4] Potassium tetrabromocuprate(II)

NOMENCLATURE trans-[Co(en)2I(H2O)](NO3)2 trans-aquabis(ethylenediamine)iodocobalt(III) nitrate mer-[Ru(PPh3)3Cl3] mer-trichlorotris(triphenylphosphine)ruthenium(III)

NOMENCLATURE The name of the coordination compound (neutral, cationic or anionic) begins with the names of the ligands. The metal is listed next, following in parentheses by the oxidation state of the metal.

NOMENCLATURE When more than one of a given ligand is bound to the same metal atom or ion, the number of such ligands is designated by the following prefixes: 2 di 6 hexa 10 deca 3 tri 7 hepta 11 undeca 4 tetra 8 octa 12 dodeca 5 penta 9 nona

NOMENCLATURE However, when the name of the ligand in question already contains one of these prefixes or ligands with complicated names (generally ligand names that are three syllables or longer), then a prefix from the following list is used instead: 2 bis 6 hexakis 3 tris 7 heptakis 4 tetrakis 8 octakis 5 pentakis 9 ennea

NOMENCLATURE Neutral ligands are given the same name as the uncoordinated molecule, but with spaces omitted. Some examples are: (CH3)3SO dimethylsulfoxide (DMSO) (NH2)2CO urea C5H5N pyridine terpy terpyridine bpy 2,2’-bipyridine en ethylenediamine PCl3 trichlorophosphine PPh3 triphenylphopshine

NOMENCLATURE EXCEPTIONS: Some neutral molecules, when serving as ligands are given special names. These are: NH3 ammine H2O aqua NO nitrosyl CO carbonyl CS thiocarbonyl

NOMENCLATURE Anionic ligands are given names that end in the letter “o”. When the name of the free, uncoordinated anion ends in “ate”, the ligand name is changed to end in “ato”. Some examples are : CH3CO2- (acetate) acetato SO42- (sulfate) sulfato CO32- (carbonate) carbonato acac acetylacetonato

NOMENCLATURE When the name of the free, uncoordinated anion ends in “ide”, the ligand name is changed to end in “ido”. Some examples are: N3- (nitride) nitrido N3- (azide) azido NH2- (amide) amido

NOMENCLATURE When the name of the free, uncoordinated anion ends in “ite”, the ligand name is changed to end in “ito”. Some examples are: NO2- (nitrite) nitrito SO32- (sulfite) sulfito ClO3- (chlorite) chlorito

NOMENCLATURE Certain anionic ligands are given special names, all ending in “o”: CN- cyano F- fluoro Cl- chloro Br- bromo I- iodo O2- oxo O2- superoxo OH- hydroxo H- hydrido CH3O- methoxo

NOMENCLATURE The ligands are named alphabetically, ignoring the prefixes bis, tris, etc…

NOMENCLATURE When the coordination entity is either neutral or cationic, the usual name of the metal is used, followed in parentheses by the oxidation state of the metal. However, when the coordination entity is an anion, the name of the metal is altered to end in “ate”. This is done for some metals by simply changing the ending “ium” to “ate”: Scandium scandate Titanium titanate Chromium chromate

NOMENCLATURE Geometrical isomers are designated by cis- or trans- and mer- or fac-, the latter two standing for meridional or facial, respectively.

NOMENCLATURE Bridging ligands are designated with the prefix -. When there are two bridging ligands of the same kind, the prefix di-- is used. Bridging ligands are listed in order with other ligands, and set off between hypens. An important exception arises when the molecule is symmetrical, and a more compact name can be given by listing the bridging ligand first.

NOMENCLATURE Example:

NOMENCLATURE

NOMENCLATURE Ligands that are capable of linkage isomerism are given specific names for each mode of attachment. -SCN- thiocyanato (S-thiocyanato) -NCS- isothiocyanto (N-thiocyanto) -NCSe- isoselenocyanato (N-selenocyanato) -NO2- nitro -ONO- nitrito

EXAMPLES [Co(NH3)5CO3]Cl Potassium pentachloronitridoosmate(VI)
[Cr(H2O)4Cl2]Cl Potassium pentacyanonitrosylferrate(II) 5. K4[Mn(CN)6] 6. [Ni(bipy)3(NO3)2] 7. [Co(N3)(NH3)5]SO4

NOMENCLATURE Bridging ligands between two metal ions have the prefix ‘ ’. -amido--hydroxobis(tetraamminecobalt)(IV)

ISOMERISM

ISOMERISM Four-coordinate complexes
Square-planar complexes may have cis and trans isomers. No chiral isomers (enantiomers) are possible when the molecule has a mirror plane. cis- and trans-diamminedichloroplatinum(II) How about tetrahedral complexes? Chelate rings commonly impose a ‘cis’ structure. Why

ISOMERISM

CHIRALITY Mirror images are nonsuperimposable.
A molecule can be chiral if it has no rotation-reflection axes (Sn) Chiral molecules have no symmetry elements or only have an axes of proper rotation (Cn). CBrClFI, Tetrahedral molecule (different ligands) Octahedral molecules with bidentate or higher chelating ligands Octahedral species with [Ma2b2c2], [Mabc2d2], [Mabcd3], [Mabcde2], or [Mabcdef]

CHIRALITY

SIX-COORDINATE OCTAHEDRAL COMPLEXES
ML3L3’ Fac isomers have three identical ligands on the same face. Mer isomers have three identical ligands in a plane bisecting the molecule.

ISOMERISM

SIX-COORDINATE OCTAHEDRAL COMPLEXES
The maximum number of isomers can be difficult to calculate (repeats). Placing a pair of ligands in the notation <ab> indicates that a and b are trans to each other. [M<ab><cd><ef>], [Pt<pyNH3><NO2Cl><BrI>] How many diastereoisomers in the above platinum compound (not mirror images)? Identify all isomers belonging to Ma3bcd.

COMBINATIONS OF CHELATE RINGS
Propellers and helices Left- and right-handed propellers Examine the movement of a propeller required to move it in a certain direction. For a left-handed propeller, rotating it ccw would cause it to move away (). For a right-handed propeller, rotating it cw would cause it to move away (). This is called ‘handedness’. Many molecules possess it.

Tris(ethylenediamine)cobalt(III)
this molecule can be treated like a three-bladed propeller. look down a three fold axis to determine the ‘handedness’ of this complex ion. the direction of rotation required to pull the molecule away from you determines the handedness ( or ). do this with you molecule set and rubber bands.

DETERMINING HANDEDNESS FOR CHIRAL MOLECULES
Complexes with two or more nonadjacent chelate rings may have chiral character. Any two noncoplanar and nonadjacent chelate rings can be used. Look at Figure 9-14 (Miessler and Tarr). Molecules with more than one pair of rings may require more than one label. Ca(EDTA)2+ Three labels would be required. Remember that the chelate rings must be noncoplanar, nonadjacent, and not connected at the same atom.

LINKAGE (AMBIDENTATE) ISOMERISM
A few ligands may bond to the metal through different atoms. SCN- and NO2- How would you expect hard acids to bond to the thiocyanate ligand? Solvents can also influence bonding. High and low dielectric constants. Steric effects of linkage isomerism Intramolecular conversion between linkages. [Co(NH3)5NO2]+2, Figure 9-19.

COORDINATION NUMBERS AND STRUCTURES
Factors considered when determining structures. The number of bonds. Bond formation is exothermic; the more the better. VSEPR arguments Occupancy of d orbitals. Steric interference by large ligands. Crystal packing effect. It may be difficult to predict shapes.

LOW COORDINATION NUMBERS (C.N.)
CN 1 is rare except in ion pairs in the gas phase. CN 2 is also rare. [Ag(NH3)2]+, Ag is d10 (how?) VSEPR predicts a linear structure. Large ligands help force a linear or near-linear arrangment. [Mn(N[SiMePh2]2)2] in Figure 9-22. C.N. 3 is more likely with d10 ions. Trigonal-planar structure is the most common. [Cu(SPPh3)3]+, adopts a low C.N. due to ligand crowding.

COORDINATION NUMBER 4 Tetrahedral and square planar complexes are the most common. Small ions and/or large ligands prevent high coordination numbers (Mn(VII) or Cr(VI)). Many d0 or d10 complexes have tetrahedral structures (only consider bonds). MnO4- and [Ni(CO)4] Jahn-Teller distortion (Chapter 10)

COORDINATION NUMBER 4 Square-planar geometry
d8 ions (Ni(II), Pd(II), and Pt(III)) [Pt(NH3)2Cl2] The energy difference between square-planar and tetrahedral structures can be quite small. Can depend on both the ligand and counterion. More in chapter 10.

COORDINATION NUMBER 5 Common structures are trigonal bipyramid and square pyramid. The energy difference between the two is small. In many measurements, the five ligands appear identical due to fluxional behavior. How would you modify the experiment to differentiate between the two structures? Five-coordinate compounds are known for the full range of transition metals. Figure 9-27.

COORDINATION NUMBER 6 This is the most common C.N. with the most common structure being octahedral. If the d electrons are ignored, this is the predicted shape. [Co(en)3]3+ This C.N. exists for all transition metals (d0 to d10).

DISTORTIONS OF COMPLEXES CONTAINING C.N. 6
Elongation and compression (Fig. 9-29). Produces a trigonal antiprism structure when the angle between the top and bottom triangular faces is 60. Trigonal prism structures are produced when the faces are eclipsed. Most trigonal prismatic complexes have three bidentate ligands (Figure 9-30).  interactions may stabilize some of these structures. The Jahn-Teller effect is useful in predicting observed distortions.

HIGHER COORDINATION NUMBERS
C.N. 7 is not common C.N. 8 There are many 8-coordinate complexes for large transition elements. Square antiprism and dodecahedron C.N.’s up to 16 have been observed.

MAGNETIC SUSCEPTIBILITY
Diamagnetic versus paramagnetic complexes. Commonly provides mass susceptibility per gram. Magnetic moment 

CONTRIBUTIONS TO THE MAGNETIC MOMENT
Spin magnetic moment S = maximum total spin in the complex O atom Orbital angular momentum Characterized by the quantum number L which is equal the maximum possible sum of ml values.

CONTRIBUTIONS TO THE MAGNETIC MOMENT
Usually, the spin-only moment is sufficient to calculate the magnetic moment. Especially for the first transition series where g (gyromagnetic ratio) is approximated to be 2 and n is the number of unpaired electrons. Determine the spin-only and complete magnetic moment for Fe.

Calculate the spin-only magnetic moment For the following atoms/ions: Fe+2 (observed: 5.1), Fe, Cr, Cr+3 (observed = 3.8)

ELECTRONIC SPECTRA Orbital energy levels can be obtained directly from electron spectra (will be covered later). This chapter illustrates simple energy level diagrams that are commonly more complex. Based upon subtle differences in electronic spectra, the structure may be predicted with some success.

THEORIES OF ELECTRONIC STRUCTURE
Valence Bond Theory – Not commonly used, but the hybrid notation is still common. Crystal Field Theory – An electrostatic approach used to describe the splitting in metal d-orbital energies. Does not describe bonding. Ligand Field Theory – A more complete description of bonding in terms of the electronic energy levels of the frontier orbitals. Commonly does not include energy of the bonding orbitals. Angular Overlap Method – Used to estimate the relative magnitude of the orbital energies in a MO calculation.

VALENCE BOND THEORY (HYBRIDIZATION)
A set of hybrid orbitals is produced to explain the bonding. Octahedral – d2sp3 (6 hybrid orbitals of equal energy) Tetrahedral - ?? Uses ‘inner’ and ‘outer’ orbitals to explain the experimentally determined unpaired electrons. The magnetic behavior determines which d orbitals (e.g. 3d or 4d) are used for bonding (Figure 10-2).

VALENCE BOND DESCRIPTION
Two configurations are possible for d4-d7 ions. Fe(III) has 5 electrons in the d-orbitals. One unpaired electron, the ligands are ‘strong’ and force the metal d electrons to pair up. Strong-field (bind strongly)  low spin complex The hybridization orginates from the 3d inner orbitals (d2sp3).

VALENCE BOND DESCRIPTION
Five unpaired electrons, the ligands are ‘weak’ and cannot force the metal d electrons to pair up. Weak-field (bind weakly) high spin The hybridization originates from the 4d outer orbitals (sp3d2).

VALENCE BOND THEORY Structure, hybridization, and magnetism [Co(NH3)6]3+, diamagnetic, octahedral [CoF6]3-, paramagnetic, octahedral [PtCl4]2-, diamagnetic, sq. planar [NiCl4]2-, pamagnetic, tetrahedral

SAMPLE PROBLEM: The complexes [Mn(H2O)6]2+, [Fe(H2O)6]3+, [MnCl4]2-, and [FeCl4]- have all magnetic moments. What does this tell about the geometric and electronic structures of these complexes?

CRYSTAL FIELD THEORY Focus: energies of the d orbitals Assumptions
1. Ligands: negative point charges 2. Metal-ligand bonding: entirely ionic strong-field (low-spin): large splitting of d orbitals weak-field (high-spin): small splitting of d orbitals

D = crystal field splitting

High spin Low spin

CRYSTAL FIELD THEORY The average energy of the d-orbitals in the presence of the octahedral field is greater than than of the free ion. Energy difference between the two sets is equal to O. The t2g set is lowered by 0.4 O and the eg set is raised by 0.6 O. Crystal field stabilization energy (CFSE) – The energy difference between the actual distribution of electrons and that for all electrons in the uniform field. Equal to LFSE (later) Drawbacks

LIGAND FIELD THEORY – OCTAHEDRAL COMPLEXES
Consider -type bonding between the ligands and the metal atom/ion. Construct LGOs (performed previously). What is the reducible representation? Construct the LGOs (pictures). Construct the molecular orbitals with the metal orbitals. Same symmetry types. A group of metal orbitals do not have the appropriate symmetry? Which orbitals are these? Symmetry type? Bonding? Look at Figure 10-5.

SF6 = A1g + T1u + Eg

The six bonding orbitals are largely filled by the electrons from the ligands. The higher MOs (e.g. t2g and eg) are largely filled by the electrons on the metal atom/ion. The ligand field treatment largely focuses on the t2g and higher orbitals. The split between the two sets of orbitals, t2g and eg, is called O.

Ligands whose orbitals interact strongly with the metal orbitals are called strong-field ligands. Strong-field  large O  low spin (why?) Ligands with small interactions are called weak-field ligands. Weak-field  small O  high spin (why?) For d0 – d3 and d8-d10 only one electron configuration is possible (no difference in net spin). For d4 – d7 there is a difference between strong- and weak-field cases.

LOW SPIN VERSUS HIGH SPIN
Energy of pairing electrons c is the Coulombic energy of repulsion (always positive when pairing) and e is the quantum mechanical exchange energy (always negative). e relates to the number of exchangeable pairs in a particular electron configuration. This term is negative and depends on the number of possible states. Determine c and e for a d5 metal complex (low and high spin).

The relationship between O, c, and e determines the orbital configuration.  is largely independent on the ligands while O is strongly dependent. Look at Table 10-6 which gives these parameters for aqueous (aqua) ions. O for 3+ ions is larger than O for 2+ ions. O values for d5 are smaller than d4 and d6.

If O>, there is a lower energy upon pairing in the lower levels (low spin). If O<, there is a lower energy with unpaired electrons in the lower levels (high spin). In Table 10-6, [Co(H2O)6]3+ is probably the only complex that could be low spin.

Ligand Field Stabilization Energies (LFSE)
The difference (1) the total energy of a coordination complex with the electron configuration resulting from ligand field splitting of the orbitals and (2) the total energy for the same complex with all the orbitals equally populated is the LFSE. -2/5O + 3/5O (d4 to d7 complexes) Table 10-7

 BONDING IN OCTAHEDRAL COMPLEXES
The x and z axes must be taken as a single set producing a combined LGO set. Why? Be able to derive the reducible representation.  = T1g + T2g + T1u + T2u How will the LGOs combine with orbitals from the metal atom/ion? Discuss the overlap between the -bonding LGOs and the p-orbitals of T1u symmetry.

PI BONDING IN OCTAHEDRAL COMPLEXES
The main addition to the interaction diagram is between the t2g orbitals of the metal and LGOs. These were nonbonding when only considering -type bonding (look at Figure 10-5). Pi bonding may occur when the ligands have available p or * molecular orbitals.

LIGANDS WITH EMPTY * ORBITALS
Examine the example for the CN- ligand in the book (Figure 10-9). The HOMO forms the LGOs from -type bonding (already discussed previously). The LUMO, 1*, also forms a reducible set of LGOs (T1g + T2g + T1u + T2u). Examine Figure to illustrate effectiveness of overlap.

LIGANDS WITH EMPTY * ORBITALS
The resulting t2g LGOs are generally higher in energy than the initial t2g orbitals on he metal. Bonding/antibonding t2g orbitals will result. What will this do to O and the bond strength? Figure This is termed as metal-to-ligand  bonding or  back-bonding. Some of the electron density in the d orbitals on the metal is donated back to the ligands. The ligands are termed as -acceptor ligands.

LIGANDS WITH FILLED -TYPE ORBITALS
Ligands such as F- or Cl- will possess molecular  orbitals that possess electrons. This set of ‘t2g’ orbitals are generally lower in energy than the t2g orbitals on the metal. What are the consequences? Examine Figure Ligand-to-metal  bonding (-donor ligands). This bonding is generally less favorable.

SQUARE-PLANAR COMPLEXES
The y-axis is pointed toward the center atom. LGOs for sigma-type bonding. The -bonding orbitals on the x- and z-axes have to be considered separately? Why? These are termed as  (px) and  (pz) Examine Table 10-9. What is the symmetry of a square-planar complex?

SQUARE-PLANAR COMPLEXES SIGMA-TYPE BONDING ONLY
Finding the LGOs. red = A1g + B1g + Eu What are the orbitals on the central metal atom that can interact with these LGOs? Inspecting the character table reveals that the metal d-orbitals are split into three representations. Why? Examine Figure The energy difference between the eg/b2g nonbonding orbitals and the a1g antibonding is .

SQUARE-PLANAR COMPLEXES INCLUDING PI-BONDING
px = A2g + B2g + Eu () What are the interacting orbitals on the metal? pz = A2u + B2u + Eg () The effective overlap of the p orbitals on the metal to form  bonds is small. Examine Figure

THE ‘SETS’ OF ORBITALS IN FIGURE 10-15
The 1st set contains bonding orbitals (mostly sigma). 8 electrons from the ligands largely fill these orbitals. The 2nd set contains 8 -donor orbitals of the ligands. This interaction is small and decreases the energy differences in orbitals the next higher set. The 3rd set is primarily metal d-orbitals with some modifications due to interactions with the ligands. 3, 2, and 1 are in this set. The 4th set largely originates from the * orbitals of the ligands (if present). One of the main effects of these orbitals is the increase in the gap energy labeled 1.

ANGULAR OVERLAP (CRYSTAL FIELD)
Estimates the strength of interaction between individual ligand orbitals and d-orbitals based on the overlap between them. These values are then combined for all ligands and d-orbitals. The value for a given d-orbital is the sum of the numbers for the appropriate ligands in a column. This number can be positive or negative depending on location of the ligand and d-orbitals. The value for a given ligand is the sum of the numbers for all d-orbitals in the row. This number can also be positive or negative depending on location of the ligand and d-orbitals.

ANGULAR OVERLAP Sigma-donor interaction (no pi-orbitals are available). [M(NH3)6]n+ The strongest interaction is between the metal dz2 orbital and a ligand p-orbital (or appropriate MO). Describe the interaction based on this method. Table and Figure

ANGULAR OVERLAP Pi-acceptor ligands (available -type orbitals).
Strongest interaction is between dxz and * on the ligand. The * orbitals are almost always higher in energy. Reverse the signs. Figure and Table 10-12 There is a lowering of 4e due to this interaction. Why is magnitude e always smaller than that of e? Understand -donor interactions.

SAMPLE PROBLEM Using the angular overlap model, determine the splitting pattern of the d orbitals for a tetrahedral complex of formula ML4.where L is a capable of  interactions only.

SAMPLE PROBLEM Determine the energies of the d orbitals predicted by the angular overlap model for square planar complexes a) considering  interactions only b) considering both -donor and -acceptor interactions

THE SPECTROCHEMICAL SERIES
 depends on the relative energies and the degree of overlap. How ligands effect  -donor ligands -donating -accepting (or back bonding) Understand the spectrochemical series (page 368)

MAGNITUDE OF E, E, AND  Changing the metal and/or ligand effects the magnitudes of e and e, thereby changing the value of . Aqua species of Co2+ and Co3+ [Fe(H2O)6]2+ versus [Fe(H2O)6]3+ Tables and (Angular Overlap) e > e (always) Values decrease with increasing size and decreasing electronegativity Negative values for e. Why?

THE JAHN TELLER EFFECT There cannot be unequal occupation of orbitals with identical energies. The molecule will distort so that these orbitals are no longer degenerate. Cu(II) d9 ion, The complex will distort. How? The low-spin Cr(II) complex is octahedral with tetragonal distortion (Oh  D4h) Two absorption bands are observed instead of one.

DETERMINING FOUR- AND SIX-COORDINATE PREFERENCES
General angular overlap calculations of the energies expected for different number of d electrons and different geometries can give us some indication of relative stabilities. Larger number of bonds usually make the octahedral complexes more stable. Why are the energies equal in the d5, d6, and d7 cases? Figure

DETERMINING FOUR- AND SIX-COORDINATE PREFERENCES
The success of these simplistic calculations is variable. The s- and p-orbitals of the metal are not included. No -type interactions are included in Figure The orbital potential energies for the metals change with increasing atomic number (more negative). Can add –0.3e  (increase in Z) as a rough correction to the total enthalpy.

THE PROCESS FOR A COMPLEX OF D3h SYMMETRY
Construct the sigma-type bonding LGOs for the complex. Determine the interacting orbitals on the center atom. Construct a table to determine e (and e if appropriate). Construct the MO diagram and overlap energy figure. Homework: Determine the e contribution.

Symmetry and Group Theory
The symmetry properties of molecules and how they can be used to predict vibrational spectra, hybridization, optical activity, etc.

POINT GROUPS Molecules are classified and grouped based on their symmetry. Molecules with similar symmetry are but into the same point group. A point group contains all objects that have the same symmetry elements.

SYMMETRY ELEMENTS Symmetry elements are mirror planes, axis of rotation, centers of inversion, etc. A molecule has a given symmetry element if the operation leaves the molecule appearing as if nothing has changed (even though atoms and bonds may have been moved.)

Symmetry Elements Element Symmetry Operation Symbol Identity E n-fold axis Rotation by 2π/n Cn Mirror plane Reflection σ Center of inversion Inversion i n-fold axis of Rotation by 2π/n Sn improper rotation followed by reflection perpendicular to the axis of rotation

C3 C3 or three-fold rotational axis of the
ammonia molecule. If we rotate the ammonia molecule by 360/3 or 120º about this axis, its appearance is unchanged.

Rotational axes of BF3 . principal axis (highest value of Cn)
C C C C2 . three-fold axis three-fold axis two-fold axis two-fold axis viewed from viewed from viewed from viewed from above the side the side above Note: there are 3 C2 axes

How many axes of rotation does borazine possess?
SAMPLE PROBLEM How many axes of rotation does borazine possess? Ethane in the eclipsed conformation?

Mirror planes (σ) of BF3:
Mirror planes can contain the principal axis (σv) or be at right angles to it (σh). BF3 has one σh and three σv planes: (v = vertical, h = horizontal) σv mirror plane σh mirror plane C3 principal axis C3 principal axis σv mirror plane contains the C3 axis σh mirror plane is at right angles to the C3 axis

SAMPLE PROBLEM Mirror planes of symmetry for Borazine, naphtlalene, diborane, dxy orbital?

center of symmetry center of symmetry (Note: The center of symmetry is important in deciding whether orbitals are g or u (lecture 2.))

SAMPLE PROBLEM Which of the following flourine compounds has center of inversion? BF3, SiF4, PF5, XeF5-, SF6, C2F4,

The S4 improper rotation axis here is also a C2 axis
rotate by 360o/4 The S4 improper rotation axis here is also a C2 axis

Rotational axes and mirror planes of the water molecule:
principal axis σv mirror plane C2 σv mirror plane C2 The water molecule has only one rotational axis, its C2 axis, which is also its principal axis. It has two mirror planes that contain the principal axis, which are therefore σv planes. It has no σh mirror plane, and no center of symmetry.

Rotational axes and mirror planes of benzene C2 C2
principal axis C2 C2 C2 C2 C6 σh σv σv C6 principal axis C6 principal axis

Rotational axes and mirror planes of boron trifluoride
C3 principal axis C2 C2 C2 σh σv σv σh boron trifluoride has a C3 principal axis and three C2 axes, a σh mirror plane three σv mirror planes, but no center of inversion C3 principal axis

Identity, E All molecules have Identity. This operation leaves the entire molecule unchanged. A highly asymmetric molecule such as a tetrahedral carbon with 4 different groups attached has only identity, and no other symmetry elements.

Improper Rotation An improper rotation is rotation, followed by reflection in the plane perpendicular to the axis of rotation.

Improper Rotation The staggered conformation of ethane has an S6 axis that goes through both carbon atoms.

Improper Rotation Note that an S1 axis doesn’t exist; it is same as a mirror plane.

Improper Rotation Likewise, an S2 axis is a center of inversion.

Sample problem Draw the structure for the following showing the correct geometry and identify all the symmetry elements present in each: a) SCN- b) S2O32-, c) IF4- d) 1,8-dichloronaphthalene e) formaldehyde

Point Groups Molecules with the same symmetry elements are placed into point groups. Group theory, the mathematical treatment of the properties of groups can be used to determine the molecular orbitals, vibrations, and other properties of the molecule.

∞ ∞

Point Groups In general, you will not need to assign a molecule to its point group. Recognition of the features of some common point groups is useful.

Point Groups Water and ammonia both belong to the Cnv class of molecules. These have vertical planes of reflection, but no horizontal planes.

Point Groups Y The Dnh groups have a horizontal plane in addition to vertical planes. Many inorganic complexes belong to these symmetry groups. X X X X Y

POINT GROUPS Highly symmetrical molecules, such as identically substituted tetrahedrons or octahedrons belong to their own point groups (Td or Oh respectively).

Point Groups In assigning a point group, we typically ignore the fine detail, such as conformation isomers, of the ligands. In working problems using group theory, the point group of the molecule will usually be provided to you.

Example: PF5, SF6, IOF3, XeF4, ethane (eclipsed and staggered), ethylene and chloroethylene. Ferrocene (eclipsed and staggered)

COORDINATION CHEMISTRY III: REACTIONS OF METAL COMPLEXES
The ability to predict products and choose appropriate reaction condition to obtain the desired products is still a matter of art as well as science.

substitution reactions
kinetic consequences of reaction pathways experimental evidence in octahedral substitution substitution reactions of square-planar complexes the trans effect oxidation-reduction reactions reactions of coordinated ligand

SUBSTITUTION REACTIONS

SUBSTITUTION REACTIONS
Labile complexes <==> Fast substitution reactions (< few min) Inert complexes <==> Slow substitution reactions (>h) a kinetic concept Not to be confused with stable and unstable (a thermodynamic concept DGf <0)

INERT, LABILE vs STABLE, UNSTABLE
SUBSTITUTION REACTIONS – INERT AND LABILE INERT, LABILE vs STABLE, UNSTABLE kinetic terms thermodynamic terms Stable but labile unstable but inert

Mechanisms of ligand EXCHANGE REACTIONS IN OCTAHEDRAL COMPLEXES
Dissociative (D) Associative (A) Interchange (I) Ia if association is more important Id if dissociation is more important

OF DISSOCIATIVE REACTIONS
KINETICS OF DISSOCIATIVE REACTIONS

of interchange reactions
Kinetics of interchange reactions Fast equilibrium K1 = k1/k-1 k2 << k-1 For [Y] >> [ML5X]

Kinetics of associative reactions

Principal mechanisms of ligand exchange in octahedral complexes
Dissociative Associative

(5-coordinated intermediate)
MOST COMMON Dissociative pathway (5-coordinated intermediate) Associative pathway (7-coordinated intermediate)

Experimental evidence for dissociative mechanisms
Rate is independent of the nature of L

Experimental evidence for dissociative mechanisms
Rate is dependent on the nature of L

Inert and labile complexes
Some common thermodynamic and kinetic profiles Exothermic (favored, large K) Large Ea, slow reaction Exothermic (favored, large K) Large Ea, slow reaction Stable intermediate Endothermic (disfavored, small K) Small Ea, fast reaction

Labile or inert? LFAE = LFSE(sq pyr) - LFSE(oct)

Why are some configurations inert and some are labile?

Other metal on factors that affect reaction rates
Oxidation state of the central atom: Central atom with higher oxidation states have slower ligand exchange rates [AlF6]- > [SiF6]- > [PF6]- > SF6 Ionic radius. Smaller ions have slower exchange rates [Sr(H2O)6]2+ > [Ca(H2O)6]2+ > [Mg(H2O)6]2+ 112 pm pm 66 pm Both effects due to higher electrostatic attraction between central atom and attached ligands.

Substitution reactions in square-planar complexes
the trans effect (the ability of T to labilize X)

Synthetic applications
of the trans effect

Mechanisms of ligand exchange reactions in square planar complexes

THE trans EFFECT SIGMA-BONDING EFFECTS
Sigma-Bonding Effect. A strong  bond between Pt and T weakens the Pt-X bond. H- > PR3 > SCN- ~ CH3- ~ CO ~ CN- > Br- > Cl- > NH3 > OH-

PI-BONDING EFFECTS If back donation occurs to a ligand, the flow of electron density from the metal leaves less electron density to be donated in the opposite direction. C2H4 ~ CO > CN- > NO2- > SCN- > I- > Br- > Cl- > NH3 > OH-

Overall trans effect: CO ~ CN- ~ C2H4 > PR3 ~ H- > CH3- ~ SC(NH2)2 > C6H5- >NO2- ~ SCN ~ I- >Br- > Cl- > py , NH3 ~ OH- ~H2O

SAMPLE PROBLEM: Predict the products of the reactions (there may be one product when there are conflicting preferences) [PtCl4-] NO2- → (a) (a) + NH3 → (b) [PtCl3NH3] O2- → (c) (c) NO2- → (d)

SAMPLE PROBLEM: Is it possible to prepare different isomers of Pt(II) complexes with 4 different ligands? Predict the products expected if 1 mole of [PtCl4]- is reacted successively with the following reagents: (the product of reaction a is used in reaction b) 2 moles NH3 2 moles py 2 moles Cl- 1 mole NO2-

Electron transfer (redox) reactions
-1e (oxidation) M1(x+)Ln + M2(y+)L’n M1(x +1)+Ln + M2(y-1)+L’n +1e (reduction) Very fast reactions (much faster than ligand exchange) May involve ligand exchange or not Very important in biological processes (metalloenzymes)

REDOX MECHANISMS: Inner sphere mechanism: When two molecules are connected by a common ligand which the electron is transferred, in which case the reaction is called bridging or innersphere reaction. Outer sphere mechanism: Exchange may occur between two separate coordination sphere in outersphere reaction.

Outer sphere mechanism
[Fe(CN)6] [IrCl6]2- [Fe(CN)6] [IrCl6]3- [Co(NH3)5Cl] [Ru(NH3)6]2+ [Co(NH3)5Cl]+ + [Ru(NH3)6]3+ Reactions ca times faster than ligand exchange (coordination spheres remain the same) r = k [A][B] Tunneling mechanism

Inner sphere mechanism
[Co(NH3)5Cl)] [Çr(H2O)6]2+ [Co(NH3)5Cl)]2+:::[Çr(H2O)6]2+ [Co(NH3)5Cl)]2+:::[Çr(H2O)6]2+ [CoIII(NH3)5(m-Cl)ÇrII(H2O)6]4+ [CoII(NH3)5(m-Cl)ÇrIII(H2O)6]4+ [CoIII(NH3)5(m-Cl)ÇrII(H2O)6]4+ [CoII(NH3)5(m-Cl)ÇrIII(H2O)6]4+ [CoII(NH3)5(H2O)]2+ + [ÇrIII(H2O)5Cl]2+ [CoII(NH3)5(H2O)]2+ [Ço(H2O)6] NH4+

Inner sphere mechanism
Reactions much faster than outer sphere electron transfer (bridging ligand often exchanged) r = k’ [Ox-X][Red] k’ = (k1k3/k2 + k3) Tunneling through bridge mechanism

COORDINATION CHEMISTRY STRUCTURES AND ISOMERS

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COORDINATION CHEMISTRY STRUCTURES AND ISOMERS

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