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Coordination Compounds. Master Werner's Theory, VBT, CFT, Isomerism, and IUPAC Nomenclature.
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Basic Terms & Werner's Theory
Double Salts: These are addition compounds that completely dissociate into their constituent simple ions when dissolved in water (e.g., Mohr's salt, Potash alum). They lose their solid-state identity in solution.
Complex Salts: These compounds contain complex ions that do not dissociate into simple ions in water (e.g., $K_4[Fe(CN)_6]$). They retain their identity in both solid and solution states.
- Primary Valency: It corresponds to the oxidation state of the central metal ion. It is ionizable and is satisfied strictly by negative ions.
- Secondary Valency: It corresponds to the coordination number of the metal. It is non-ionizable, directional in space (giving geometry), and can be satisfied by negative ions or neutral molecules.
Coordination Entity: It constitutes a central metal atom/ion bonded to a fixed number of oppositely charged ions or neutral molecules enclosed in square brackets, e.g., $[Co(NH_3)_6]^{3+}$.
Coordination Number (CN): The total number of ligand donor atoms to which the central metal is directly bonded via coordinate covalent bonds.
Ligands are atoms, ions, or molecules that can donate at least one pair of electrons to the central metal atom/ion to form a coordinate bond.
Types by charge:
1. Neutral: $H_2O$ (aqua), $NH_3$ (ammine), $CO$ (carbonyl).
2. Anionic: $Cl^-$ (chlorido), $CN^-$ (cyanido).
3. Cationic: $NO^+$ (nitrosonium).
Denticity refers to the number of ligating (donor) groups in a single ligand molecule that simultaneously bind to a central metal ion.
A multidentate ligand can bind through two or more donor atoms simultaneously. Examples: ethylenediamine (bidentate), EDTA (hexadentate).
Ambidentate ligands are unidentate ligands that contain more than one coordinating atom, but they can coordinate to the central metal ion through only one of them at a time.
Examples:
1. $NO_2^-$ (can link through N as nitrito-N, or through O as nitrito-O).
2. $SCN^-$ (can link through S as thiocyanato-S, or through N as thiocyanato-N).
When a multidentate ligand coordinates to a central metal ion using two or more donor atoms, it forms a cyclic ring structure. Such cyclic complexes are thermodynamically much more stable than non-cyclic complexes formed by similar unidentate ligands. This enhanced stability is called the chelate effect.
Let oxidation number of Co be $x$.
'en' (ethylenediamine) is a neutral ligand, charge = 0.
'Cl' is chlorido, charge = -1.
$x + 2(0) + 2(-1) = +1 \implies x - 2 = +1 \implies x = \mathbf{+3}$.
Homoleptic: Complexes in which the central metal is bound to only one kind of donor group/ligand (e.g., $[Co(NH_3)_6]^{3+}$).
Heteroleptic: Complexes in which the central metal is bound to more than one kind of donor group/ligand (e.g., $[Co(NH_3)_4Cl_2]^+$).
Primary Valency (Oxidation State): $4(+1) + x + 6(-1) = 0 \implies x = +2$. Primary valency is 2.
Secondary Valency (Coordination Number): There are 6 unidentate $CN^-$ ligands bonded to Fe. Secondary valency is 6.
IUPAC Nomenclature & Isomerism
Diamminechloridonitrito-N-platinum(II).
Rules: Name ligands alphabetically (ammine, chlorido, nitrito). Oxidation state of Pt is +2. The complex is neutral, so the metal name is unchanged.
Potassium trioxalatoferrate(III).
Rules: Cation is named first. The complex is anionic, so the metal suffix becomes "-ate" (ferrate). Oxalate is a bidentate ligand, so we use 'tri' for three of them. Oxidation state of Fe is +3.
Tris(ethane-1,2-diamine)cobalt(III) sulphate.
Rules: 'en' has a numerical prefix in its name, so we use 'tris'. The complex is a cation. The external anion is simply 'sulphate' (no numerical prefix needed for counter ions).
Tetracarbonylnickel(0).
Rules: Carbon monoxide acts as a neutral ligand called 'carbonyl'. Since the complex is neutral, the oxidation state of Ni is strictly zero, written as (0).
$K_2[Ni(CN)_4]$.
Reasoning: Nickel is +2, four cyanido ligands bring a total charge of -4. The complex sphere has a net charge of $(+2) + (-4) = -2$. To balance this, we need two Potassium ($K^+$) ions.
Ionization isomerism occurs when the counter ion in a complex salt acts as a ligand and displaces a ligand from within the coordination sphere, producing different ions in aqueous solution.
Example: $[Co(NH_3)_5Br]SO_4$ (red-violet, yields $SO_4^{2-}$ ions) and $[Co(NH_3)_5SO_4]Br$ (red, yields $Br^-$ ions).
Linkage isomerism arises exclusively in coordination compounds containing ambidentate ligands. It occurs when the same ligand coordinates to the metal through different donor atoms.
Example: $[Co(NH_3)_5(NO_2)]^{2+}$ features a metal-N bond (yellow complex), while its linkage isomer $[Co(NH_3)_5(ONO)]^{2+}$ features a metal-O bond (red complex).
This is a square planar $MA_2B_2$ complex showing Cis-Trans isomerism.
Cis-isomer: The two $Cl^-$ and two $NH_3$ are adjacent (90°). Known as Cisplatin, it is a highly effective anti-cancer drug used in chemotherapy.
Trans-isomer: The identical ligands are opposite each other (180°). It has no anti-cancer properties.
The $[Co(en)_3]^{3+}$ complex is an octahedral complex containing three bidentate ligands. Its spatial arrangement completely lacks a plane of symmetry. Therefore, its mirror image is non-superimposable on the original molecule, making it a chiral molecule that exists as a pair of enantiomers (dextro and laevo rotatory).
They occur in octahedral complexes of the type $MA_3B_3$ (e.g., $[Co(NH_3)_3(NO_2)_3]$).
Facial (fac): Three identical ligands occupy the corners of one triangular face of the octahedron.
Meridional (mer): Three identical ligands are arranged around the meridian (equator) of the octahedron, meaning two are trans (180°) to each other.
Valence Bond Theory (VBT)
VBT postulates that the central metal ion uses empty s, p, and d orbitals, which undergo hybridization to form a set of equivalent, directional hybrid orbitals. These empty hybrid orbitals overlap with the filled orbitals of the ligands (containing lone pairs) to form coordinate covalent bonds.
In octahedral geometry (CN = 6):
Inner orbital complex: The metal utilizes inner $(n-1)d$ orbitals for hybridization ($d^2sp^3$). These are typically formed with strong field ligands that force electron pairing.
Outer orbital complex: The metal utilizes outer $nd$ orbitals for hybridization ($sp^3d^2$). These are typically formed with weak field ligands.
By observing the electronic configuration of the metal after hybridization. If all electrons are paired, the complex is diamagnetic. If there are one or more unpaired electrons left in the d-orbitals, the complex is paramagnetic.
$Co$ is in +3 state, giving a $3d^6$ configuration. Ammonia ($NH_3$) acts as a strong ligand for $Co^{3+}$, forcing the 6 electrons to pair up against Hund's rule, filling only three 3d orbitals. This leaves two 3d orbitals completely empty. The metal undergoes $d^2sp^3$ hybridization. Since no unpaired electrons remain, it is diamagnetic.
$Co$ is in +3 state ($3d^6$). Fluoride ($F^-$) is a weak field ligand and cannot force pairing. The electrons remain unpaired (4 unpaired electrons). To accommodate 6 ligands, Co uses its outer 4d orbitals, resulting in $sp^3d^2$ hybridization. Due to 4 unpaired electrons, it is strongly paramagnetic.
In both, Ni is +2 ($3d^8$).
In $[Ni(CN)_4]^{2-}$, $CN^-$ is a strong ligand, causing the 8 electrons to pair up completely, leaving one 3d orbital empty for $dsp^2$ (square planar) hybridization. No unpaired electrons = diamagnetic.
In $[NiCl_4]^{2-}$, $Cl^-$ is a weak ligand. Electrons do not pair up (2 unpaired electrons remain). It undergoes $sp^3$ (tetrahedral) hybridization. Unpaired electrons = paramagnetic.
- $sp^3$ Hybridization: Tetrahedral shape (e.g., $[Zn(NH_3)_4]^{2+}$).
- $dsp^2$ Hybridization: Square Planar shape (e.g., $[PtCl_4]^{2-}$).
- It cannot explain the vividly distinct colors of coordination compounds.
- It does not give a quantitative interpretation of magnetic data.
- It cannot clearly explain or predict whether a specific ligand will behave as strong (forcing pairing) or weak (no pairing).
Low Spin (Spin Paired): Formed when strong field ligands force electron pairing, minimizing the number of unpaired electrons. Usually corresponds to inner orbital ($d^2sp^3$) complexes.
High Spin (Spin Free): Formed when weak field ligands do not force pairing, maximizing the number of unpaired electrons. Usually corresponds to outer orbital ($sp^3d^2$) complexes.
Fe is in +3 state, giving $3d^5$ configuration. $H_2O$ is a weak field ligand, so no pairing occurs. There are 5 unpaired electrons ($n=5$).
$\mu = \sqrt{n(n+2)} = \sqrt{5(5+2)} = \sqrt{35} \approx \mathbf{5.92 \text{ B.M.}}$
Crystal Field Theory (CFT)
CFT treats the metal-ligand bond as purely ionic (electrostatic) interactions. It considers ligands as point charges (or point dipoles). When these negatively charged ligands approach the central metal ion, they cause repulsion with the metal's d-electrons, breaking the degeneracy (equal energy) of the five d-orbitals.
In an octahedral field, six ligands approach along the x, y, and z axes. The two d-orbitals lying exactly on these axes ($d_{x^2-y^2}$, $d_{z^2}$) experience maximum repulsion and their energy is raised, forming the $e_g$ set. The three orbitals lying between the axes ($d_{xy}$, $d_{yz}$, $d_{zx}$) experience less repulsion, dropping in energy to form the $t_{2g}$ set.
The spectrochemical series is an experimentally determined sequence in which ligands are arranged in increasing order of their crystal field splitting power ($\Delta$).
Halides are weak field ligands ($I^- < Br^- < Cl^- < F^-$).
Carbon and Nitrogen donors are strong field ligands ($CN^- \approx CO > en > NH_3$).
For a $d^4$ ion, the first 3 electrons fill the $t_{2g}$ level. The 4th electron has a choice:
1. If $\Delta_o < P$ (Weak field ligand): It jumps to the $e_g$ level. Configuration is $t_{2g}^3 e_g^1$ (High spin).
2. If $\Delta_o > P$ (Strong field ligand): The gap is too large to cross. It pairs up in $t_{2g}$. Configuration is $t_{2g}^4 e_g^0$ (Low spin).
In a tetrahedral field, ligands approach between the axes. This completely reverses the splitting pattern seen in octahedral fields. The orbitals between the axes ($d_{xy}$, $d_{yz}$, $d_{zx}$) experience more repulsion and rise in energy to form the $t_2$ set, while the axial orbitals drop to form the $e$ set. Also, $\Delta_t \approx \frac{4}{9} \Delta_o$.
Because there are only 4 ligands (instead of 6) and they do not point directly at any d-orbitals, the splitting energy in tetrahedral complexes ($\Delta_t$) is inherently very small. Since $\Delta_t$ is almost always less than the pairing energy ($P$), electrons will readily occupy the higher $t_2$ orbitals rather than pairing up.
According to CFT, light absorption promotes an electron from the lower energy d-orbital set ($t_{2g}$) to the higher energy d-orbital set ($e_g$). This is called a d-d transition. The complex absorbs a specific wavelength of visible light to fuel this jump, and what we see is the complementary color of the absorbed light.
In the hydrated complex, the water ligands cause crystal field splitting, allowing d-d transitions that absorb yellow-green light and transmit violet. On heating, the water ligands evaporate, destroying the crystal field. The five d-orbitals become degenerate (equal energy) again, making d-d transitions impossible, hence it becomes colorless.
Zinc is in the +2 oxidation state with a $3d^{10}$ configuration. Its d-orbitals are completely full. Because there are no vacant spaces in the higher energy d-orbitals ($e_g$ or $t_2$), a d-d electron transition is physically impossible. Without light absorption for transition, the complex appears colorless.
- It treats ligands strictly as point charges, completely ignoring any covalent character in the metal-ligand bond.
- If it were purely electrostatic, anionic ligands ($OH^-, Cl^-$) should create the strongest field. Ironically, they are at the bottom of the spectrochemical series, while neutral molecules like $CO$ are at the top.
Stability & Applications
Complexes are formed in a stepwise manner, each with a stepwise stability constant ($K_1, K_2...$). The overall stability constant ($\beta_n$) is the equilibrium constant for the net formation reaction. It is equal to the product of all the stepwise stability constants: $\beta_n = K_1 \times K_2 \times ... \times K_n$. A higher $\beta$ means a more stable complex.
Stability is directly proportional to the charge density of the central metal ion. A higher positive charge and a smaller ionic radius result in a stronger electrostatic pull on the incoming ligands, leading to a much more stable complex.
This is driven by thermodynamics, specifically entropy ($\Delta S$). When a multidentate ligand displaces several unidentate ligands (like water) from the metal, the total number of free molecules in the solution increases. This massive increase in randomness/entropy makes the formation of the chelate ring extremely favorable.
In the Macarthur-Forrest cyanide process, finely powdered Gold or Silver ore is leached with a dilute solution of $NaCN$ in the presence of air. The metal dissolves by forming a highly stable, water-soluble coordination complex like $[Au(CN)_2]^-$. The pure metal is later recovered by displacement with Zinc.
The compound is Cisplatin, chemically known as cis-diamminedichloridoplatinum(II) or $cis\text{-}[Pt(NH_3)_2Cl_2]$. It binds to the DNA of rapidly dividing cancer cells, inhibiting their replication.
Hardness of water is caused by $Ca^{2+}$ and $Mg^{2+}$ ions. EDTA (Ethylene Diamine Tetraacetic Acid) is a hexadentate ligand that forms highly stable, colorless chelate complexes with these ions. A simple volumetric titration using a standard EDTA solution (with an indicator like Eriochrome Black T) accurately determines their concentration.
- Chlorophyll (essential for photosynthesis) is a coordination complex of Magnesium ($Mg^{2+}$).
- Hemoglobin (oxygen carrier in blood) is a coordination complex of Iron ($Fe^{2+}$).
- Vitamin B12 (cyanocobalamin) is a coordination complex of Cobalt ($Co^{3+}$).
In metal carbonyls, the metal-carbon bond possesses both $\sigma$ and $\pi$ character. First, CO donates its lone pair into an empty d-orbital of the metal to form a $\sigma$ bond. Simultaneously, the filled d-orbitals of the metal back-donate electron density into the empty anti-bonding $\pi^*$ orbital of CO. This mutual reinforcement heavily strengthens the bond, known as the synergic effect.
Both are mononuclear metal carbonyls with the metal in a zero oxidation state.
$Ni(CO)_4$ is completely tetrahedral (using $sp^3$ hybridization).
$Fe(CO)_5$ assumes a trigonal bipyramidal structure (using $dsp^3$ hybridization).
Wilkinson's catalyst is a coordination complex of rhodium: Chloridotris(triphenylphosphine)rhodium(I), formula $[RhCl(PPh_3)_3]$.
It is extensively used industrially as a highly efficient homogeneous catalyst for the hydrogenation of alkenes to alkanes at room temperature and pressure.
Chapter 5 Mastered!
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