Exhaustive Guide to Coordination Compounds

Introduction to Coordination Compounds

• In the transition metals, a large number of complex compounds are formed where metal atoms are bound to anions or neutral molecules through the sharing of electrons.

• These are specifically known as coordination compounds in modern terminology.

• Coordination chemistry is a vital area of modern inorganic chemistry, providing insights into biological systems, industrial processes, and metallurgical techniques.

• Biological Significance:

  • Chlorophyll: A coordination compound of magnesium ($Mg$).

  • Haemoglobin: A coordination compound of iron ($Fe$).

  • Vitamin B12: A coordination compound of cobalt ($Co$).

• Applications:

  • Used as industrial catalysts and analytical reagents.

  • Applied in electroplating, textile dyeing, and medicinal chemistry.

• Learning Objectives: Understanding Werner’s postulates, nomenclature (IUPAC), coordination terms (ligand, coordination number, etc.), isomerism types, and bonding theories (Valence Bond Theory and Crystal Field Theory).

Werner’s Theory of Coordination Compounds

• Alfred Werner (1866-1919), a Swiss chemist, was the first to formulate structures for coordination compounds, earning the Nobel Prize in 1913.

• Werner proposed two types of valences for metal ions:

  • Primary Valence: Ionisable, satisfied by negative ions. This corresponds to the oxidation state of the metal. For example, in $CrCl_3$, $CoCl_2$, and $PdCl_2$, the primary valences are $3$, $2$, and $2$ respectively.

  • Secondary Valence: Non-ionisable, satisfied by neutral molecules or negative ions. This corresponds to the coordination number and is fixed for a metal.

• Experimental Basis: Werner reacted a series of cobalt(III) chloride-ammonia complexes with excess silver nitrate ($AgNO_3$) in the cold. Observations were as follows:

  • $1\text{ mol } CoCl_3 \cdot 6NH_3$ (Yellow) yielded $3\text{ mol AgCl}$. Formulation: $[Co(NH_3)_6]Cl_3$.

  • $1\text{ mol } CoCl_3 \cdot 5NH_3$ (Purple) yielded $2\text{ mol AgCl}$. Formulation: $[CoCl(NH_3)_5]Cl_2$.

  • $1\text{ mol } CoCl_3 \cdot 4NH_3$ (Green) yielded $1\text{ mol AgCl}$. Formulation: $[CoCl_2(NH_3)_4]Cl$.

  • $1\text{ mol } CoCl_3 \cdot 4NH_3$ (Violet) yielded $1\text{ mol AgCl}$. Formulation: $[CoCl_2(NH_3)_4]Cl$.

• Isomers: The Green and Violet forms have the same empirical formula ($CoCl_3 \cdot 4NH_3$) but different properties, making them isomers.

• Modern Formulations: Species within square brackets ($[…]$) are coordination entities or complexes that do not dissociate in solution. Ions outside the brackets are counter ions.

• Postulates of Werner’s Theory:

  1. Metals show two types of linkages (valences): primary and secondary.

  2. Primary valences are satisfied by negative ions and are ionisable.

  3. Secondary valences are satisfied by neutral or negative groups and are non-ionisable; the number of these is fixed for a metal.

  4. Ions/groups bound by secondary linkages have characteristic spatial arrangements/geometries (coordination polyhedra).

Definitions of Key Terms 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 groups are bound in a specific geometry. These are considered Lewis acids. Examples: $Ni^{2+}$ in $[NiCl_2(H_2O)_4]$, $Co^{3+}$ in $[CoCl(NH_3)_5]^{2+}$, and $Fe^{3+}$ in $[Fe(CN)_6]^{3-}$.

• Ligands: Ions or molecules bound to the central atom. They can be:

  • Unidentate: Bound via a single donor atom (e.g., $Cl^-$, $H_2O$, $NH_3$).

  • Didentate: Bound via two donor atoms (e.g., ethane-1,2-diamine/$en$ or oxalate/$C_2O_4^{2-}$).

  • Polydentate: Bound via several donor atoms. Example: $EDTA^{4-}$ (ethylenediaminetetraacetate ion) is hexadentate, binding via two nitrogen and four oxygen atoms.

  • Chelate Ligand: A di- or polydentate ligand using two or more donor atoms simultaneously to bind a single metal ion, forming a ring. Chelate complexes are generally more stable (chelate effect).

  • Ambidentate Ligand: Ligands with two different donor atoms through either of which they can coordinate. Examples: $NO_2^-$ (N or O) and $SCN^-$ (S or N).

• Coordination Number (CN): The number of ligand donor atoms directly bonded to the metal. Only sigma ($\sigma$) bonds are counted; pi ($\pi$) bonds are excluded.

• Coordination Sphere: The central atom and ligands enclosed in square brackets. Ions outside are counter ions.

• Coordination Polyhedron: The spatial arrangement of ligands around the metal (e.g., octahedral, tetrahedral, square planar).

• Oxidation Number: The charge the metal would carry if all ligands and shared electron pairs were removed. Indicated by Roman numerals in parentheses.

• Homoleptic Complexes: Metal is bound to only one kind of donor group (e.g., $[Co(NH_3)_6]^{3+}$).

• Heteroleptic Complexes: Metal is bound to more than one kind of donor group (e.g., $[Co(NH_3)_4Cl_2]^+$).

Difference Between Double Salts and Complexes

• Double Salts: Formed by stable compounds in stoichiometric ratios but dissociate completely into simple ions in water. Examples: Carnallite ($KCl \cdot MgCl_2 \cdot 6H_2O$), Mohr’s salt ($FeSO_4 \cdot (NH_4)_2SO_4 \cdot 6H_2O$), Potash alum ($KAl(SO_4)_2 \cdot 12H_2O$).

• Complexes: Do not dissociate completely into simple ions. The complex ion (e.g., $[Fe(CN)_6]^{4-}$) remains intact in solution.

Nomenclature of Coordination Compounds

• Formulas for Mononuclear Coordination Entities:

  1. Central atom is listed first.

  2. Ligands follow in alphabetical order regardless of charge.

  3. Polydentate ligands are also alphabetical; abbreviations are used for ordering.

  4. The entity is enclosed in square brackets. Polyatomic ligands are in parentheses.

  5. No space between ligands and metal within the sphere.

  6. Charge is indicated outside the bracket as a superscript (e.g., $[Co(CN)_6]^{3-}$).

• Naming Rules:

  1. Cation is named before the anion.

  2. Ligands are named alphabetically first.

  3. Anionic ligands end in ‘-o’ or ‘-ido’ (e.g., chlorido). Neutral ligands use specific names: $H_2O$ (aqua), $NH_3$ (ammine), $CO$ (carbonyl), $NO$ (nitrosyl).

  4. Prefixes (di, tri, tetra) indicate quantity. If the ligand name has a prefix, use bis, tris, tetrakis.

  5. Oxidation state is shown in Roman numerals.

  6. If the complex is an anion, the metal name ends in ‘-ate’. Latin names are often used: ferrate ($Fe$), argentate ($Ag$), aurate ($Au$), cuprate ($Cu$), stannate ($Sn$), plumbate ($Pb$).

• Example: $[Cr(NH_3)_3(H_2O)_3]Cl_3$ is triamminetriaquachromium(III) chloride.

• Example: $Ag(NH_3)_2][Ag(CN)_2]$ is diamminesilver(I) dicyanidoargentate(I).

Isomerism in Coordination Compounds

• Stereoisomerism:

  • Geometrical Isomerism: Arises in heteroleptic complexes.

  • Square Planar ($[MX_2L_2]$ or $MABXL$): Cis (adjacent) and Trans (opposite) forms. Tetrahedral complexes do not show geometrical isomerism.

  • Octahedral ($[MX_2L_4]$ or $[MX_2(L-L)_2]$): Cis and Trans forms.

  • Facial (fac) and Meridional (mer): Occurs in $[Ma_3b_3]$ complexes like $[Co(NH_3)_3(NO_2)_3]$. fac-isomer has three donor atoms on adjacent octahedral corners; mer-isomer are around the meridian.

  • Optical Isomerism: Molecules that are non-superimposable mirror images (enantiomers/chiral). Dextro ($d$) rotates light right; Laevo ($l$) rotates light left. Common in didentate octahedral complexes.

• 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., $[Co(NH_3)_6][Cr(CN)_6]$ vs $[Cr(NH_3)_6][Co(CN)_6]$).

  • Ionisation Isomerism: Counter ion displaces a ligand (e.g., $[Co(NH_3)_5(SO_4)]Br$ and $[Co(NH_3)_5Br]SO_4$).

  • Solvate/Hydrate Isomerism: Differs by whether solvent water is bonded to metal or free in crystal lattice (e.g., $[Cr(H_2O)_6]Cl_3$ [violet] vs $[Cr(H_2O)_5Cl]Cl_2 \cdot H_2O$ [grey-green]).

Valence Bond Theory (VBT)

• Metal uses $(n-1)d$, $ns$, $np$, or $ns$, $np$, $nd$ orbitals for hybridisation.

• Hybridisation and Geometry:

  • $CN = 4$: $sp^3$ (Tetrahedral) or $dsp^2$ (Square planar).

  • $CN = 5$: $sp^3d$ (Trigonal bipyramidal).

  • $CN = 6$: $sp^3d^2$ (Outer orbital/High spin octahedral) or $d^2sp^3$ (Inner orbital/Low spin octahedral).

• Magnetic Properties: Complexes with unpaired electrons are paramagnetic; those without are diamagnetic.

  • Inner Orbital Complex: Uses $(n-1)d$ orbitals (low spin, e.g., $[Co(NH_3)_6]^{3+}$).

  • Outer Orbital Complex: Uses $nd$ orbitals (high spin, e.g., $[CoF_6]^{3-}$).

• Limitations: Does not give quantitative magnetic data, explain colour, explain vary thermodynamic stabilities, or distinguish between weak and strong ligands consistently.

Crystal Field Theory (CFT)

• Electrostatic model treating ligands as point charges (anions) or dipoles (neutrals).

• Octahedral Splitting: Repulsion between metal d-electrons and ligand electrons splits the degenerate d-orbitals.

  • Two $e_g$ orbitals ($d_{x^2-y^2}$, $d_{z^2}$) point toward axes and increase in energy by $\frac{3}{5} \Delta_o$.

  • Three $t_{2g}$ orbitals ($d_{xy}$, $d_{yz}$, $d_{xz}$) point between axes and decrease in energy by $\frac{2}{5} \Delta_o$.

• Spectrochemical Series: Ligands arranged 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

• High vs. Low Spin:

  • If \Delta_o < P (Pairing energy): Weak field ligand, high spin electron configuration.

  • If \Delta_o > P: Strong field ligand, low spin (electrons pair up in $t_{2g}$).

• Tetrahedral Splitting: Inverted and smaller than octahedral ($\Delta_t = \frac{4}{9} \Delta_o$). Low spin tetrahedral complexes are rare.

• Colour: Explained by d-d transitions. Energy absorption from visible light excites an electron from a lower to higher d-state. complementary colour is seen.

• Gemstones:

  • Ruby: Al_2_O_3 with $0.5\text{-}1\% Cr^{3+}$ ions in octahedral sites.

  • Emerald: $Cr^{3+}$ ions in octahedral sites of the mineral beryl ($Be_3Al_2Si_6O_{18}$).

Bonding in Metal Carbonyls

• Metal carbonyls feature a metal-carbon bond with both $\sigma$ and $\pi$ character.

• Synergic Effect:

  • $M-C \sigma$ bond: Carbonyl carbon donates a lone pair into a vacant metal orbital.

  • $M-C \pi$ bond: Back-donation of electrons from filled metal d-orbitals into vacant antibonding $\pi^*$ orbitals of $CO$.

• Structures:

  • $Ni(CO)_4$: Tetrahedral.

  • $Fe(CO)_5$: Trigonal bipyramidal.

  • $Cr(CO)_6$: Octahedral.

  • $Mn_2(CO)_{10}$: Two square pyramidal units with an $Mn-Mn$ bond.

  • $Co_2(CO)_8$: $Co-Co$ bond with two bridging $CO$ groups.

Applications and Importance

• Analytical Chemistry: Reagents like $EDTA$, $DMG$, and $\alpha\text{-nitroso-}\beta\text{-naphthol}$ are used for metal detection and estimation.

• Water Hardness: Estimated via titration with $Na_2EDTA$ which complexes $Ca^{2+}$ and $Mg^{2+}$.

• Metallurgy: Extraction of gold and silver via cyanide complexes. Purification of nickel via the Mond process ($Ni(CO)_4$ decomposition).

• Medicine:

  • Chelate Therapy: $EDTA$ for lead poisoning; $D\text{-penicillamine}$ for excess copper ($Cu$); $desferrioxime B$ for iron ($Fe$).

  • Antitumour: $cis\text{-platin}$ inhibiting tumor growth.

• Industrial Catalysts: Wilkinson catalyst $[(Ph_3P)_3RhCl]$ for alkene hydrogenation.

• Photography: Hypo solution fixes film by dissolving undecomposed $AgBr$ as $[Ag(S_2O_3)_2]^{3-}$.

Questions & Discussion

• Example 5.7: $[MnBr_4]^{2-}$ has a magnetic moment of $5.9\,BM$. Predict geometry. Solution: $5.9\,BM$ indicates $5$ unpaired electrons, implying $sp^3$ hybridisation and tetrahedral geometry.

• Isomerism in [Co(en)3]3+: This complex exhibits optical isomerism ($d$ and $l$ forms).

• Difference in Magnetism: $[Ni(CN)_4]^{2-}$ is diamagnetic (square planar, $dsp^2$), while $[NiCl_4]^{2-}$ is paramagnetic (tetrahedral, $sp^3$) because $CN^-$ is a strong field ligand and $Cl^-$ is weak.

• Stability of Complexes: $[Fe(C_2O_4)_3]^{3-}$ is more stable than $[Fe(H_2O)_6]^{3+}$ due to the chelate effect provided by the oxalate ligand.