Class 12 Chemistry · NCERT chapter 9 · 7 topics · free to read
Werner's Theory & Terminology
Alfred Werner proposed that metal ions have two types of valences: primary (ionizable, satisfied by anions) and secondary (non-ionizable, satisfied by ligands in a fixed geometry). This framework introduces coordination number, coordination sphere, ligands, denticity, and the distinction between ionizable and non-ionizable species in a complex.
[Co(NH_3)_6]^ 3+ C.N. = 6 (octahedral)
Werner's Postulates
Werner proposed that metal ions show two kinds of valence: primary valence (ionizable, now called oxidation state) and secondary valence (non-ionizable, now called coordination number). Secondary valence is directional and fixed for a given metal ion.
In 1893, Alfred Werner proposed a revolutionary theory to explain the puzzling behavior of coordination compounds. Before Werner, chemists couldn't explain why CoCl_3 6NH_3 behaved differently from CoCl_3 5NH_3 in solution — both had the same composition type, yet gave different numbers of ions.
Primary valence = ionizable valence = oxidation state of the central metal ion (satisfied by negative ions)
Secondary valence = non-ionizable valence = coordination number (satisfied by ligands, which can be neutral or anionic)
Secondary valence is fixed for a given metal ion (e.g., Co³⁺ almost always has secondary valence = 6)
Secondary valence is directional — ligands arrange in a definite geometry around the metal
Werner correctly predicted octahedral geometry for coordination number 6 and square planar/tetrahedral for 4
Central Metal Ion
The central metal atom or ion is a Lewis acid (electron pair acceptor) that bonds to surrounding ligands. It is typically a transition metal with empty d-orbitals available for bonding.
The central metal atom/ion sits at the center of the coordination compound and acts as a Lewis acid — it has empty orbitals ready to accept electron pairs from ligands. Transition metals are the most common central atoms because they have partially filled d -orbitals, variable oxidation states, and a strong tendency to form complexes.
Central atom is a Lewis acid (electron pair acceptor)
Most commonly a transition metal (e.g., Fe, Co, Ni, Cu, Pt, Cr)
Can also be a main-group metal: [AlF_6]^ 3- , [SiF_6]^ 2-
The charge on the central metal = total complex charge minus sum of all ligand charges
Common oxidation states in complexes: Co(III), Fe(II), Fe(III), Cr(III), Pt(II), Pt(IV), Ni(II), Cu(II)
Quick recall
What is the oxidation state of Fe in [Fe(CN)_6]^ 4- ?
Each CN^- has charge -1 , so total ligand charge = 6 (-1) = -6 . Complex charge = -4 . So Fe oxidation state = -4 - (-6) = +2 . It's Fe^ 2+ .
Ligands & Denticity
Ligands are Lewis bases (electron pair donors) that bond to the central metal. They are classified by the number of donor atoms: monodentate (1), bidentate (2), polydentate (multiple). Chelating ligands form ring structures with the metal.
A ligand is any molecule or ion that has at least one lone pair of electrons and can donate it to the central metal ion. Ligands are Lewis bases. The number of donor atoms a ligand uses to bind to the metal is called its denticity.
Ambidentate ligands can bind through different atoms: NO_2^- (N-bonded = nitro, O-bonded = nitrito), SCN^- (S-bonded = thiocyanato, N-bonded = isothiocyanato)
1 donor atom per ligand
2 or more donor atoms per ligand
No ring formed
Ring structure formed with metal
Less stable complexes
More stable (chelate effect)
Example: [Co(NH_3)_6]^ 3+
Example: [Co(en)_3]^ 3+
6 separate ligands for C.N. 6
Fewer ligands needed for same C.N.
Coordination Number & Coordination Sphere
Coordination number (C.N.) is the total number of donor atoms directly bonded to the central metal. The coordination sphere (inside square brackets) contains the metal and all directly bonded ligands. Species outside the bracket are counter ions.
The coordination number (C.N.) is the number of ligand donor atoms directly bonded to the central metal ion. It is NOT simply the number of ligands — a bidentate ligand contributes 2 to the C.N. The coordination sphere is everything inside the square brackets [ ] in the formula.
C.N. = (number of ligands denticity of each ligand)
C.N. = total number of donor atoms bonded to central metal, NOT the number of ligands
[Co(en)_3]^ 3+ : 3 bidentate ligands 2 = C.N. = 6
[PtCl_4]^ 2- : 4 monodentate ligands 1 = C.N. = 4
Common C.N. values: 4 (tetrahedral or square planar), 6 (octahedral)
Species outside [ ] are counter ions — they dissociate in solution and are not part of the coordination sphere
Quick recall
In [CoCl_2(en)_2]^+ , what is the coordination number of Co?
The charge on the coordination sphere equals the sum of the central metal's oxidation state and all ligand charges. Counter ions outside the brackets balance this charge to give an electrically neutral compound.
Every coordination compound is electrically neutral overall. The charge on the coordination sphere (inside brackets) is balanced by counter ions (outside brackets). To find the oxidation state of the central metal, set up: metal charge + sum of ligand charges = charge on complex ion.
O.S. of metal + (charge of each ligand) = charge on complex ion
K_4[Fe(CN)_6] : complex charge = -4 , Fe + 6(-1) = -4 , so Fe = +2
[Co(NH_3)_5Cl]Cl_2 : 2 Cl^- outside means complex charge = +2 ; Co + 5(0) + (-1) = +2 , so Co = +3
[Pt(NH_3)_2Cl_2] : neutral complex (no counter ions), Pt + 2(0) + 2(-1) = 0 , so Pt = +2
IUPAC Nomenclature
IUPAC naming of coordination compounds follows strict rules: ligands are named alphabetically before the metal, anionic ligands get '-o' suffix, neutral ligands keep their names (except aqua, ammine, carbonyl, nitrosyl), and the metal gets '-ate' suffix in anionic complexes. Oxidation state is shown in Roman numerals in parentheses.
A systematic set of rules governs the naming of coordination compounds. The order is: cation before anion (for ionic complexes), ligands alphabetically before metal, and oxidation state in Roman numerals.
Naming coordination compounds is like following a recipe — get the steps right and you'll never go wrong. The IUPAC system ensures every complex has exactly one unambiguous name.
Rule 1: Name the cation first, then the anion (same as any ionic compound: sodium chloride, not chloride sodium)
Rule 2: Within the coordination sphere, name ligands first (alphabetically), then the metal
Rule 3: Ligand names are listed in alphabetical order (ignore prefixes like di-, tri-, tetra- when alphabetizing)
Rule 4: Use prefixes di-, tri-, tetra-, penta-, hexa- for simple ligands; bis-, tris-, tetrakis- for complex/already-prefixed ligands
Rule 5: Anionic ligands end in '-o': chlorido, bromido, cyanido, hydroxido, nitrito/nitro, oxalato
Rule 6: Neutral ligands keep their names EXCEPT: H_2O = aqua, NH_3 = ammine, CO = carbonyl, NO = nitrosyl
Rule 7: Metal in anionic complex gets '-ate' suffix (often Latin name): ferrate, cuprate, cobaltate, platinate, argentate
Rule 8: Metal in cationic/neutral complex keeps its normal name
Rule 9: Oxidation state of metal in Roman numerals in parentheses after metal name: cobalt(III), iron(II)
Rule 10: No space between ligand names; one space between last ligand name and metal name? No — actually no space within the complex name, and the entire coordination sphere name is one word
Anionic Ligand Names
Anionic ligands replace their '-ide' ending with '-o' (sometimes '-ato' for '-ate' endings). This is the most tested naming detail in NEET.
When a negative ion acts as a ligand, its name changes. The general pattern: '-ide' becomes '-ido', '-ite' becomes '-ito', '-ate' becomes '-ato'. (Older IUPAC used '-o' directly: chloro, bromo, cyano. Newer recommendations use '-ido': chlorido, bromido, cyanido. NEET accepts both.)
Practice is the only way to master nomenclature. Here are worked examples covering cationic, anionic, and neutral complexes with increasing complexity.
Let's name some complexes step by step. Follow the recipe: identify the coordination sphere, list ligands alphabetically with correct names, name the metal (with '-ate' if anionic), and add oxidation state.
Given an IUPAC name, write the formula by identifying the metal, ligands, charges, and counter ions. The formula convention: cation first, then anion; within brackets, metal first, then anionic ligands, then neutral ligands.
When writing the formula from a name, the convention inside brackets is: metal first, then anionic ligands (alphabetically by symbol), then neutral ligands (alphabetically by symbol). This is the OPPOSITE of naming order (where ligands come first).
Ligands listed BEFORE metal
Metal written BEFORE ligands
Ligands in alphabetical order of name
Anionic ligands first, then neutral (by symbol)
Prefixes: di, tri, tetra...
Subscripts: _2 , _3 , _4 ...
'-ate' for anionic complex
Counter cation written first: K_3[...]
Potassium trioxalatoferrate(III): K is counter ion, Fe(III) is metal, oxalato ( C_2O_4^ 2- ) 3 . Charge: 3(-2) + 3 = -3 . Formula: K_3[Fe(C_2O_4)_3]
In the formula, the cation is ALWAYS written first, whether or not it's the complex ion
Isomerism in Coordination Compounds
Coordination compounds show two broad types of isomerism: structural isomerism (ionization, linkage, coordination, solvate) where the connectivity differs, and stereoisomerism (geometrical cis-trans and optical) where connectivity is the same but spatial arrangement differs. Geometrical isomerism is common in square planar and octahedral complexes; optical isomerism requires non-superimposable mirror images.
cis-[Pt(NH_3)_2Cl_2] trans-[Pt(NH_3)_2Cl_2]
Structural Isomerism
Structural isomers have the same molecular formula but differ in which species are inside vs outside the coordination sphere, or which donor atom of an ambidentate ligand is bonded to the metal.
In structural isomerism, the bonding connectivity itself is different. There are four types relevant to coordination compounds, and NEET loves testing all four.
Ionization isomerism: Same formula, but different ions are inside/outside the coordination sphere. They give different ions in solution. Example: [Co(NH_3)_5Br]SO_4 (gives SO_4^ 2- test) vs [Co(NH_3)_5SO_4]Br (gives Br^- test)
Linkage isomerism: Ambidentate ligand bonds through different donor atoms. Example: [Co(NH_3)_5NO_2]^ 2+ (nitro, N-bonded) vs [Co(NH_3)_5ONO]^ 2+ (nitrito, O-bonded)
Coordination isomerism: Both cation and anion are complex ions, and ligands are exchanged between them. Example: [Co(NH_3)_6][Cr(CN)_6] vs [Cr(NH_3)_6][Co(CN)_6]
Solvate (hydrate) isomerism: Water molecules swap between coordination sphere and crystal lattice. Example: [Cr(H_2O)_6]Cl_3 (violet) vs [Cr(H_2O)_5Cl]Cl_2 H_2O (green)
Geometrical (cis-trans) Isomerism
Geometrical isomers have the same connectivity but differ in the spatial arrangement of ligands. In square planar and octahedral complexes, identical ligands can be adjacent (cis) or opposite (trans).
Geometrical isomerism arises when ligands can occupy different spatial positions relative to each other without breaking any bonds. The two arrangements are cis (same side/adjacent) and trans (opposite side).
Square planar [Ma_2b_2] : cis (both 'a' adjacent, 90°) and trans (both 'a' opposite, 180°). Classic example: cisplatin [Pt(NH_3)_2Cl_2]
Square planar [Mabcd] : 3 geometrical isomers possible
Square planar [Ma_2bc] : cis and trans isomers based on position of the two 'a' ligands
Octahedral [Ma_4b_2] : cis (both 'b' adjacent) and trans (both 'b' opposite). Example: [Co(NH_3)_4Cl_2]^+
Octahedral [Ma_3b_3] : fac (facial — three identical ligands on one triangular face) and mer (meridional — three identical ligands in one plane)
Tetrahedral complexes do NOT show geometrical isomerism — all positions are equivalent
Identical ligands on SAME side (90°)
Identical ligands on OPPOSITE side (180°)
Has a net dipole moment (usually)
Dipole moments cancel (zero or smaller)
Cisplatin: anticancer drug
Transplatin: biologically inactive
More soluble in polar solvents
Less soluble in polar solvents
Optical Isomerism
Optical isomers (enantiomers) are non-superimposable mirror images of each other. They rotate the plane of polarized light in opposite directions. Common in octahedral complexes with bidentate ligands and cis-square planar complexes with unsymmetric ligands.
A coordination compound is optically active if its mirror image is non-superimposable on the original — just like your left and right hands. The two mirror images are called enantiomers: one rotates plane-polarized light clockwise (dextro, d or + ) and the other rotates it anticlockwise (laevo, l or - ).
Optical isomers are non-superimposable mirror images (enantiomers)
They have identical physical properties (m.p., b.p., solubility) EXCEPT rotation of plane-polarized light
d (dextro/ + ) rotates light clockwise; l (laevo/ - ) rotates light anticlockwise
A 1:1 mixture of d and l forms = racemic mixture (optically inactive, rotations cancel)
Common optically active complexes: cis- [Co(en)_2Cl_2]^+ , [Co(en)_3]^ 3+ , [Cr(C_2O_4)_3]^ 3-
trans isomers of octahedral complexes generally have a plane of symmetry → NOT optically active
Only cis-octahedral complexes with bidentate ligands typically show optical isomerism
Quick recall
Does [Co(en)_3]^ 3+ show optical isomerism?
Yes! It has 3 bidentate en ligands in an octahedral arrangement. The complex has no plane of symmetry — its mirror image is non-superimposable. It exists as d and l enantiomers. This is one of the most frequently asked examples in NEET.
Bonding: Valence Bond Theory (VBT)
VBT explains bonding in coordination compounds through hybridization of the central metal's orbitals. The metal provides empty hybrid orbitals; ligands donate electron pairs into them. The type of hybridization (sp³, dsp², sp³d², d²sp³) determines geometry. Inner orbital complexes use (n-1)d orbitals and are usually low-spin; outer orbital complexes use nd orbitals and are high-spin.
In VBT, the central metal ion provides empty orbitals (after hybridization) and each ligand donates an electron pair into one of these empty hybrid orbitals. The hybrid orbitals determine the geometry of the complex.
Valence Bond Theory (VBT) for coordination compounds is an extension of the hybridization concept. The central metal ion hybridizes its atomic orbitals to create a set of equivalent empty hybrid orbitals. Each ligand then donates its lone pair into one of these empty orbitals, forming a coordinate (dative) bond.
Metal provides empty orbitals → Lewis acid
Ligand provides electron pairs → Lewis base
Bond formed = coordinate covalent bond (both electrons from ligand)
Inner orbital complexes use (n-1)d orbitals for hybridization and are formed with strong field ligands. Outer orbital complexes use nd orbitals and are formed with weak field ligands. This distinction directly determines magnetic properties.
The crucial question in VBT is: which d-orbitals does the metal use for hybridization? If the ligand is strong enough to force pairing of existing d-electrons (making inner d-orbitals available), we get an inner orbital complex. If not, the metal uses higher-energy outer d-orbitals → outer orbital complex.
Uses (n-1)d orbitals
Uses nd orbitals
Hybridization: d^2sp^3 (octahedral) or dsp^2 (sq. planar)
Hybridization: sp^3d^2 (octahedral) or sp^3 (tetrahedral)
Strong field ligands (CN⁻, CO, NH₃, en)
Weak field ligands (F⁻, Cl⁻, Br⁻, H₂O)
Electrons are PAIRED → fewer unpaired electrons
Electrons remain UNPAIRED → more unpaired electrons
Low spin (usually diamagnetic or weakly paramagnetic)
High spin (strongly paramagnetic)
Example: [Co(NH_3)_6]^ 3+ ( d^2sp^3 )
Example: [CoF_6]^ 3- ( sp^3d^2 )
Magnetic Properties from VBT
The magnetic moment of a coordination compound is determined by the number of unpaired electrons, which VBT predicts based on hybridization type. The spin-only formula gives magnetic moment in Bohr magnetons.
VBT's most powerful prediction is magnetic behavior. By determining the hybridization and filling of orbitals, we can count unpaired electrons and calculate the magnetic moment using the spin-only formula.
_ spin-only = n(n+2) BM
n = 0 → = 0 BM → Diamagnetic (repelled by magnetic field)
n = 1 → = 1.73 BM
n = 2 → = 2.83 BM
n = 3 → = 3.87 BM
n = 4 → = 4.90 BM
n = 5 → = 5.92 BM → Maximum for first-row transition metals
Quick recall
[Ni(CN)_4]^ 2- is diamagnetic but [NiCl_4]^ 2- is paramagnetic. Explain using VBT.
Ni^ 2+ is d^8 . With strong field CN^- : electrons pair up → dsp^2 hybridization → square planar → 0 unpaired electrons → diamagnetic. With weak field Cl^- : no pairing → sp^3 hybridization → tetrahedral → 2 unpaired electrons → paramagnetic ( = 2.83 BM).
Bonding: Crystal Field Theory (CFT)
CFT treats metal-ligand bonding as purely electrostatic. Ligands create an electric field that splits the d-orbitals into two energy levels. In octahedral complexes: t₂g (lower) and eg (higher), split by Δo. In tetrahedral complexes: e (lower) and t₂ (higher), split by Δt ≈ 4/9 Δo. The magnitude of splitting determines whether electrons pair (strong field/low spin) or remain unpaired (weak field/high spin), explaining color and magnetic properties.
_t = 4 9 _o
CFT Basics
Crystal Field Theory assumes ligands are point negative charges (or dipoles) that interact electrostatically with the d-orbitals of the central metal ion. This interaction splits the five degenerate d-orbitals into two groups of different energies.
Crystal Field Theory (CFT) takes a completely different approach from VBT. Instead of covalent bonding, CFT treats the metal-ligand interaction as purely electrostatic — ligands are just negative point charges (for anions) or point dipoles (for neutral molecules like NH_3 , H_2O ). These negative charges create an electric field around the metal ion.
CFT assumes purely electrostatic (ionic) interaction — no covalent bonding
Ligands = negative point charges that repel d -electrons
d -orbitals pointing AT ligands → higher energy (more repulsion)
d -orbitals pointing BETWEEN ligands → lower energy (less repulsion)
The energy gap between the two sets = crystal field splitting energy ( )
CFT successfully explains: color, magnetic properties, and stability of complexes
Octahedral Crystal Field Splitting
In an octahedral field, 6 ligands approach along the x, y, and z axes. The d-orbitals split into a lower-energy set t₂g (dxy, dyz, dxz) and a higher-energy set eg (dx²-y², dz²). The energy gap is Δo (crystal field splitting energy for octahedral).
In an octahedral complex, 6 ligands sit along the x , y , and z axes. The d -orbitals split into two groups:
t_ 2g set (lower energy): d_ xy , d_ yz , d_ xz — these point BETWEEN the axes (between ligands) → less repulsion → stabilized
e_g set (higher energy): d_ x^2-y^2 , d_ z^2 — these point ALONG the axes (directly at ligands) → more repulsion → destabilized
Energy gap = _o (subscript 'o' for octahedral)
t_ 2g is lowered by -0.4 _o per electron (relative to average energy)
e_g is raised by +0.6 _o per electron (relative to average energy)
The average energy of the five d -orbitals remains unchanged (barycentre rule): 3(-0.4 _o) + 2(+0.6 _o) = 0
CFSE = (-0.4n_ t_ 2g + 0.6n_ e_g ) _o + P p
where n_ t_ 2g , n_ e_g = electrons in each set; P = pairing energy; p = number of forced pairings
Tetrahedral Crystal Field Splitting
In a tetrahedral field, the splitting pattern is inverted compared to octahedral: e set is lower and t₂ set is higher. The splitting energy Δt is only about 4/9 of Δo, so tetrahedral complexes are almost always high-spin.
In a tetrahedral complex, 4 ligands approach from alternate corners of a cube. Now the d -orbitals that point BETWEEN the axes ( d_ xy , d_ yz , d_ xz ) are CLOSER to the ligands, and the ones pointing along axes ( d_ x^2-y^2 , d_ z^2 ) are farther away. The splitting is inverted compared to octahedral.
e set (lower energy): d_ x^2-y^2 , d_ z^2
t_2 set (higher energy): d_ xy , d_ yz , d_ xz
Splitting is inverted compared to octahedral
_t = 4 9 _o — tetrahedral splitting is only ~44% of octahedral
Because _t is small, it rarely exceeds pairing energy → almost all tetrahedral complexes are high-spin
No 'g' subscript in tetrahedral labels ( e and t_2 , not e_g and t_ 2g ) because tetrahedra lack a center of symmetry (no inversion center → no 'gerade')
_t = 4 9 _o 0.44 _o
Quick recall
Why are tetrahedral complexes almost always high-spin?
Because _t 4 9 _o is so small that it almost never exceeds the pairing energy P . Since _t < P for almost all cases, electrons prefer to occupy higher orbitals (Hund's rule) rather than pair up. Hence, high-spin.
Spectrochemical Series
The spectrochemical series arranges ligands in order of increasing crystal field splitting energy (Δ). Strong field ligands (CN⁻, CO, en, NH₃) cause large splitting → low spin. Weak field ligands (I⁻, Br⁻, Cl⁻, F⁻, H₂O) cause small splitting → high spin.
Not all ligands split the d -orbitals by the same amount. The spectrochemical series is an experimentally determined ordering of ligands from weakest to strongest field (smallest to largest ).
Strong field ligands (large ): NH_3 , en, NO_2^- , CN^- , CO → low spin
If > P (pairing energy) → electrons pair up → low spin
If < P → electrons fill higher orbitals → high spin
H_2O is a borderline ligand — it's at the middle of the series
Color of Coordination Compounds
The color of transition metal complexes arises from d-d transitions: an electron absorbs visible light of energy equal to Δ and jumps from the lower to the upper set of d-orbitals. The observed color is complementary to the absorbed color.
The beautiful colors of coordination compounds — the deep blue of [Cu(NH_3)_4]^ 2+ , the yellow of K_2[CrO_4] , the violet of [Ti(H_2O)_6]^ 3+ — all arise from d - d transitions. When white light hits the complex, a photon of energy exactly equal to is absorbed, promoting an electron from the lower set to the upper set of d -orbitals.
= h = hc _ absorbed
Electron absorbs a photon of energy = and jumps from t_ 2g to e_g (in octahedral)
The color we see is the complementary color of the absorbed light
Absorbs red → appears green/blue. Absorbs violet → appears yellow.
Stability & Applications of Coordination Compounds
The stability of a coordination compound is quantified by its stability constant (formation constant) Kf — the larger the Kf, the more stable the complex. Chelate complexes are more stable than analogous non-chelate complexes (chelate effect). Coordination compounds have widespread applications in analytical chemistry (qualitative/quantitative analysis), medicine (cisplatin, EDTA therapy), extraction of metals, electroplating, and photography.
K_f = [ML_n] [M][L]^n
Stability Constant ( K_f )
The stability constant (or formation constant) Kf is the equilibrium constant for the formation of a complex from the metal ion and ligands. A larger Kf means a more stable complex that exists predominantly in the associated form in solution.
When a metal ion M reacts with ligands L to form a complex [ML_n] , the equilibrium constant for this reaction is the stability constant (also called formation constant), K_f .
M + nL leftharpoons [ML_n] K_f = [ML_n] [M][L]^n
Larger K_f → more stable complex → equilibrium lies far to the right
Smaller K_f → less stable → complex dissociates easily
The reciprocal 1 K_f is called the instability constant (or dissociation constant)
K_f for [Cu(NH_3)_4]^ 2+ = 10^ 12 — extremely stable
K_f for [Fe(CN)_6]^ 4- ≈ 10^ 35 — one of the most stable complexes known
Complexes with chelating (polydentate) ligands are significantly more stable than analogous complexes with monodentate ligands. This is the chelate effect, driven primarily by a favorable entropy change — releasing many monodentate ligands when fewer polydentate ligands bind.
The chelate effect is the observation that complexes containing chelate rings (from polydentate ligands) are much more stable than similar complexes with monodentate ligands. For example, [Ni(en)_3]^ 2+ is far more stable than [Ni(NH_3)_6]^ 2+ , even though both have 6 Ni-N bonds.
Chelate complexes are thermodynamically MORE stable (larger K_f ) than non-chelate analogues
Main driving force: entropy — replacing 6 monodentate ligands with 3 bidentate ones releases more free molecules → S > 0
Five-membered and six-membered chelate rings are most stable (least ring strain)
EDTA (hexadentate) forms extremely stable chelate complexes — forms 5 five-membered rings with one metal ion!
The more chelate rings formed, the greater the stability
Applications
Coordination compounds have diverse applications: analytical chemistry (hardness testing, qualitative analysis), medicine (cisplatin, EDTA chelation therapy), metallurgy (cyanide leaching of gold/silver), electroplating, and photography.
Coordination compounds aren't just textbook curiosities — they are central to medicine, industry, biology, and analytical chemistry. Here are the most important applications that NCERT covers and NEET tests.
Analytical Chemistry: EDTA titrations for water hardness ( Ca^ 2+ , Mg^ 2+ ). DMG (dimethylglyoxime) — test for Ni^ 2+ (bright red precipitate). K_4[Fe(CN)_6] — test for Fe^ 3+ (Prussian blue)
Extraction of metals: Gold and silver extraction using cyanide process: 4Au + 8CN^- + O_2 + 2H_2O 4[Au(CN)_2]^- + 4OH^-
Electroplating: Silver plating uses K[Ag(CN)_2] bath. Gold plating uses K[Au(CN)_2] . Complexing the metal gives a smooth, uniform deposit
Photography: Black & white photography uses AgBr crystals. The 'fixer' ( Na_2S_2O_3 , hypo) dissolves unreacted AgBr by forming [Ag(S_2O_3)_2]^ 3-
Why is EDTA used in chelation therapy for lead poisoning?
EDTA is hexadentate — it wraps around Pb^ 2+ with 6 donor atoms, forming an extremely stable chelate complex ( K_f very large). This complex is water-soluble and non-toxic, so it's safely excreted by the kidneys, removing the toxic lead from the body.
Organometallic Compounds
Organometallic compounds contain at least one metal-carbon bond. Metal carbonyls (e.g., Ni(CO)₄, Fe(CO)₅) are the most important class for NEET. The M-CO bond involves synergic bonding: CO donates a lone pair to the metal (σ-bond) and the metal back-donates electron density from filled d-orbitals into empty π* orbitals of CO (π-back bonding). The 18-electron rule predicts stability of metal carbonyls.
Ni(CO)_4: Ni(28) = 10e^- + 4 2e^- = 18e^-
Organometallic Compounds
An organometallic compound has at least one direct metal-carbon (M-C) bond. They are classified by the nature of the M-C bond: σ-bonded (e.g., Grignard reagents), π-bonded (e.g., ferrocene), and σ- and π-bonded (e.g., metal carbonyls).
An organometallic compound is defined as a compound containing at least one metal-carbon (M-C) bond. This definition is strict: metal cyanides ( KCN ) and metal carbides ( CaC_2 ) are borderline cases but are generally NOT considered organometallic because the carbon is part of an ionic lattice, not a direct covalent M-C bond in the traditional sense.
-bonded: Ferrocene [Fe( ^5-C_5H_5)_2] , Zeise's salt K[PtCl_3( ^2-C_2H_4)]
and -bonded: Metal carbonyls like Ni(CO)_4 , Fe(CO)_5
Metal carbonyls are the most important class for NCERT/NEET
Metal Carbonyls
Metal carbonyls are complexes where CO (carbon monoxide) is the ligand. They can be mononuclear (one metal atom) or polynuclear (multiple metal atoms bridged by CO). The bonding involves synergic interaction between the metal and CO.
Metal carbonyls are coordination compounds where CO molecules are the ligands. CO is a remarkably versatile ligand — it can act as a terminal ligand (bonded to one metal) or a bridging ligand (connecting two metals, written as -CO).
In Fe_2(CO)_9 : 3 bridging CO groups ( -CO) + 6 terminal CO groups
Metal carbonyls are typically diamagnetic (all electrons paired)
Ni(CO)_4 is a volatile, toxic, colorless liquid — used in the Mond process for nickel purification
Quick recall
What is the structure of Fe(CO)_5 ?
Trigonal bipyramidal. Fe has 8 electrons ( d^8 ), gains 10 from 5 CO ligands → 18 electrons total (satisfies 18-electron rule). The 5 CO groups arrange in a trigonal bipyramidal geometry.
Synergic Bonding
In M-CO bonding, two simultaneous processes occur: (1) CO donates a lone pair to the metal via σ-bond (M←CO), and (2) the metal donates electron density from its filled d-orbitals into the empty π* antibonding orbitals of CO via π-back bonding (M→CO). These two processes reinforce each other — hence 'synergic'.
The M-CO bond is special because it involves two-way electron traffic that mutually strengthens each other. This is called synergic bonding (synergic = working together).
Step 1 ( -donation): The lone pair on carbon of CO is donated INTO an empty metal orbital → M arrow CO
Step 2 ( -back donation): Filled d -orbitals of the metal donate electron density INTO the empty ^* (antibonding) orbital of CO → M arrow CO
The -donation increases electron density on the metal, which ENHANCES the -back donation
The -back donation removes excess electron density from the metal, which ENHANCES the -donation
Net result: a very strong M-CO bond and a weakened C-O bond (bond order of CO decreases from 3 in free CO)
18-Electron Rule
The 18-electron rule states that stable metal carbonyls tend to have 18 electrons in the metal's valence shell (counting metal d-electrons + electrons donated by ligands). This corresponds to filling all 9 valence orbitals (one s + three p + five d).
Just as main-group elements follow the octet rule (8 electrons), transition metal complexes often follow the 18-electron rule. The metal has 9 valence orbitals ( 1s + 3p + 5d ), and when all 9 are filled with 2 electrons each, you get 18 electrons — a stable, noble-gas-like configuration.
Total valence electrons = metal d-electrons + electrons donated by ligands
Ni(CO)_4 : Ni = d^ 10 (10 electrons) + 4 CO 2 = 8 electrons → 18 electrons (stable, exists!)
Fe(CO)_5 : Fe = d^ 8 (8 electrons) + 5 CO 2 = 10 electrons → 18 electrons (stable, exists!)
V(CO)_6 : V = d^ 5 (5 electrons) + 6 CO 2 = 12 electrons → 17 electrons (exists but is paramagnetic and less stable — dimerizes to get to 18e)
The 18-electron rule works best for metal carbonyls and organometallics, less reliably for classical Werner complexes
Quick recall
Predict the formula of manganese carbonyl if it obeys the 18-electron rule.
Mn is group 7 → 7 valence electrons. Need 18 - 7 = 11 more electrons. But each CO gives 2, and 11 is odd — can't get 18 with CO alone. Solution: Mn forms a Mn-Mn bond (contributes 1e each) → Mn_2(CO)_ 10 . Each Mn: 7 + 1 (Mn-Mn) + 5(CO) 2 = 18.