Comprehensive Study Guide for Coordination Compounds

Fundamental Concepts of Coordination Chemistry

• Addition Compounds: These are formed when solutions containing two or more salts in stoichiometric ratios are allowed to evaporate, resulting in the formation of crystals.

• Coordination Compounds: These are compounds in which a central metal atom is linked to ions or neutral molecules by coordinate bonds.

  • Example: $[Cr(H_2O)_5Cl]^{2+}$.

• Coordination Entity / Coordination Sphere / Counter Sphere: A coordination entity consists of a central metal atom or ion bonded to a fixed number of ions or molecules.

  • In $[CoCl_3(NH_3)_3]$, the cobalt ion is the central entity surrounded by three ammonia molecules and three chloride ions.
  • Other examples: $[Ni(CO)_4]$, $[PtCl_2(NH_3)_2]$, $[Fe(CN)_6]^{4-}$, $[Co(NH_3)_6]^{3+}$.

• Central Metal Ion: This is an acceptor atom containing vacant orbitals. A fixed number of ligands are attached to it via coordinate bonds in a definite geometrical arrangement.

• Ionization Sphere / Counter Ions: These are the ions present outside the coordination sphere.

  • Example: In $K_4[Fe(CN)_6]$, $K^+$ represents the counter ions.

• Coordination Number (C.N.): The total number of ligands attached to the central metal atom through coordinate bonds.

  • Example: In $[Ag(CN)_2]^- \text{, the C.N. of } Ag^+ = 2$.

• Oxidation Number (O.N.) or State: This represents the charge the central atom would carry if all ligands were removed along with the shared electron pairs.

  • $[Cu(NH_3)_4]^{2+} \text{: O.N. of } Cu \text{ is } +2$.
  • $[Ni(CO)_4] \text{: O.N. of } Ni \text{ is } 0$.
  • $[Fe(CN)_6]^{3-} \text{: O.N. of } Fe \text{ is } +3$.

• Charge on the Metal Ion: This is the algebraic sum of the charges carried by the central metal ion and its coordinated ligands.

  • Example: In $[Ag(CN)_2]^-$, $Ag^+ \text{ has } +1 \text{ charge and } CN^- \text{ has } -1$. The net charge is $-1$.

Classification of Ligands

• Ligands: Ions or molecules bound to the central atom/ion in the coordination entity.

• Unidentate Ligand: A ligand bound to a metal ion through a single donor atom.

  • Examples: $Cl^-$, $H_2O$, $NH_3$.

• Bidentate Ligand: A ligand that binds through two donor atoms.

  • Examples: $H_2NCH_2CH_2NH_2$ (ethane-1,2-diamine or "en"), $C_2O_4^{2-}$ (oxalate or "ox").

• Tridentate Ligands: Ligands possessing three donor atoms per ligand.

  • Example: (dien) diethylenetriamine.

• Hexadentate Ligands: Ligands possessing six donor atoms per ligand.

  • Example: EDTA (Ethylene diamine tetra acetic acid).

• Bridging Ligands: Mono-dentate ligands that simultaneously coordinate through two metal atoms, acting as a bridge. Complexes containing these are called bridged complexes.

Formula Writing and Nomenclature Rules

• Rules for Formula Writing:

  1. The formula of the cation (simple or complex) is written first.
  2. The coordination entity is enclosed in square brackets.
  3. Inside the coordination sphere, the metal atom is written first, followed by ligands in alphabetical order of their names.
  4. For abbreviated ligands (e.g., en, ox), the first letter of the abbreviation is considered for alphabetical order.
  5. Polyatomic ligands are enclosed in parentheses (e.g., $(SCN)$, $(PPh_3)$).
  6. There is no space between ligands and the metal.
  7. For a charged coordination entity, the charge is indicated as a right superscript outside the square brackets with the number before the sign.
  8. The charge of the cation must be balanced by the charge of the anion.

• Rules for Naming Mononuclear Coordination Compounds:

  1. The complex cation is named first, followed by the anion.
  2. Ligands are named in alphabetical order before the metal atom or ion.
  3. Anionic ligands end in "-o" (e.g., Cyanido/Cyano, Chlorido/chloro).
  4. Neutral ligands often have special names: $H_2O$ is aqua, $NH_3$ is ammine, $CO$ is carbonyl, $NO$ is nitrosyl.
  5. Positive ligands end in "-ium" (e.g., $NO^+$ is nitrosonium).
  6. Prefixes such as mono, di, and tri indicate the number of individual ligands.
  7. If a ligand name includes a numerical prefix, terms like bis, tris, and tetrakis are used, and the ligand name is placed in parentheses.
  8. If the complex is a cation or neutral, the metal is named as the element. If the complex is an anion, the metal name ends with the suffix "-ate".

Nomenclature Examples and Chemical Formulas

• $[PtCl(NH_2CH_3)(NH_3)_2]Cl$: diammine chlorido methyl amine platinum (II) chloride
• $[Pt(NH_3)_4Cl_2]Cl_2$: tetraaminedichlorido platinum (IV) chloride (CBSE 2011C)
• $[Cr(NH_3)_6][Co(CN)_6]$: hexaammine chromium (III) hexacyanocobaltate (III)
• $[Pt(NH_3)_4Cl_2(NO)]Br_3$: tetraamminedichloridonitrosyl platinum (IV) bromide
• $[Co(CN)_2(NH_3)_4]Cl$: tetraammine dicyano cobalt (III) chloride
• $[Cr(NH_3)_5(NCS)][ZnCl_4]$: penta ammine isothiocyanato chromium (II) tetrachlorido Zincate (II) (CBSE 2011, Delhi 2010C)
• $[Co(NH_3)_5Cl]Cl_2$: pentaammine chlorido cobalt (III) chloride
• $[Cr(NH_3)_2Cl_2(en)]Cl$: diamminedichlorido (ethane-1,2-diamine) chromium (III) chloride
• $K_4[Fe(CN)_6]$: Potassium hexacyanoferrate (II) (CBSE 2015)
• $[NiCl_4]^{2-}$: Tetrachlorido nickelate (II) ion (Delhi 2015)
• $[Co(NH_3)_5(NO_2)]^{2-}$: Pentaammine nitrito-N cobalt (III) ion (CBSE 2015 / Delhi 2014C)
• $[Pt(NH_3)_2Cl_2]$: Diammine dichlorido platinum (II)
• $[Co(C_2O_4)_3]^{3-}$: Trioxalato Cobaltate (III) ion
• $[Cr(CO)_6]$: Hexacarbonyl chromium (0)
• $[PtCl_3(C_2H_4)]^-$: Trichlorido ethene platinum (II)
• $[CoBr_2(en)_2]^+$: Dibromido bis (ethane-1,2-diamine) Cobalt (III) ion
• $K_3[Fe(C_2O_4)_3]$: Potassium trioxalato ferrate (III) ion
• $[CoCl_2(en)_2]^+$: Dichlorido bis (ethane-1,2-diamine) Cobalt (III) ion
• $[Co(NH_3)_5(ONO)]^{2+}$: Pentaammine nitrito-O cobalt (III) ion
• Sodium dicyanido aurate (I): $Na[Au(CN)_2]$
• Tetraammine chloridonitrito-N-platinum (IV) sulphate: $[Pt(NH_3)_4Cl(NO_2)]SO_4$
• Mercury tetrathiocyanato cobaltate (II): $Hg[Co(SCN)_4]$
• Potassium trioxalato aluminate (III): $K_3[Al(C_2O_4)_3]$

Werner's Theory of Coordination Compounds

• Core Postulates:

  1. Metals in coordination compounds exhibit two types of valencies: primary and secondary.
  2. Primary Valencies: Normally ionisable and satisfied by negative ions. These correspond to the Oxidation State.
  3. Secondary Valencies: Non-ionisable and satisfied by neutral molecules or negative ions. These correspond to the Coordination Number and are fixed for a metal.
  4. Ions or groups bound by secondary linkages have characteristic spatial arrangements depending on the coordination number.

• Experimental Evidence (Precipitation with $AgNO_3$):

  • CoCl_3 · 6NH_3: Reacts with excess $AgNO_3$ to yield $3 \text{ mol AgCl}$. Formula: $[Co(NH_3)_6]Cl_3$. (Secondary valency satisfied by $6 NH_3$; Primary by $3 Cl^-$).
  • CoCl_3 · 5NH_3: Yields $2 \text{ mol AgCl}$. Formula: $[Co(NH_3)_5Cl]Cl_2$. (Secondary valency satisfied by $5 NH_3$ and $1 Cl^-$).
  • CoCl_3 · 4NH_3: Yields $1 \text{ mol AgCl}$. Formula: $[Co(NH_3)_4Cl_2]Cl$. (Secondary valency satisfied by $4 NH_3$ and $2 Cl^-$).

• Conductivity and Ion Counts:

  • Conductivity depends on the number of ions per mole of electrolyte.
  • Experimental order: CoCl_3 · 6NH_3 (4 \text{ ions}) > CoCl_3 · 5NH_3 (3 \text{ ions}) > CoCl_3 · 4NH_3 (2 \text{ ions}).

Valence Bond Theory (VBT)

• Key Features:

  • Coordination number depends on the number of vacant orbitals available for bonding in the metal ion.
  • The metal ion uses $(n-1)d$, $ns$, $np$, or $ns$, $np$, $nd$ orbitals for hybridisation to yield equivalent orbitals for specific geometries.

• Hybridisation and Geometry:

  • C.N. 4:
    • $sp^3$: Tetrahedral geometry.
    • $dsp^2$: Square planar geometry.
  • C.N. 6:
    • $sp^3d^2$: Octahedral (Outer orbital / high spin complex).
    • $d^2sp^3$: Octahedral (Inner orbital / low spin complex).

• Magnetic Properties:

  • Diamagnetic: All electrons are paired.
  • Paramagnetic: Unpaired electrons are present.

• Orbital Configurations for Specific Complexes:

  • $[Cr(H_2O)_6]^{3+}: d^3 \text{ config, } d^2sp^3 \text{ hybridisation, Octahedral, 3 unpaired electrons, Paramagnetic.}$
  • $[FeF_6]^{3-}: d^5 \text{ config, } sp^3d^2 \text{ hybridisation, Octahedral, 5 unpaired electrons, Paramagnetic.}$
  • $[Fe(CN)_6]^{4-}: d^6 \text{ config, } d^2sp^3 \text{ hybridisation, Octahedral, 0 unpaired electrons, Diamagnetic.}$

• Limitations of VBT:

  1. Involves numerous assumptions.
  2. Does not explain the colour of coordination compounds.
  3. Fails to explain kinetic and thermodynamic stabilities.
  4. Does not reliably distinguish between weak and strong ligands (e.g., $[Mn(CN)_6]^{3-}$ is paramagnetic but VBT predicts diamagnetic).
  5. Does not explain experimental hybridisation findings, such as $[Cu(NH_3)_4]^{2+}$ being $dsp^2$ instead of $sp^3$.

Isomerism in Coordination Compounds

• Structural Isomerism: Compounds with different bonds.

  • Linkage Isomerism: Arises from ambidentate ligands (e.g., $-NO_2$ vs $-ONO$).
  • Coordination Isomerism: Arises from the interchange of ligands between cationic and anionic entities of different metals.
  • Ionisation Isomerism: Arises when a counter ion displaces a ligand from the coordination sphere.
    • Example: One compound yields no precipitate with $BaCl_2$ but a white precipitate with the second; one yields yellow precipitate with $AgNO_3$ while the other yields white.
  • Solvate Isomerism: Involves solvent molecules (like water in "hydrate isomerism") being either directly bonded to the metal or present as free molecules in the lattice.
  • Ligand Isomerism: Different isomers of the same ligand attach to the metal (e.g., pn = 1,2-diaminopropane vs tn = 1,3-diaminopropane).

• Stereoisomerism: Same chemical formula and bonds but different spatial arrangements.

  • Geometrical Isomerism: Found in heteroleptic complexes.
    • Cis: Same kind of ligands occupy adjacent positions.
    • Trans: Same kind of ligands are opposite to each other.
    • Octahedral (Fac/Mer): Facial (fac) if three donor atoms are at corners of an octahedral face; Meridional (mer) if they are around the meridian.
  • Tetrahedral Complexes: Do not show geometrical isomerism because relative ligand positions are always adjacent.

Crystal Field Theory (CFT)

• Principles:

  1. Bonds between ligands and central metal are assumed to be purely ionic, representing electrostatic attraction.
  2. Metals/ions and ligands are viewed as point charges (or electric dipoles for neutral ligands).
  3. CFT focuses on the splitting of d-orbitals.
  4. Splitting Process: In an octahedral field, ligands approach along the axes ($x, y, z$). Orbitals along the axes ($d_{x^2-y^2}$ and $d_{z^2}$, the eg set) experience stronger repulsion and rise in energy. Orbitals between the axes ($d_{xy}$, $d_{yz}$, $d_{zx}$, the t2g set) have lower energy.

• Spectrochemical Series: Ligands on the right (e.g., Carbonyl) cause larger splitting ($Δ_o$) and are strong field ligands (SFL). Ligands on the left are weak field ligands (WFL).

Calculation of Crystal Field Stabilization Energy (CFSE)

• Formula: $CFSE = x \times (+0.6 Δ_o) + y \times (-0.4 Δ_o) + (Z \times P)$

  • $x$: Electrons in eg set.
  • $y$: Electrons in t2g set.
  • $Z$: Number of pairs formed (difference between pairs before and after splitting).
  • $P$: Pairing energy.

• Case Studies in Octahedral Field ($C.N. = 6$):

  • $d^1$ Configuration: Electron enters t2g. $CFSE = -0.4 Δ_o$.
  • $d^4$ Configuration (SFL): $Δ_o > P$, all 4 electrons in t2g. $CFSE = -1.6 Δ_o + P$.
  • $d^4$ Configuration (WFL): $Δ_o < P$, 4th electron enters eg. $CFSE = -0.6 Δ_o$.
  • $d^5$ Configuration (SFL): $Δ_o > P$, all 5 electrons in t2g. $CFSE = -2 Δ_o + P$.
  • $d^5$ Configuration (WFL): $Δ_o < P$, 4th and 5th electrons enter eg. $CFSE = 0$.
  • $d^6$ Configuration (SFL): $Δ_o > P$, all 6 electrons in t2g. $CFSE = -2.4 Δ_o + 2P$.
  • $d^6$ Configuration (WFL): $Δ_o < P$. $CFSE = -0.4 Δ_o$.

Color and Applications

• Color of Complexes: A substance appears colored when it absorbs specific visible light wavelengths and transmits others.

  • Transition metal color arises from d-d transitions: excitation of electrons from t2g to eg levels.
  • Complementary colors: Absorbing orange light results in a blue appearance (e.g., hydrated copper(II) ion).

• Practical Applications:

  1. Hardness estimation: EDTA is used to estimate $Ca^{2+}$ and $Mg^{2+}$ in water.
  2. Medicine: EDTA treats lead poisoning; cis-platin $[Pt(NH_3)_2Cl_2]$ is an anti-tumor agent for cancer treatment.
  3. Biology: Chlorophyll contains Magnesium complexes, Hemoglobin contains Iron, and Vitamin B12 contains Cobalt.
  4. Industry: Electroplating uses gold and silver complexes like $K[Ag(CN)_2]$.
  5. Metallurgy: Gold and silver are extracted using cyanide complexes ($[Ag(CN)_2]^-$ and $[Au(CN)_4]^-$).
  6. Analysis: Ni(II) is estimated as red glyoxime using dimethyl glyoxime (DMG).