Chapter 20. Coordination Chemistry (Inorganic Chemistry, Coordination Chemistry)
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4. Example : Synthesis Process of Platinum Coordination Compounds
1. Transition Elements
⑴ Definition : Elements with outermost orbital being d orbital, from Group 3 to Group 12, also known as d-block elements.
① Essential Memorization Transition Elements : Scandium, Titanium, Vanadium, Chromium, Manganese, Iron, Cobalt, Nickel, Copper, Zinc
⑵ Characteristics
① Varied oxidation states (exception : Zinc has only one oxidation state)
② Good at forming coordination ions
⑶ Types of d orbitals : Five types in total
Figure 1. Types of d orbitals
Since ℓ = 2, they all have two nodal planes
2. Coordination Ions
⑴ Coordination Ions : Ions formed by the combination of a metal ion and ligands
Mn+ : Central metal
xL : Ligand
MLxn+ : Coordination ion
① Coordination number = Number of bonds formed by the central metal = Steric number = Number of orbitals participating in hybridization
② Oxidation state of central metal = Charge of coordination compound - Sum of charges of ligands coordinated
○ Oxidation state of central metal determines the number of electrons forming the MOs of the coordination compound
③ Ligand in slight excess, excess
○ With a slight amount of ligand, the formation constant for complex formation is satisfied even with a small amount of metal ion reacting
○ With an excess of ligand, the concentration of metal ion becomes 0 for the complex formation constant equation to be satisfied
⑵ Formation constant Kf : Equilibrium constant for the formation of coordination ions
① Ionization of complex compounds : Generally, ligands do not ionize in solution (∵ Kf ≫ 1)
② Size comparison : The stronger the bond between the central metal and the ligand, the larger the formation constant
○ Difference in formation constants is mainly due to differences in entropy change
○ Chelates have larger formation constants
○ Common ligands ＜ Ethylenediamine (en) ＜ Diethylenetriamine (trien)
○ Larger crystal field splitting leads to larger formation constants
○ Exception : Formation constant of OH- for Zn2+ is larger than that of NH3
⑶ Chelate (Chelator) : Ligand that binds to multiple sites simultaneously
① Major chelates
○ (Note) The 1st coordination in the list is included for reference and is not a chelating ligand
○ 1st coordination : H2O, CN-, SCN- (thiocyanate), X- (halogens), NH3, NO2- (nitrite), OH-, etc.
○ 2nd coordination : Oxalate ion, ethylenediamine (en), ethyleneglycol, acac, CO32-, etc.
○ Oxalate : HO2CCO2H
○ Ethylenediamine : H2NCH2CH2NH2
○ Ethyleneglycol : HO(CH2)2OH
○ Acac : (CH3COCHCOCH3)-
○ (Note) m -phenylenediamine is not a 2nd coordination chelate
○ 3rd coordination : Diethylenetriamine (dien)
○ 6th coordination : Ethylenediaminetetraacetic acid (EDTA)
○ EDTA = (HOOCCH2)2NCH2CH2N(CH2COOH)2
② Chelate effect
○ Definition : When there are two competing reactions to form equivalent bonds, the one involving the chelating ligand occurs more easily
○ In the case of a 1st coordination ligand binding to a metal ion in a 6-fold manner, the total number of degrees of freedom is reduced from 7 to 1, leading to an unfavorable entropy change
○ Chelate ion formation reactions are favorable in terms of entropy since they require fewer ligands
○ The formation constant of a chelate is larger than that of a non-chelating ligand
⑷ Hybrid Orbitals
① To understand hybrid orbitals, one must know the orbital arrangement of the central metal ion
○ Note : According to the Aufbau principle, the 4s orbital is filled before the 3d orbital because it is lower in energy. However, once orbitals are formed, the 3d orbital becomes lower than the 4s orbital, and electrons are removed from the 4s orbital
○ Example : The electron configuration of Fe is 1s22s22p63s23p64s23d6, and the electron configuration of Fe2+ is 1s22s22p63s23p63d6.
○ Why it’s necessary to determine the oxidation state of the central metal ion in a given situation
② 2nd coordination
○ sp hybridization : Linear. [Ag(NH3)2]+, etc.
③ 4th coordination
○ sp3 hybridization : Tetrahedral. [Zn(NH3)4]2+, etc.
○ dsp2 hybridization : Square planar. [Ni(CN)4]2-, [Pt(NH3)4]2+, mostly d8 complexes
④ 6th coordination
○ sp3d2 (d2sp3) hybridization : Octahedral. XeF4, [Co(NH3)6]3+, [PtCl6]2+, etc.
⑸ Isomerism : Compounds with the same chemical formula but different atomic arrangements
① Structural Isomerism : Compounds with different bonds between atoms
○ Linkage Isomerism : Isomers with different metal-ligand bonds
○ Ionization Isomerism : Isomers that form different ions in solution
② Stereoisomerism : Compounds where atoms are bonded in the same way but have different spatial arrangements
○ Type 1. Geometric Isomerism : cis/trans isomers, fac/mer isomers, etc.
○ Type 1-1. cis/trans Isomerism : If two substituents are on the same side, it’s cis; if on opposite sides, it’s trans
○ Type 1-2. fac/mer Isomerism
○ fac Isomerism : When three substituents are mutually orthogonal
○ (Note) fac comes from ‘facial,’ and it refers to three substituents lying in a single plane
○ mer Isomerism : When a pair of substituents are opposite each other
○ (Note) mer comes from ‘meridional,’ and it refers to a pair of substituents forming a meridian
○ Type 2. Optical Isomerism (Enantiomerism)
○ Example 1. FeCl3(H2O)3 (ferric chloride) has fac/mer isomerism. There’s a symmetry plane within the molecule, so there’s no optical isomerism
Figure 2. Stereoisomerism of FeCl3(H2O)3
○ Example 2. Co(en)2Cl2+ has two cis-Co(en)2Cl2+ and one trans-Co(en)2Cl2+ that are mirror images of each other
Figure 3. Stereoisomerism of Co(en)2Cl2+
○ Example 3. Co(en)2BrCl+ has two cis-Co(en)2BrCl+ and one trans-Co(en)2BrCl+ that are mirror images of each other
○ (Note) Observation reveals this
Figure 4. Stereoisomerism of Co(en)2BrCl+
○ Example 4. [Co(en)3]3+ with symmetric ligands attached in sets of 3 has mirror image isomers : There are a total of 2 stereoisomers
Figure 5. Stereoisomerism of [Co(en)3]3+
○ Example 5. [Co((R)-pn)3]3+ has four types of isomers: Λ-fac, Δ-fac, Λ-mer, Δ-mer
Figure 6. Stereoisomerism of [Co((R)-pn)3]3+
○ Example 6. [Fe(gly)3]3- with a chelating ligand gly at 2 positions (N-terminal and O-terminal) has four types of stereoisomers as in Example 4
3. Crystal Field Theory (CFT)
⑴ Definition : Theory that explains the color and magnetic properties of coordination ions through changes in the energy of d orbitals of the central metal ion forming coordination complexes
⑵ Crystal field splitting
① Definition : When the nonbonding electron pairs of ligands repel the d orbital electrons of the central metal ion, the energy levels of the d orbitals rise significantly. Depending on the distance between d orbitals and ligands, different energy differences emerge, causing the d orbitals to split
② The larger the repulsion, the more unstable it is, resulting in higher energy levels
⑶ Crystal field splitting according to structure
① Overview
○ The splitting of the crystal field affects the d orbital electron arrangement.
○ The d orbitals are important in determining paramagnetic / diamagnetic properties, so a good understanding of the crystal field splitting based on structure is necessary.
② Assumptions
○ Metal-ligand bonding is considered as complete ionic bonding.
○ The greater the overlap between ligands and orbitals, the higher the energy level of the d orbitals as the repulsion strength increases.
③ Octahedral Structure
○ Ligands approach from the x, y, and z axes.
○ Strong repulsion with dz2 and dx2-y2 orbitals, which have a high probability of finding electrons along the ligand’s negatively charged axis.
○ Comparison: d z2 = dx2-y2 > dxy = dyz = dxz (memorize)
○ Memorization tip: Think of it like fractions.
○ Jahn-Teller effect: In CrF2 crystals, the 6 Cr-F bonding pairs around Cr consist of 2 long and 4 short bonds.
**Figure 7. Crystal Field Theory for Octahedral Structures]
④ Tetrahedral Structure
○ Ligands approach the space between the x, y, and z axes.
○ Strong repulsion with dxy, dyz, and dxz orbitals along with dx2-y2 and dz2 orbitals.
○ Comparison: d xy = dyz = dxz > dz2 = dx2-y2 (memorize)
○ Memorization tip: Think of it like fractions.
**Figure 8. Crystal Field Theory for Tetrahedral Structures]
⑤ Square Planar Structure
○ Example: Cisplatin (cis-Pt(NH3)2Cl2)
○ Ligands approach from the x and y axes.
○ Strongest repulsion with the dx2-y2 orbital, which has higher electron density along the x and y axes.
○ Comparison: d x2-y2 > dxy > dz2 > dxz = dyz (memorize)
○ Note that dxz and dyz have lower energy levels than dz2.
**Figure 9. Crystal Field Theory for Square Planar Structures]
⑥ Linear Structure
○ Example: I3-
○ Repulsion increases as ligands approach along the z-axis, causing higher energy levels.
○ Comparison: dz2 > dxz = dyz > dxy = dx2-y2
○ (Note) dxz and dyz are always paired.
**Figure 10. Crystal Field Splitting of d Orbitals in Linear ML2 Complexes’ Central Metal Ion M]
㈎ : dz2, ㈏ : dyz, ㈐ : dx2-y2
⑷ Magnitude of Crystal Field Splitting
① Principles
○ The stronger the central metal-ligand bond, the greater the aggregation of ligand electrons towards the central metal, leading to increased crystal field splitting.
○ Larger crystal field splitting leads to increased formation constants.
② Factor 1: Ligand Field Theory (LFT)
○ Different ligands cause different degrees of bonding with the central metal, resulting in varying crystal field splitting sizes.
○ Difference from Crystal Field Theory
○ Crystal Field Theory understands metal-ligand interaction as electrostatic forces.
○ Ligand Field Theory considers metal-ligand interaction as covalent bonding.
○ Crystal Field Theory assumes constant d orbital energy levels based on ligand type.
○ Ligand Field Theory assumes variable d orbital energy levels based on ligand type.
○ Comparison
CO, CN- > NO2- > ONO- > en > NH3 > SCN- ＞ H2O > OH- > F- > Cl- > Br- > I-
○ Halide ions have increasing crystal field splitting as the period decreases, due to smaller radii leading to stronger electron repulsion.
○ First four (C, N, E, N) are essential to memorize.
○ Strong-field ligands, weak-field ligands
○ Exception: In the case of Zn2+ as the central metal, the formation constant of OH- is greater than NH3.
③ Factor 2: Central Metal Charge
○ Greater central metal charge leads to stronger attraction of electrons to the nucleus, resulting in increased repulsion and crystal field splitting.
○ Greater oxidation state of the central metal leads to larger crystal field splitting → shorter absorption wavelength.
○ Exception: CuCl2 is blue, while CuCl is white due to full occupancy of d orbitals reflecting all light.
④ Factor 3: Hybrid Orbital Structure
○ Octahedral structures have more electron repulsion than tetrahedral structures, leading to larger crystal field splitting.
○ Ligand Field Theory doesn’t apply here: Octahedral almost always has larger crystal field splitting than tetrahedral.
○ Octahedral structures absorb high-energy short wavelengths (visible and ultraviolet light).
○ Tetrahedral structures absorb lower-energy, relatively longer wavelengths.
○ Square planar structures absorb the shortest wavelengths. (ref)
⑤ Crystal Field Stabilization Energy (CFSE)
○ Definition: Energy stabilized when a metal ion is placed in a crystal field created by ligands in a complex compound.
○ Represented as Δ or Δo.
○ Appears in an octahedral structure as follows:
⑥ Pairing Energy (P)
⑸ Classification of sp3d2 Orbitals According to Crystal Field Splitting
① sp3d2 orbitals are divided into high-spin and low-spin configurations based on the size of crystal field splitting (Δ).
② High Spin:
○ High spin configurations maximize spin pairing in dxy, dyz, dxz orbitals and dx2-y2, dz2 orbitals.
○ Situation: P > Δ, corresponds to weak-field ligands.
○ In high-spin, the energy difference between eg and t2g is small, so both are considered at the same energy level, maximizing the placement of unpaired electrons.
○ Even in high-spin, the dx2-y2 and dz2 orbitals are occupied last due to energy level differences.
③ Low Spin:
○ Low spin configurations have fewer electron pairings due to larger differences between dxy, dyz, dxz orbitals and dx2-y2, dz2 orbitals.
○ Situation: P < Δ, corresponds to strong-field ligands.
○ Since eg > t2g is distinct, electrons are placed in t2g orbitals first.
④ Other:
○ K2[MnCl4] is square planar and high spin, resulting in 5 unpaired electrons.
⑹ Color of Complexes
① Larger crystal field splitting (Δ) results in shorter maximum absorption wavelength (λmax).
② The color of a solution is determined by the absorbed light’s complementary color.
○ Larger absorption wavelengths yield complementary colors closer to blue.
○ Smaller absorption wavelengths yield complementary colors closer to red.
○ Larger crystal field splitting → smaller absorption wavelengths → complementary colors determined using a color chart.
**Figure 11. Absorption Wavelength and Complementary Color]
③ Some complexes like Ni(H2O)62+ and Ni(en)32+ don’t follow the simple principle of complementary colors.
**Figure 12. Absorption Spectra of Ni(H2O)62+ and Ni(en)32+]
○ Ni(H2O)62+ appears green.
○ Ni(en)32+ appears blue.
○ Expected: Increased crystal field splitting (en > H2O) → decreased absorption wavelength (en < H2O) → increased observed wavelength (en > H2O).
○ Reality: Observed wavelength of Ni(en)32+ is shorter than Ni(H2O)62+.
○ Reason: Interference by other peaks in the case of Ni complexes.
○ In Ni(H2O)62+ absorption, another peak absorbs at 400 nm, resulting in an observed color around 500 nm.
○ In Ni(en)32+ absorption, less interference by other peaks, leading to an observed color around 450 nm according to the principle of complementary colors.
⑺ Magnetic Properties of Complexes
① Principle: Presence of unpaired electrons leads to paramagnetism, absence leads to diamagnetism.
○ Examples: d0 and d10 complexes are always diamagnetic.
○ Reason: Unpaired electrons are not possible, fundamentally due to the impossibility of t2g → eg electron transitions.
② 4-Coordinate Complexes
○ Square planar: Paramagnetic
○ Square planar (d8): Diamagnetic
③ 6-Coordinate Complexes
○ d6 complex + Low Spin: Diamagnetic
○ d6 complex + High Spin: Paramagnetic
○ Others: Paramagnetic
4. Example: Synthesis of Platinum Complex Compounds
⑴ Step 1: Synthesize complex compound A, a cis isomer, by adding a small amount of ammonia solution to a solution of potassium hexachloroplatinate (K2PtCl4).
⑵ Step 2: Synthesize complex compound B by adding excess ammonia solution to solution A.
⑶ Step 3: Synthesize complex compound C, a trans isomer, by adding a suitable amount of HCl(aq) to solution B.
Input: January 2, 2019, 19:42
Modified: June 3, 2023, 20:49