Coordination Compounds revision notes

Rules for Writing Formulas

1. Central atom is listed first.
2. Ligands are listed in alphabetical order, irrespective of their charge (2004 IUPAC Draft Update) JEE Tip.
3. Polydentate ligands and ligand abbreviations are enclosed in parentheses.
4. No space between ligands and the metal within the square brackets.
5. Charge of the complex is indicated outside the square brackets as a right superscript.

Rules for Naming Compounds

1. Cation first: The cation is named before the anion.
2. Counter-ion Multipliers JEE Tip: You never indicate the number of cations or anions in the name of an ionic coordination compound. For example, $[\text{Co}(\text{H}_2\text{NCH}_2\text{CH}_2\text{NH}_2)_3]_2(\text{SO}_4)_3$ is named tris(ethane-1,2–diamine)cobalt(III) sulphate, not trisulphate.
3. Alphabetical ligands: Ligands are named in alphabetical order before the name of the central atom (reverse of formula writing).
4. Ligand Suffixes (2004 IUPAC Draft Update): Anionic ligands end in -ido (e.g., chloro becomes chlorido, cyano becomes cyanido). Neutral/cationic ligands keep their names, with strict exceptions: aqua ($\text{H}_2\text{O}$), ammine ($\text{NH}_3$) JEE Tip note the double 'm', carbonyl ($\text{CO}$), and nitrosyl ($\text{NO}$).
5. Numerical Prefixes: mono, di, tri are used for simple ligands. If the ligand name contains a prefix itself (like ethane-1,2-diamine), use bis, tris, tetrakis and enclose the ligand name in parentheses.
6. Oxidation State: Indicated by a Roman numeral in parentheses.
7. Complex Anion Suffix: If the complex is an anion, the metal name ends in -ate (e.g., cobaltate, ferrate, argentate, zincate). If it's a cation or neutral, the regular element name is used.

Isomers are compounds with the same chemical formula but different arrangements of atoms.

1. Stereoisomerism (Same bonds, different spatial arrangement)

Geometrical Isomerism

• Arises in heteroleptic complexes due to different geometric arrangements.
• Square Planar (CN=4): $[\text{MX}_2\text{L}_2]$ type forms cis (adjacent) and trans (opposite). $[\text{MABXL}]$ type forms 3 isomers (2 cis, 1 trans) JEE Tip. Tetrahedral geometry does NOT show geometrical isomerism because all positions are adjacent to each other.
• Octahedral (CN=6): $[\text{MX}_2\text{L}_4]$ or $[\text{MX}_2(\text{L-L})_2]$ forms cis and trans (e.g., $[\text{CoCl}_2(\text{en})_2]^+$).
• Facial and Meridional: Occurs in octahedral $[\text{Ma}_3\text{b}_3]$ complexes (e.g., $[\text{Co}(\text{NH}_3)_3(\text{NO}_2)_3]$). Facial (fac): three identical donor atoms occupy adjacent positions at the corners of an octahedral face. Meridional (mer): positions are around the meridian of the octahedron JEE Tip.

Optical Isomerism

• Mirror images that cannot be superimposed (chiral/enantiomers). Defined as dextro ($d$) or laevo ($l$).
• Common in octahedral complexes involving didentate ligands (e.g., $[\text{Co}(\text{en})_3]^{3+}$).
• JEE Tip In $[\text{M}(\text{L-L})_2\text{X}_2]$ type complexes, ONLY the cis-isomer is optically active; the trans-isomer has a plane of symmetry and is optically inactive.

2. Structural Isomerism (Different bonds/connections)

• Linkage Isomerism: Arises in complexes containing ambidentate ligands (e.g., $[\text{Co}(\text{NH}_3)_5(\text{NO}_2)]\text{Cl}_2$ – red form maps to M-ONO, yellow form maps to M-NO$_2$).
• Coordination Isomerism: Interchange of ligands between cationic and anionic entities of different metal ions present in a single complex salt (e.g., $[\text{Co}(\text{NH}_3)_6][\text{Cr}(\text{CN})_6]$ and $[\text{Cr}(\text{NH}_3)_6][\text{Co}(\text{CN})_6]$).
• Ionisation Isomerism: The counter ion itself is a potential ligand and displaces a ligand inside the sphere, which then becomes the counter ion (e.g., $[\text{Co}(\text{NH}_3)_5(\text{SO}_4)]\text{Br}$ and $[\text{Co}(\text{NH}_3)_5\text{Br}]\text{SO}_4$).
• Solvate (Hydrate) Isomerism: Similar to ionization isomerism, but the differing molecule is a solvent (usually water). Free solvent vs. directly bonded solvent (e.g., $[\text{Cr}(\text{H}_2\text{O})_6]\text{Cl}_3$ violet vs $[\text{Cr}(\text{H}_2\text{O})_5\text{Cl}]\text{Cl}_2\cdot\text{H}_2\text{O}$ grey-green).

Valence Bond Theory (VBT)

• Metal atom uses empty $(n-1)d, ns, np$ or $ns, np, nd$ orbitals to hybridize and yield equivalent orbitals for overlapping with ligand orbitals.

Coordination Number 4

• Tetrahedral ($sp^3$): Outer $s$ and $p$ orbitals hybridize. Example: $[\text{NiCl}_4]^{2-}$ ($d^8$, weak field $\text{Cl}^-$, paramagnetic, 2 unpaired electrons) JEE Tip.
• Square Planar ($dsp^2$): Example: $[\text{Ni}(\text{CN})_4]^{2-}$ ($d^8$, strong field $\text{CN}^-$, causes pairing, diamagnetic).

Coordination Number 6

• Inner Orbital Complex ($d^2sp^3$): Uses inner $(n-1)d$ orbitals. Generally low spin/spin-paired. Example: $[\text{Co}(\text{NH}_3)_6]^{3+}$ (diamagnetic).
• Outer Orbital Complex ($sp^3d^2$): Uses outer $nd$ orbitals. Generally high spin/spin-free. Example: $[\text{CoF}_6]^{3-}$ (paramagnetic).

Crystal Field Theory (CFT)

• An electrostatic model considering metal-ligand bonds as purely ionic (point charges for anions, point dipoles for neutral molecules). Removes degeneracy of the five $d$ orbitals.

Octahedral Splitting ($\Delta_o$)

• Ligands approach along axes ($x, y, z$). Repulsion raises the energy of $d_{x^2-y^2}$ and $d_{z^2}$ (the $e_g$ set) by $\frac{3}{5}\Delta_o$.
• Energy of $d_{xy}, d_{yz}, d_{xz}$ (the $t_{2g}$ set) decreases by $\frac{2}{5}\Delta_o$.
• JEE Tip If $\Delta_o < P$ (pairing energy), $4^{th}$ electron enters $e_g$ $\rightarrow t_{2g}^3 e_g^1$ (Weak field ligand, High spin). If $\Delta_o > P$, $4^{th}$ electron pairs in $t_{2g}$ $\rightarrow t_{2g}^4 e_g^0$ (Strong field ligand, Low spin).

Tetrahedral Splitting ($\Delta_t$)

• Splitting is inverted: $e$ set is lower, $t_2$ set is higher.
• $\Delta_t = \frac{4}{9}\Delta_o$. JEE Tip Because $\Delta_t$ is small, pairing energy is rarely overcome. Tetrahedral complexes are almost always high spin; low spin configurations are rarely observed.

Color & Wavelength Relationships

• Originates from $d-d$ electron transitions from a lower to higher split $d$-orbital (e.g., $t_{2g}$ to $e_g$).
• The complex absorbs a specific wavelength, and the observed color is the complementary color.
• $[Co(CN)_6]^{3-}$ absorbs in the ultraviolet region (310 nm) and appears pale yellow.
• $[Cu(H_2O)_4]^{2+}$ absorbs red light (600 nm) and appears blue.
• $[Ti(H_2O)_6]^{3+}$ absorbs blue-green light (498 nm) and appears violet.
• Progressive substitution of ligands changes $\Delta_o$ and therefore the color. E.g., $[\text{Ni}(\text{H}_2\text{O})_6]^{2+}$ (green) + en $\rightarrow$ pale blue $\rightarrow$ blue/purple $\rightarrow$ $[\text{Ni}(\text{en})_3]^{2+}$ (violet).
• Gemstones JEE Tip: Ruby's color is due to $\text{Cr}^{3+}$ ($d^3$) replacing $\text{Al}^{3+}$ in $\text{Al}_2\text{O}_3$ lattice. Emerald's color is due to $\text{Cr}^{3+}$ occupying octahedral sites in beryl ($\text{Be}_3\text{Al}_2\text{Si}_6\text{O}_{18}$).

Advanced Bonding Theories

• While VBT and CFT are studied in detail, Ligand Field Theory (LFT) and Molecular Orbital Theory (MOT) are used to explain the covalent character of metal-ligand bonding, which CFT completely fails to address.

Bonding in Metal Carbonyls (Synergic Bonding)

• Homoleptic carbonyls contain only CO ligands (e.g., $[\text{Ni}(\text{CO})_4]$ - tetrahedral, $[\text{Fe}(\text{CO})_5]$ - trigonal bipyramidal, $[\text{Cr}(\text{CO})_6]$ - octahedral).

Dinuclear Metal Carbonyl Structures

• Decacarbonyldimanganese(0) $[Mn_2(CO)_{10}]$: Made up of two square pyramidal $Mn(CO)_5$ units joined together by a single Mn–Mn bond.
• Octacarbonyldicobalt(0) $[Co_2(CO)_8]$: Contains a Co–Co bond that is bridged by two CO groups.

• Spectrochemical Series JEE Tip: Ligands arranged in increasing order of field strength (crystal field splitting $\Delta_o$): $\text{I}^- < \text{Br}^- < \text{SCN}^- < \text{Cl}^- < \text{S}^{2-} < \text{F}^- < \text{OH}^- < \text{C}_2\text{O}_4^{2-} < \text{H}_2\text{O} < \text{NCS}^- < \text{edta}^{4-} < \text{NH}_3 < \text{en} < \text{CN}^- < \text{CO}$
• Stability of Chelates: Chelating ligands (en, ox, EDTA) form much more stable complexes compared to unidentate analogues (Chelate Effect).

• Crystal Field Splitting Relation: $\Delta_t = \frac{4}{9} \Delta_o$
• Magnetic Moment (Spin Only): $\mu = \sqrt{n(n+2)} \text{ BM}$ (where $n$ = number of unpaired electrons).

• Werner's Primary vs Secondary Valence Rule: Primary valences are ionizable and satisfied by anions. Secondary valences are non-ionizable, satisfied by neutral or anionic species, and determine the coordination polyhedron.
• Hund's Rule in CFT: In $d^1, d^2, d^3$, electrons singly occupy $t_{2g}$ levels. Pairing decisions start only at the $d^4$ configuration depending on the magnitude of $\Delta_o$ vs $P$.

• Analytical Chemistry: Complex formation used in qualitative/quantitative analysis. Reagents include EDTA, DMG (dimethylglyoxime), $\alpha$-nitroso-$\beta$-naphthol, cupron.
• Water Hardness Estimation: Simple titration with $\text{Na}_2\text{EDTA}$. Forms stable complexes with $\text{Ca}^{2+}$ and $\text{Mg}^{2+}$.

Metallurgy

• Macarthur-Forrest Cyanide Process: Extraction of Au and Ag. Gold + cyanide + $\text{O}_2$ + water $\rightarrow$ $[\text{Au}(\text{CN})_2]^-$. Pure metal is retrieved by displacement with Zinc.
• Mond Process: Purification of Nickel relies on the formation and subsequent decomposition of $[\text{Ni}(\text{CO})_4]$.

Biological Systems

• Chlorophyll: Coordination compound of Magnesium.
• Haemoglobin: Coordination compound of Iron (oxygen carrier).
• Vitamin B12 (Cyanocobalamine): Coordination compound of Cobalt (anti-pernicious anaemia factor).
• Enzymes: Carboxypeptidase A and Carbonic anhydrase require coordinated metal ions.
• Industrial Catalysts: Wilkinson’s Catalyst $[\text{RhCl}(\text{PPh}_3)_3]$ is used for the hydrogenation of alkenes.
• Electroplating: Electroplating Ag and Au is smoother using complex solutions $[\text{Ag}(\text{CN})_2]^-$ and $[\text{Au}(\text{CN})_2]^-$ rather than simple metal ions.
• Photography: Hypo solution (sodium thiosulfate) "fixes" developed film by dissolving unreacted $\text{AgBr}$ to form complex ion $[\text{Ag}(\text{S}_2\text{O}_3)_2]^{3-}$.

Medicinal Chemistry (Chelate Therapy)

• Excess Cu and Fe removal: D-penicillamine and desferrioxime B.
• Lead poisoning: Treated with EDTA.
• Cancer therapy: cis-platin $[Pt(NH_3)_2Cl_2]$ inhibits tumor growth.

• Failure of Crystal Field Theory regarding Anionic Ligands: CFT assumes ligands act as point charges, meaning highly negatively charged anionic ligands should create the strongest repulsions and the greatest splitting. Anomaly: Anionic ligands (like $I^-, Br^-, Cl^-$) are actually found at the lowest end of the spectrochemical series.
• Tetrahedral Symmetry Subscript Exception: In Crystal Field Theory, the energy levels for octahedral and square planar complexes use the subscript 'g' (e.g., $t_{2g}$, $e_g$) because they possess a centre of symmetry. Exception: Tetrahedral complexes lack a center of symmetry, so the 'g' subscript is strictly omitted (energy levels are just $t_2$ and $e$).
• Absence of Low-Spin Tetrahedral Complexes: Because the crystal field splitting in tetrahedral complexes ($\Delta_t$) is significantly smaller than in octahedral complexes ($\Delta_t = \frac{4}{9} \Delta_o$), orbital splitting energies are almost never large enough to force electron pairing. Anomaly: Therefore, low spin configurations in tetrahedral geometries are rarely observed.
• VBT's Failure to Predict 4-Coordinate Geometry: Valence Bond Theory cannot make exact predictions regarding whether a 4-coordinate complex will adopt a tetrahedral ($sp^3$) or square planar ($dsp^2$) structure without prior knowledge of its magnetic moment.
• Failure of Valence Bond Theory to Explain Spin States: VBT cannot explain why $[\text{MnCl}_6]^{3-}$ ($d^4$) is high-spin while $[\text{Mn}(\text{CN})_6]^{3-}$ ($d^4$) is low spin. VBT simply attributes this to "inner vs outer orbital" without explaining the fundamental cause (ligand strength).
• Ligand Naming Exceptions: While most neutral ligands retain their molecule's name, there are four strict exceptions: $H_2O$ becomes aqua, $NH_3$ becomes ammine (note the double 'm'), $CO$ becomes carbonyl, and $NO$ becomes nitrosyl.
• Colorless Complex Exceptions: If crystal field splitting does not occur, $d-d$ transitions cannot happen, and the substance is colorless. Special case: Removal of water from violet $[Ti(H_2O)_6]Cl_3$ by heating removes the ligands, rendering it completely colourless. Anhydrous $CuSO_4$ is white for the same reason.

• IUPAC Nomenclature: Identifying the correct name, specifically the suffix rules (-ate), multiplier rules, and alphabetical ordering of complex, heavily substituted ligands (including the new -ido endings).
• Isomerism: Counting total geometrical and optical isomers (especially for $M(A-A)_2B_2$ and $M A_3 B_3$ complexes). Identifying fac/mer and cis/trans configurations, and knowing that trans and square planar complexes are optically inactive.
• VBT & Magnetic Moments: Predicting geometry ($sp^3$ vs $dsp^2$ vs $sp^3d^2$ vs $d^2sp^3$) and calculating spin-only magnetic moments. Distinguishing inner vs outer orbital complexes.
• Crystal Field Theory (CFT): Calculating CFSE (Crystal Field Stabilization Energy), arranging ligands based on the Spectrochemical Series, and relating $\Delta_o$ to pairing energy ($P$) to predict high/low spin state.
• Synergic Bonding: Understanding how back-bonding affects the C-O bond length and stretching frequency in IR spectroscopy (more back-bonding = weaker C-O bond = longer C-O bond length).
• Applications: Identifying Wilkinson's catalyst, Cis-platin, Hypo solution reactions, and the role of EDTA in volumetric analysis and lead poisoning.