Comprehensive Notes on Coordination Compounds

Overview of Coordination Compounds

• Definition: Coordination compounds are substances in which metal atoms are bound to a number of anions or neutral molecules by the sharing of electrons. In modern terminology, these are referred to as coordination compounds.
• Importance: They represent a challenging area of modern inorganic chemistry, providing insights into chemical bonding and molecular structure.
• Biological Significant Compounds:     * Chlorophyll: A coordination compound of magnesium ($Mg$).     * Haemoglobin: A coordination compound of iron ($Fe$).     * Vitamin $B_{12}$ (Cyanocobalamine): A coordination compound of cobalt ($Co$).
• Applications: Used in metallurgy, industrial catalysts, analytical reagents, electroplating, textile dyeing, and medicinal chemistry.
• Foundations: Coordination compounds are considered the backbone of modern inorganic and bio-inorganic chemistry.

Werner’s Theory of Coordination Compounds

• Origin: Alfred Werner ($1866$–$1919$), a Swiss chemist, was the first to propose structures for coordination compounds. He received the Nobel Prize in $1913$ for his work on the linkage of atoms.
• Primary and Secondary Valence:     * Primary Valence: Generally ionisable and satisfied by negative ions. For binary compounds like $CrCl_3$, $CoCl_2$, or $PdCl_2$, the primary valences are $3$, $2$, and $2$ respectively.     * Secondary Valence: Non-ionisable. These are satisfied by neutral molecules or negative ions. The secondary valence is equal to the coordination number ($CN$) and is fixed for a particular metal.
• Werner’s Experimental Observations on Cobalt(III) Chloride Complexes:     By adding excess silver nitrate ($AgNO_3$) to a series of cobalt(III) chloride-ammonia complexes, different amounts of silver chloride ($AgCl$) precipitated, indicating different numbers of ionisable chloride ions:     * Yellow (1 mol\[Co(NH_3)_6\]Cl_3): Yields $3 mol$ of $AgCl$. Conductivity corresponds to a $1:3$ electrolyte.     * Purple (1 mol\[CoCl(NH_3)_5\]Cl_2): Yields $2 mol$ of $AgCl$. Conductivity corresponds to a $1:2$ electrolyte.     * Green (1 mol\[CoCl_2(NH_3)_4\]Cl): Yields $1 mol$ of $AgCl$. Conductivity corresponds to a $1:1$ electrolyte.     * Violet (1 mol\[CoCl_2(NH_3)_4\]Cl): Yields $1 mol$ of $AgCl$. Conductivity corresponds to a $1:1$ electrolyte. Green and Violet forms are isomers.
• Postulates of Werner’s Theory:     1. Metals show two types of linkages: primary and secondary.     2. Primary valences are ionisable and satisfied by negative ions.     3. Secondary valences are non-ionisable and satisfied by neutral molecules or negative ions. Secondary valence is the coordination number.     4. Groups bound by secondary linkages have characteristic spatial arrangements called coordination polyhedra.
• Modern Formulation: Species inside the square bracket are the coordination entity/complex. Ions outside are counter ions. Common shapes for transition metal complexes include octahedral, tetrahedral, and square planar.

Difference Between Double Salts and Complexes

• Double Salts: Formed by the combination of stable compounds in stoichiometric ratios. They dissociate completely into simple ions when dissolved in water. Examples: Carnallite ($KCl \times MgCl_2 \times 6H_2O$), Mohr’s salt ($FeSO_4 \times (NH_4)_2SO_4 \times 6H_2O$), and Potash alum ($KAl(SO_4)_2 \times 12H_2O$).
• Complex Ions: Also formed in stoichiometric ratios but do not dissociate into simple component ions. For example, the complex ion $[Fe(CN)_6]^{4-}$ of $K_4[Fe(CN)_6]$ does not dissociate into $Fe^{2+}$ and $CN^-$ ions.

Important Definitions in Coordination Chemistry

• Coordination Entity: A central metal atom or ion bonded to a fixed number of ions or molecules. Example: $[CoCl_3(NH_3)_3]$.
• Central Atom/Ion: The atom or ion to which ligands are bound in a definite geometry. These act as Lewis acids. Examples: $Ni^{2+}$ in $[NiCl_2(H_2O)_4]$.
• Ligands: The ions or molecules bound to the central atom.     * Unidentate: Bound through a single donor atom (e.g., $Cl^-$, $H_2O$, $NH_3$).     * Didentate: Can bind through two donor atoms (e.g., ethane-$1,2$-diamine ($H_2NCH_2CH_2NH_2$) or oxalate ($C_2O_4^{2-}$)).     * Polydentate: Several donor atoms are present. Example: ethylenediaminetetraacetate ($EDTA^{4-}$) is a hexadentate ligand binding via two nitrogen and four oxygen atoms.     * Chelate Ligand: A di- or polydentate ligand that uses two or more donor atoms simultaneously to bind a single metal ion. This creates a chelate complex, which is generally more stable. The number of such groups is the denticity.     * Ambidentate Ligand: Ligands with two different donor atoms that can coordinate at either site ($NO_2^-$ can bind via $N$ or $O$; $SCN^-$ can bind via $S$ or $N$).
• Coordination Number (CN): The number of ligand donor atoms directly bonded to the metal ion. It is determined specifically by the number of sigma bonds; pi bonds are not counted.
• Coordination Sphere: The central atom and attached ligands enclosed in square brackets. The ionisable groups outside are counter ions.
• Coordination Polyhedron: The spatial arrangement of the ligand atoms directly attached to the central atom ($octahedral$, $tetrahedral$, $square planar$).
• Oxidation Number: The charge the central atom would carry if all ligands and shared electron pairs were removed. Represented by a Roman numeral in parentheses.
• Homoleptic vs. Heteroleptic:     * Homoleptic: Metal bound to only one kind of donor group (e.g., $[Co(NH_3)_6]^{3+}$).     * Heteroleptic: Metal bound to more than one kind of donor group (e.g., $[Co(NH_3)_4Cl_2]^+$).

Nomenclature of Coordination Compounds

• Formulas of Mononuclear Entities:     1. Central atom is listed first.     2. Ligands are listed in alphabetical order (ignoring charge).     3. Polydentate ligands are also listed alphabetically (use the first letter of abbreviation).     4. The entire entity is enclosed in square brackets ($[ ]$).     5. Polyatomic ligands are enclosed in parentheses ($( )$).     6. No space between ligands and metal within the sphere.     7. Charge is indicated as a right superscript outside brackets (e.g., $[Co(CN)_6]^{3-}$).
• Naming Rules:     1. Cation is named first.     2. Ligands are named in alphabetical order before the metal.     3. Anionic ligand names end in $-o$ (e.g., $chlorido$). Neutral ligands: $aqua$ ($H_2O$), $ammine$ ($NH_3$), $carbonyl$ ($CO$), and $nitrosyl$ ($NO$).     4. Multipliers: $di-, tri-, tetra-$ etc. If ligand name contains a prefix, use $bis-, tris-, tetrakis-$ with ligand in parentheses.     5. Oxidation state in Roman numerals in parentheses.     6. If complex is a cation or neutral, use element name (e.g., cobalt). If it is an anion, use suffix $-ate$ (e.g., cobaltate); some use Latin roots (ferrate for iron).

Isomerism in Coordination Compounds

Isomers have the same chemical formula but different structural arrangements, leading to distinct physical or chemical properties.

Stereoisomerism

• Geometrical Isomerism: Arises in heteroleptic complexes with different spatial arrangements of ligands.     * Square Planar ($MX_2L_2$): $cis$ (adjacent) and $trans$ (opposite) isomers. Example: $[Pt(NH_3)_2Cl_2]$.     * Octahedral ($MX_2L_4$): Also shows $cis$ and $trans$.     * Octahedral ($Ma_3b_3$): Shows facial (fac) (three donor atoms on adjacent corners of an octahedral face) and meridional (mer) (positions around the meridian) isomers. Example: $[Co(NH_3)_3(NO_2)_3]$.     * Note: Geometrical isomerism is not possible in tetrahedral complexes as all relative ligand positions are identical.
• Optical Isomerism: Mirror images that are non-superimposable (enantiomers). Molecules are called chiral.     * Dextro (d): Rotates plane-polarised light to the right.     * Laevo (l): Rotates plane-polarised light to the left.     * Common in octahedral complexes with didentate ligands (e.g., $[Co(en)_3]^{3+}$ and $cis-[PtCl_2(en)_2]^{2+}$).

Structural Isomerism

• Linkage Isomerism: Occurs with ambidentate ligands ($SCN^-$ vs $NCS^-$; $NO_2^-$ vs $ONO^-$).
• Coordination Isomerism: Interchange of ligands between cationic and anionic entities (e.g., $[Co(NH_3)_6][Cr(CN)_6]$ and $[Cr(NH_3)_6][Co(CN)_6]$).
• Ionisation Isomerism: Occurs when the counter ion acts as a potential ligand and displaces a ligand (e.g., $[Co(NH_3)_5(SO_4)]Br$ and $[Co(NH_3)_5Br]SO_4$).
• Solvate Isomerism: Also called 'hydrate isomerism' if the solvent is water. Differs by whether solvent is directly bonded to the metal or free in the lattice (e.g., $[Cr(H_2O)_6]Cl_3$ (violet) and $[Cr(H_2O)_5Cl]Cl_2 \times H_2O$ (grey-green)).

Bonding in Coordination Compounds

Valence Bond Theory (VBT)

• Metal atom or ion under the influence of ligands uses $(n-1)d$, $ns$, $np$, or $ns$, $np$, $nd$ orbitals for hybridisation.
• Hybridisation for Coordination Numbers:     * $4$: $sp^3$ (Tetrahedral)     * $4$: $dsp^2$ (Square planar)     * $5$: $sp^3d$ (Trigonal bipyramidal)     * $6$: $sp^3d^2$ (Octahedral, Outer orbital)     * $6$: $d^2sp^3$ (Octahedral, Inner orbital)
• Magnetic Character: Paramagnetic (unpaired electrons), Diamagnetic (no unpaired electrons).
• Complex Types:     * Inner Orbital/Low Spin/Spin Paired: Uses inner $(n-1)d$ orbitals (e.g., $[Co(NH_3)_6]^{3+}$).     * Outer Orbital/High Spin/Spin Free: Uses outer $nd$ orbitals (e.g., $[CoF_6]^{3-}$).
• Limitations of VBT: Many assumptions; lacks quantitative interpretation of magnetic data, thermodynamic/kinetic stability, and cannot explain color.

Crystal Field Theory (CFT)

• An electrostatic model where metal-ligand bonds are purely ionic. Ligands are point charges (anions) or point dipoles (neutrals).
• Octahedral Splitting ($\Delta_o$):     * The five degenerate $d$ orbitals split into two sets: $t_{2g}$ (lower energy: $d_{xy}$, $d_{yz}$, $d_{xz}$) and $e_g$ (higher energy: $d_{x^2-y^2}$, $d_{z^2}$).     * $e_g$ energy increases by 3/5 \times \text{ } elta_o (or 0.6 \times \text{ } elta_o).     * $t_{2g}$ energy decreases by 2/5 \times \text{ } elta_o (or 0.4 \times \text{ } elta_o).
• Spectrochemical Series: Arranges ligands by field strength: $I^- < Br^- < SCN^- < Cl^- < S^{2-} < F^- < OH^- < C_2O_4^{2-} < H_2O < NCS^- < edta^{4-} < NH_3 < en < CN^- < CO$.
• Electron Configuration Outcomes:     * If $\Delta_o < P$ (Pairing energy), weak field ligands form high spin complexes ($t_{2g}^3 e_g^1$ for $d^4$).     * If $\Delta_o > P$, strong field ligands form low spin complexes ($t_{2g}^4 e_g^0$ for $d^4$).
• Tetrahedral Splitting ($\Delta_t$): Splitting is inverted and smaller; \Delta_t = (4/9) \times \text{ } elta_o. Low spin configurations involve rare pairing due to small splitting.
• Color: Attributed to $d-d$ transitions. Light of specific wavelength is absorbed to excite an electron from a lower to a higher energy $d$ orbital. The observed color is complementary to the absorbed one. Example: $[Ti(H_2O)_6]^{3+}$ absorbs blue-green light and appears violet.
• Limitations of CFT: Treats ligands as point charges; fails to account for covalent bonding; predicts anionic ligands should have the largest splitting, but they are often at the low end of the spectrochemical series.

Bonding in Metal Carbonyls

• Homoleptic Carbonyls: Formed by most transition metals. Structures are simple: $Ni(CO)_4$ (tetrahedral), $Fe(CO)_5$ (trigonal bipyramidal), $Cr(CO)_6$ (octahedral).
• Synergic Bonding:     * Sigma (\text{\sigma}) Bond: Donation of a lone pair from carbonyl carbon to a vacant metal orbital.     * Pi (\text{\pi}) Bond: "Back-bonding" from a filled metal $d$ orbital into a vacant antibonding \text{\pi}^\text{*} orbital of $CO$.     * This effect strengthens the bond between the metal and the carbonyl.

Importance and Applications

• Analytical Chemistry: Detection and estimation of metals using reagents like EDTA, DMG, \text{\alpha-nitroso-\beta-naphthol}. Hardness of water is estimated via titration with $Na_2EDTA$.
• Metallurgy: Extraction of silver and gold involves cyanide complex formation (e.g., $[Au(CN)_2]^-$). Impure nickel is purified via the Mond process ($Ni(CO)_4$).
• Biological Systems: Includes chlorophyll ($Mg$), haemoglobin ($Fe$), and Vitamin $B_{12}$ ($Co$). Enymes like carboxypeptidase A are also complexes.
• Industry: Catalyst examples include Wilkinson catalyst ($[(Ph_3P)_3RhCl]$) for alkene hydrogenation.
• Miscellaneous: Used in electroplating (smoothness), and black and white photography (using hypo to dissolve $AgBr$ as $[Ag(S_2O_3)_2]^{3-}$).
• Medicine: Chelate therapy removes toxic proportions of metals (e.g., $D-penicillamine$ for copper, $desferrioxime B$ for iron, $EDTA$ for lead poisoning). Cis-platin is used to inhibit tumor growth.

Questions & Discussion

• Isomerism Evidence: $[Co(NH_3)_5Cl]SO_4$ and $[Co(NH_3)_5(SO_4)]Cl$ are ionisation isomers. The first reacts with $Ba^2+$ ions to give a precipitate of $BaSO_4$, while the second does not. The second reacts with $Ag^+$ to give $AgBr$, the first does not.
• Tetrahedral Magnetism: $[NiCl_4]^{2-}$ is paramagnetic because $Cl^-$ is a weak ligand and does not cause pairing of electrons in the $3d^8$ configuration ($2$ unpaired electrons). $[Ni(CO)_4]$ is diamagnetic because $CO$ is a strong ligand causing pairing of electrons (zero oxidation state nickel).
• Magnetic Moment of manganese: $[MnBr_4]^{2-}$ has a magnetic moment of $5.9 \text{ BM}$. This corresponds to $5$ unpaired electrons, implying $sp^3$ hybridisation and tetrahedral geometry.
• Stability: Stability of complexes is affected by the chelate effect; chelated complexes (like those with $C_2O_4^{2-}$) are generally more stable than non-chelated ones.