Coordination Compounds

01

🧪

Introduction to Coordination Compounds

Coordination compounds (also called complex compounds) are formed when a central metal atom or ion is bonded to a fixed number of ions or neutral molecules (ligands) through coordinate covalent bonds. They are ubiquitous in nature — from haemoglobin to chlorophyll — and form the backbone of industrial catalysis, medicine, and analytical chemistry.

Definition

A coordination compound is a compound containing a coordination entity — a central metal atom/ion surrounded by ligands that donate electron pairs to form coordinate bonds. These are written with the coordination sphere enclosed in square brackets.

[Central Metal + Ligands]^{charge} → Coordination Entity/Sphere

Historical Background

Alfred Werner (1893) first explained the nature of bonding in coordination compounds. He was awarded the Nobel Prize in Chemistry in 1913 for this contribution — the first Swiss chemist to receive it. Prior to his work, coordination compounds like cobalt ammines were known but their structure was a mystery.

Importance

Biological role: O₂ transport (haemoglobin), photosynthesis (chlorophyll), vitamin B₁₂
Catalysis: Ziegler–Natta catalyst, Wilkinson's catalyst
Medicine: Cisplatin (anticancer), EDTA chelation therapy
Electroplating, photography, pigments, extraction of metals

02

📖

Key Terminology

Central Metal Atom / Ion

The atom or ion at the center of the coordination entity. Accepts electron pairs from ligands. Usually a transition metal (d-block), though p-block and f-block metals also form complexes.

Must have: low charge, small size, vacant orbitals to accept electron pairs from ligands.

Examples: Fe²⁺ Co³⁺ Pt²⁺ Cu²⁺

Ligands

Ions or neutral molecules that donate one or more lone pairs to the central metal. They must have at least one lone pair of electrons (Lewis bases).

Based on donor atoms:

Unidentate: 1 donor atom — NH₃, Cl⁻, H₂O, CN⁻, CO
Bidentate: 2 donors — en, ox²⁻, acac⁻
Polydentate: ≥3 donors — EDTA (hexadentate)
Ambidentate: 2 possible donors — NO₂⁻ (N or O), SCN⁻ (S or N)

Coordination Number (CN)

Total number of coordinate bonds formed by the ligands with the central metal (= total donor atoms, not number of ligands).

For bidentate ligand like en: each ligand counts as 2 towards CN.

CN = Σ (denticity of each ligand)

Common CN values: CN = 2 CN = 4 CN = 6

Coordination Sphere & Counter Ions

The coordination sphere consists of the central metal + its bound ligands, enclosed in square brackets. Species outside the brackets are counter ions that balance overall charge.

[Co(NH₃)₆]Cl₃

Here [Co(NH₃)₆]³⁺ = coordination sphere; 3 Cl⁻ = counter ions. The complex ionizes to give the coordination sphere intact in solution.

Chelate Complexes & Chelate Effect

Chelation

When a polydentate ligand uses two or more of its donor atoms to bind to the same central metal, the resulting ring structure is called a chelate ring. This process is called chelation and the ligand is a chelating agent.

The chelate effect makes chelated complexes more stable than those with equivalent unidentate ligands — driven by entropy increase (more particles released during reaction).

Example: [Cu(en)₂]²⁺ is far more stable than [Cu(NH₃)₄]²⁺ because of the chelate effect. EDTA forms extremely stable chelates with many metal ions.

Oxidation Number of Central Metal

The oxidation state of the central metal is determined by assigning charges to all ligands and counter ions and applying electrical neutrality.

Example — Finding Oxidation State

[Fe(CN)₆]⁴⁻ → Fe + 6(CN⁻) = −4

Fe + 6(−1) = −4 → Fe = +2. Iron is in the +2 oxidation state.

Example 2

[CrCl₂(NH₃)₄]⁺ → Cr + 2(−1) + 4(0) = +1

Cr = +3

03

⚗️

Double Salts vs. Complex Salts

DOUBLE SALTS

Exist only in solid state (crystalline form)

Lose identity in solution — dissociate completely into constituent ions

No new properties beyond constituent salts

Example: Mohr's salt — FeSO₄·(NH₄)₂SO₄·6H₂O

Example: Potash alum — K₂SO₄·Al₂(SO₄)₃·24H₂O

In solution: gives Fe²⁺, NH₄⁺, SO₄²⁻ freely

VS

COMPLEX SALTS

Retain identity in solution — coordination sphere stays intact

Show new, characteristic properties not shown by constituent parts

Complex ion does not dissociate to give free metal ions in solution

Example: K₄[Fe(CN)₆] → gives K⁺ and [Fe(CN)₆]⁴⁻ only

Example: [Cu(NH₃)₄]SO₄ → deep blue colour, distinct from Cu²⁺ alone

Stability quantified by formation/stability constant K_f

ℹ️

Key Distinction: In a double salt like Mohr's salt, adding NaOH gives a greenish precipitate of Fe(OH)₂ — indicating free Fe²⁺ in solution. In K₄[Fe(CN)₆], adding NaOH does NOT give Fe(OH)₂ — the iron is locked inside the coordination sphere.

04

🔬

Homoleptic & Heteroleptic Complexes

Homoleptic Complexes

All ligands attached to the central metal are of the same type.

[Co(NH₃)₆]³⁺ — all NH₃ ligands

[Fe(CN)₆]⁴⁻ — all CN⁻ ligands

[Cr(H₂O)₆]³⁺ — all H₂O ligands

These complexes generally do not show geometric or optical isomerism due to their higher symmetry.

Heteroleptic Complexes

The central metal is bound to two or more different types of ligands.

[CoCl₂(NH₃)₄]⁺ — Cl⁻ and NH₃

[Pt(NH₃)₂Cl₂] — NH₃ and Cl⁻ (cisplatin!)

[Fe(CO)₄Cl₂] — CO and Cl⁻

These can show geometric isomerism (cis/trans) and optical isomerism, which is crucial for pharmacological activity (e.g., cisplatin vs transplatin).

05

📝

IUPAC Nomenclature of Coordination Compounds

General Rules — Naming Order

For the salt: name the cation first, then the anion (same as simple salts).
For the coordination sphere: name ligands first (alphabetically), then the central metal with oxidation state.

1

Cation before Anion

Name the cationic part first, followed by the anionic part. If the complex ion is cationic, counter anion is named last (and vice versa).

2

Anionic Ligands — use "-o" suffix

Anionic ligands end in -o. Chloro (Cl⁻), Bromo (Br⁻), Fluoro (F⁻), Cyano (CN⁻), Hydroxo (OH⁻), Oxalato (ox²⁻), Nitro (NO₂⁻ via N), Nitrito (NO₂⁻ via O), Thiocyanato (SCN⁻ via S), Isothiocyanato (SCN⁻ via N), Sulfato (SO₄²⁻), Carbonato (CO₃²⁻).

3

Neutral Ligands — special names

H₂O = aqua, NH₃ = ammine (double m), CO = carbonyl, NO = nitrosyl, N₂H₄ = hydrazine, ethylenediamine = ethane-1,2-diamine. All other neutral ligands use their IUPAC names.

4

Alphabetical Order of Ligands

Ligands are named in alphabetical order (ignoring multiplying prefixes: di, tri, bis, tris). ammine comes before chloro alphabetically (a before c).

5

Multiplying Prefixes

Use di, tri, tetra, penta, hexa for simple ligands. Use bis, tris, tetrakis for ligands whose names contain a multiplying prefix (e.g. ethane-1,2-diamine, dimethylglyoximate). This prevents ambiguity.

6

Central Metal Name + Oxidation State

For cationic/neutral complexes: use the regular element name. For anionic complexes: use the Latin name with -ate suffix. E.g., iron → ferrate, copper → cuprate, gold → aurate, silver → argentate, cobalt → cobaltate.

7

Oxidation State in Roman Numerals

The oxidation state of the central metal is given in parentheses in Roman numerals immediately after the metal name: iron(III), cobalt(II), platinum(IV).

Worked Examples

Example 1

[Co(NH₃)₆]Cl₃

Name: hexaamminecobalt(III) chloride
Ligands: 6 NH₃ (ammine × 6 = hexaammine) | Metal: Co(+3) = cobalt(III) | Anion: Cl⁻ = chloride

Example 2

K₂[PtCl₄]

Name: potassium tetrachloridoplatinate(II)
Cation: K⁺ = potassium | Complex anion: [PtCl₄]²⁻ = tetrachloridoplatinate(II)

Example 3

[CrCl₂(NH₃)₄]Cl

Name: tetraamminedichloridochromium(III) chloride
Ligands in alphabetical order: ammine (a) before chlorido (c) → tetraammine + dichloro | Cr = +3

Example 4

K₃[Fe(CN)₆]

Name: potassium hexacyanidoferrate(III)
Anionic complex → iron becomes ferrate | CN⁻ = cyanido | Fe = +3

Example 5 (Bis prefix)

[Co(en)₂Cl₂]⁺

Name: bis(ethane-1,2-diamine)dichloridocobalt(III)
"en" contains "diamine" so use "bis" instead of "di"

Example 6 (Ambidentate)

[Co(NO₂)(NH₃)₅]²⁺

If N-bonded: pentaamminenitro-N-cobalt(III) — If O-bonded: pentaamminenitrito-O-cobalt(III)

⚠️

Common mistake: "ammine" (with double m) for NH₃ as ligand, vs "amine" for organic amines. Never write "tetramine cobalt"; the correct form is "tetraamminecobalt".

06

🏛️

Werner's Theory of Coordination Compounds

Werner's Postulates (1893)

Metals possess two types of valency: Primary valency (ionisable) and Secondary valency (non-ionisable).
Primary valency is the normal valency / oxidation state of the metal. It is satisfied by negative ions and is ionizable.
Secondary valency = Coordination number. It is satisfied by neutral molecules or negative ions. It is non-ionisable and directional (determines geometry).
Every metal tries to satisfy both its primary and secondary valencies.
Ligands satisfying both primary AND secondary valency are written inside the bracket; those satisfying only primary valency are written outside.

Werner's Cobalt Ammines — Classic Example

Complex	Formula	Colour	Ionisable Cl⁻ (by AgNO₃)	Non-Ionisable Cl⁻	CN
Luteo cobalt	[Co(NH₃)₆]Cl₃	Yellow	3	0	6
Purpureo cobalt	[CoCl(NH₃)₅]Cl₂	Purple	2	1	6
Roseo cobalt	[CoCl₂(NH₃)₄]Cl	Red	1	2	6
Violeo cobalt	[CoCl₃(NH₃)₃]	Green/Violet	0	3	6

Werner observed that conductivity of these solutions correlated with the number of ions in solution — confirming that coordination number = 6 in all four complexes.

Evidence for Werner's Theory

Conductivity experiments: Molar conductivity proportional to number of ions → confirms ionisable vs non-ionisable nature.
Precipitation with AgNO₃: Only ionisable Cl⁻ precipitates immediately → confirms inner/outer sphere distinction.
Optical isomerism in cobalt complexes: Werner synthesized optically active [Co(en)₃]³⁺ without any carbon — first proof of octahedral geometry.
X-ray crystallography later confirmed all Werner's structural predictions.

07

🔗

Valence Bond Theory (VBT)

Assumptions of VBT

The central metal atom/ion provides empty hybrid orbitals to accept lone pairs from ligands.
The type of hybridization determines the geometry and magnetic properties of the complex.
Each metal-ligand bond is a σ-bond (coordinate covalent) formed by overlap of filled ligand orbital with empty metal hybrid orbital.
Inner orbital complexes use inner d-orbitals (n-1)d → inner orbital/low spin.
Outer orbital complexes use outer d-orbitals nd → outer orbital/high spin.

Hybridization Schemes & Geometries

Hybridization	CN	Geometry	Example
`sp`	2	Linear	[Ag(NH₃)₂]⁺, [CuCl₂]⁻
`sp²`	3	Trigonal Planar	[HgI₃]⁻, [CuCl₃]²⁻
`sp³`	4	Tetrahedral	[Ni(CO)₄], [FeCl₄]²⁻, [CoCl₄]²⁻
`dsp²`	4	Square Planar	[Ni(CN)₄]²⁻, [Pt(NH₃)₂Cl₂], [Cu(NH₃)₄]²⁺
`sp³d`	5	Trigonal Bipyramidal	[Fe(CO)₅]
`dsp³`	5	Square Pyramidal	[Ni(CN)₅]³⁻
`d²sp³`	6	Octahedral (inner)	[Co(NH₃)₆]³⁺, [Fe(CN)₆]⁴⁻
`sp³d²`	6	Octahedral (outer)	[CoF₆]³⁻, [Fe(H₂O)₆]³⁺

Inner Orbital Complexes (Low Spin)

Use (n-1)d orbitals — inner d orbitals are used for hybridization (d²sp³). Formed with strong field ligands (CN⁻, CO, NO⁺, en, NH₃).

Pairing of electrons occurs — diamagnetic or low magnetic moment. Example: [Co(NH₃)₆]³⁺ — Co³⁺ (3d⁶) → all 6 d-electrons pair up → diamagnetic, d²sp³.

Outer Orbital Complexes (High Spin)

Use nd outer d orbitals for hybridization (sp³d²). Formed with weak field ligands (F⁻, Cl⁻, Br⁻, I⁻, H₂O, OH⁻).

Electrons remain unpaired — paramagnetic, high magnetic moment. Example: [CoF₆]³⁻ — Co³⁺ (3d⁶) retains unpaired electrons, sp³d² hybridization.

Worked VBT Examples

Example: [Fe(CN)₆]⁴⁻

Fe²⁺ = 3d⁶ · Strong field ligand CN⁻

Electrons pair in 3d: 3d⁶ → t₂g⁶ with pairing. Hybridization: d²sp³ (inner). Geometry: Octahedral. Magnetic moment: 0 BM (Diamagnetic). Low Spin

Example: [Fe(H₂O)₆]³⁺

Fe³⁺ = 3d⁵ · Weak field ligand H₂O

5 unpaired electrons remain. Hybridization: sp³d² (outer). Geometry: Octahedral. μ = √(5×7) = 5.92 BM (Paramagnetic). High Spin

Limitations of VBT

Cannot explain why certain ligands are strong-field and others weak-field.
Does not account for colour in coordination compounds.
Cannot predict the exact magnetic moment — ignores orbital contribution.
Cannot explain electronic spectra (multiple bands observed).
The concept of inner vs outer orbital is somewhat artificial — no clear explanation for the preference.

08

💎

Crystal Field Theory (CFT)

Basic Assumptions of CFT

Ligands are treated as point charges (if anions) or point dipoles (if neutral molecules).
The interaction between central metal and ligands is purely electrostatic — no covalent bonding considered.
The d-orbitals of the central metal, which are degenerate in free ion, are split by the electrostatic field of ligands.
The pattern of splitting depends on the geometry/symmetry of the complex.

Crystal Field Splitting — Octahedral Complexes

In an octahedral field (6 ligands along ±x, ±y, ±z axes), d-orbitals split into two sets based on their spatial orientation relative to the approaching ligands:

e_g Set — Higher Energy

d_x²-y² and d_z² orbitals point directly toward the ligands along the axes → maximum electrostatic repulsion → raised energy.

Energy raised by +0.6Δ₀ (each orbital) = +3/5 Δ₀

t_2g Set — Lower Energy

d_xy, d_xz, d_yz orbitals point between the ligands → less repulsion → stabilized (lower energy).

Energy lowered by −0.4Δ₀ (each orbital) = −2/5 Δ₀

🔑

Key Rule: The total energy of the system is conserved. The raising of e_g by 0.6Δ₀ × 2 = +1.2Δ₀ is balanced by the lowering of t_2g by 0.4Δ₀ × 3 = −1.2Δ₀. Net stabilization of a completely filled d¹⁰ or d⁰ system = 0.

Crystal Field Stabilization Energy (CFSE) — Calculation

CFSE Formula for Octahedral Complexes

CFSE (oct) = (−0.4 × n_t₂g + 0.6 × n_eₘ) × Δ₀

where n_t₂g = electrons in t_2g, n_eₘ = electrons in e_g

Pairing energy (P) must be considered when calculating the net CFSE compared to weak field. If number of forced pairings in complex > free ion, subtract the extra pairing energy.

Net CFSE = CFSE(oct) − Pairing penalty

CFSE Table for dⁿ Configurations (Octahedral)

dⁿ	High Spin (Weak Field)	t₂g	eₘ	CFSE (Δ₀)	Low Spin (Strong Field)	t₂g	eₘ	CFSE (Δ₀)
d⁰	0	0	0	0	0	0	0	0
d¹	t₂g¹	1	0	−0.4	t₂g¹	1	0	−0.4
d²	t₂g²	2	0	−0.8	t₂g²	2	0	−0.8
d³	t₂g³	3	0	−1.2	t₂g³	3	0	−1.2
d⁴	t₂g³ eₘ¹	3	1	−0.6	t₂g⁴	4	0	−1.6
d⁵	t₂g³ eₘ²	3	2	0	t₂g⁵	5	0	−2.0
d⁶	t₂g⁴ eₘ²	4	2	−0.4	t₂g⁶	6	0	−2.4
d⁷	t₂g⁵ eₘ²	5	2	−0.8	t₂g⁶ eₘ¹	6	1	−1.8
d⁸	t₂g⁶ eₘ²	6	2	−1.2	t₂g⁶ eₘ²	6	2	−1.2
d⁹	t₂g⁶ eₘ³	6	3	−0.6	t₂g⁶ eₘ³	6	3	−0.6
d¹⁰	t₂g⁶ eₘ⁴	6	4	0	t₂g⁶ eₘ⁴	6	4	0

⚡ Interactive CFSE Calculator

d-electron count (dⁿ)

Field Strength

Geometry

CFSE Value

Crystal Field Splitting — Tetrahedral Complexes

In a tetrahedral field, the d-orbital splitting is inverted compared to octahedral: the lower set is now called e and the upper set t₂.

e set (lower, t₂g equivalent)

d_z² and d_x²-y² — these are stabilized in a tetrahedral field. Energy = −0.6Δt each (= −6/10 Δt).

t₂ set (upper, eₘ equivalent)

d_xy, d_xz, d_yz — destabilized. Energy = +0.4Δt each (= +4/10 Δt).

Δt = (4/9) × Δ₀ ≈ 0.44 Δ₀

ℹ️

Why Δt < Δ₀? In tetrahedral geometry, (1) there are only 4 ligands instead of 6, and (2) none of the d-orbitals point directly at the ligands — they only approach along corners. Combined, these factors give Δt ≈ 4/9 Δ₀. Since Δt is small, the pairing energy usually exceeds Δt, so most tetrahedral complexes are high spin.

CFSE (tet) = (−0.6 × n_e + 0.4 × n_t₂) × Δt

Limitations of Crystal Field Theory

What CFT Cannot Explain

Assumes purely ionic/electrostatic bonding — ignores covalent character of metal-ligand bonds.
Cannot explain why neutral ligands like CO, NH₃ are strong field if bonding is purely electrostatic.
Cannot account for π-back bonding (back-donation from metal to ligand π* orbitals) — essential for CO complexes.
Does not explain intensities of electronic spectra satisfactorily.
Fails for complexes showing covalent metal-ligand bonds (organometallics).
The concept of spectrochemical series ordering is empirical — CFT provides no theoretical basis for it.

🔬

Ligand Field Theory (LFT) and Molecular Orbital Theory (MOT) address these limitations by incorporating covalent bonding and π-interactions. LFT is essentially CFT extended to include some covalent character.

09

🔄

Isomerism in Coordination Compounds

Isomers are compounds with the same molecular formula but different arrangements. Coordination compounds show a rich variety of isomerism, divided into two main categories:

🔵 Ionisation Isomerism Structural

Same molecular formula but the ligands and counter ions exchange positions — the coordination sphere changes, giving different ions in solution.

Example 1:
[Co(SO₄)(NH₃)₅]Br ↔ [CoBr(NH₃)₅]SO₄
First gives Br⁻ in solution; second gives SO₄²⁻.

Example 2:
[CrCl(H₂O)(en)₂]Br₂ ↔ [CrBr(H₂O)(en)₂]BrCl

💧 Hydrate (Solvate) Isomerism Structural

Water (or solvent) molecules may either be inside the coordination sphere as ligands or outside as free solvent molecules — giving different properties.

Three isomers of CrCl₃·6H₂O:
[Cr(H₂O)₆]Cl₃ — violet — 3 Cl⁻ ionisable
[CrCl(H₂O)₅]Cl₂·H₂O — blue-green — 2 Cl⁻
[CrCl₂(H₂O)₄]Cl·2H₂O — dark green — 1 Cl⁻

🔗 Linkage Isomerism Structural

Occurs with ambidentate ligands that can bond through two different donor atoms.

NO₂⁻:
[Co(NO₂)(NH₃)₅]SO₄ — nitro (N-bonded, yellow)
[Co(ONO)(NH₃)₅]SO₄ — nitrito (O-bonded, red-brown)

SCN⁻:
[Co(NCS)(NH₃)₅]²⁺ — isothiocyanato (N-bonded)
[Co(SCN)(NH₃)₅]²⁺ — thiocyanato (S-bonded)

🔀 Coordination Isomerism Structural

When both cation and anion are complex ions, ligands can be exchanged between them.

Example:
[Co(NH₃)₆][Cr(CN)₆] ↔ [Cr(NH₃)₆][Co(CN)₆]

Another Example:
[Cu(NH₃)₄][PtCl₄] ↔ [Pt(NH₃)₄][CuCl₄]

↕️ Geometric (Cis-Trans) Isomerism Stereoisomerism

Different spatial arrangement of ligands relative to each other. Cis: same ligands adjacent (90°). Trans: same ligands opposite (180°).

Square Planar [MA₂B₂]:
e.g. [Pt(NH₃)₂Cl₂] — cis (cisplatin, anticancer drug) and trans (transplatin, inactive)

Octahedral [MA₄B₂]:
e.g. [CoCl₂(NH₃)₄]⁺ — cis (violet) and trans (green)

Octahedral [MA₃B₃]:
fac (facial) — 3 identical ligands on one face of octahedron
mer (meridional) — 3 identical ligands in a plane

🌀 Optical Isomerism Stereoisomerism

Non-superimposable mirror images (enantiomers) that rotate plane-polarised light in opposite directions. The complex must lack an improper axis of symmetry (Sₙ — includes plane, centre, alternating axis).

Common in:
[Co(en)₃]³⁺ — Δ (delta, right-handed) and Λ (lambda, left-handed) forms
cis-[CoCl₂(en)₂]⁺ — shows optical isomerism
trans-[CoCl₂(en)₂]⁺ — NOT optically active (has plane of symmetry)

Tris-chelate octahedral complexes are always chiral.

🔑

IIT JEE Focus: For square planar complexes MA₂B₂ and MA₂BC: cis-trans possible. For MABCD square planar: 3 geometric isomers, some optically active. For octahedral MA₃B₃: fac and mer. For [M(en)₃]³⁺: only optical isomers (no geometric). Always draw structures to determine isomers!

10

⚖️

Stability of Coordination Compounds

Thermodynamic Stability — Formation Constant (K_f)

The stability of a complex in solution is quantified by the formation/stability constant K_f. A higher K_f means a more stable complex.

M + nL ⇌ [MLₙ] K_f = [[MLₙ]] / ([M][L]ⁿ)

K_f is also called the association constant. Its reciprocal is the dissociation/instability constant K_d = 1/K_f.

Stepwise formation constants: K₁, K₂, K₃ ... for adding one ligand at a time. Overall K_f = K₁ × K₂ × K₃ × ...

Factors Affecting Stability

Nature of the Metal

Higher charge of central metal → more stable complex (stronger electrostatic attraction)
Smaller ionic radius → higher charge density → more stable
Metals classified as class a (hard Lewis acids) prefer F, O, N donors; class b (soft Lewis acids) prefer P, S, I donors — HSAB principle
Transition metals: stability generally increases with atomic number (Irving-Williams series for divalent ions)

Nature of the Ligand

Chelate effect: Polydentate > unidentate (entropy-driven); 5/6-membered chelate rings most stable
Basicity of ligand: stronger Lewis base → more stable complex (generally)
Size: Smaller donor atoms form stronger bonds (less steric strain)
π-donor/acceptor ability modifies stability through back-bonding

Irving–Williams Series

For high-spin divalent transition metal complexes with the same ligand, stability follows the order:

Mn²⁺ < Fe²⁺ < Co²⁺ < Ni²⁺ < Cu²⁺ > Zn²⁺

This order is explained by CFSE and increasing nuclear charge across the series. Cu²⁺ has the highest K_f due to Jahn-Teller distortion providing additional stabilization. Zn²⁺ (d¹⁰) drops because CFSE = 0.

Kinetic Stability — Labile vs Inert Complexes

Distinct from thermodynamic stability — refers to the rate of ligand substitution reactions.

Labile complexes: Rapid ligand exchange (t½ < 1 min) — most high-spin octahedral and tetrahedral complexes (e.g., [Ni(H₂O)₆]²⁺)
Inert complexes: Slow ligand exchange (t½ > 1 min) — Cr³⁺, Co³⁺, Rh³⁺, Ir³⁺, Pt²⁺, Pt⁴⁺ complexes
Inertness is related to t₂g electron configuration — d³ and low-spin d⁶ complexes are most inert.

11

🔷

Coordination Number & Geometry

CN	Geometry	Hybridization	Examples	d-config needed
2	Linear	sp	[Ag(NH₃)₂]⁺, [CuCl₂]⁻, [Au(CN)₂]⁻	d¹⁰
3	Trigonal Planar	sp²	[HgI₃]⁻, [CuCl₃]²⁻	d¹⁰
4	Tetrahedral	sp³	[Ni(CO)₄], [CoCl₄]²⁻, [FeCl₄]²⁻, [BF₄]⁻	d⁰,d¹⁰,d⁵,d⁷ high spin
4	Square Planar	dsp²	[Ni(CN)₄]²⁻, [Pt(NH₃)₂Cl₂], [Cu(NH₃)₄]²⁺, [PdCl₄]²⁻	d⁸ (low spin)
5	Trigonal Bipyramidal	sp³d / dsp³	[Fe(CO)₅], [CuCl₅]³⁻	various
5	Square Pyramidal	dsp³	[Ni(CN)₅]³⁻, [VO(acac)₂]	d⁸
6	Octahedral	d²sp³ / sp³d²	[Co(NH₃)₆]³⁺, [Fe(CN)₆]⁴⁻, [Cr(H₂O)₆]³⁺, [PtCl₆]²⁻	Most transition metals
7	Pentagonal Bipyramidal	sp³d³	[ZrF₇]³⁻, [UO₂F₅]³⁻	f-block
8	Square Antiprismatic	d⁴sp³	[Mo(CN)₈]⁴⁻, [TaF₈]³⁻	f-block
12	Icosahedral	—	[Ce(NO₃)₆]²⁻ (bidentate NO₃⁻)	lanthanides

🔑

Square Planar Rule: d⁸ metal ions (Ni²⁺, Pd²⁺, Pt²⁺, Au³⁺, Rh⁺, Ir⁺) with strong field ligands adopt square planar geometry. The CFSE gained by forming square planar (from octahedral) is very high for d⁸, making it favourable despite loss of 2 ligands.

12

🎨

Colour in Coordination Compounds

Origin of Colour — d–d Transitions

Coordination compounds are coloured because the split d-orbitals allow electrons to absorb visible light and undergo d–d transitions (from t₂g to eₘ in octahedral, or e to t₂ in tetrahedral).

The energy of absorbed light (photon) = ΔE = hν = hc/λ = Δ₀ (or Δt). The complementary colour is observed.

E = hν = hc/λ = Δ₀ → Observed colour = Complementary of absorbed

d⁰ and d¹⁰ complexes have no possible d–d transitions → generally colourless/white (e.g., [Sc(H₂O)₆]³⁺ is colourless, [Zn(NH₃)₄]²⁺ is colourless).

Colour Wheel — Absorbed vs Observed

Classic Examples

Complex	Metal Config.	Colour Observed	Colour Absorbed	λ absorbed (nm)
[Ti(H₂O)₆]³⁺	d¹	Violet/Purple	Yellow-Green	~500
[V(H₂O)₆]³⁺	d²	Green	Red	~600
[Cr(H₂O)₆]³⁺	d³	Violet	Yellow-green	~575
[Mn(H₂O)₆]²⁺	d⁵ HS	Pale pink	Green	~520 (weak, spin-forbidden)
[Fe(H₂O)₆]²⁺	d⁶ HS	Pale green	Red	~1000
[Fe(H₂O)₆]³⁺	d⁵ HS	Yellow	Violet (weak)	~400
[Co(H₂O)₆]²⁺	d⁷ HS	Pink	Blue-green	~500
[Ni(H₂O)₆]²⁺	d⁸	Green	Red/Yellow	~650
[Cu(H₂O)₆]²⁺	d⁹	Blue	Orange-red	~600
[Sc(H₂O)₆]³⁺	d⁰	Colourless	None (no d-d)	—
[Zn(NH₃)₄]²⁺	d¹⁰	Colourless	None (no d-d)	—

Factors Affecting Colour (Δ)

Nature of ligand: Stronger field ligand → higher Δ₀ → shorter λ absorbed → colour shifts toward blue/violet. Weak field → lower Δ₀ → longer λ (red end) absorbed.
Oxidation state of metal: Higher oxidation state → higher Δ → colour change. [Fe(H₂O)₆]²⁺ (pale green) vs [Fe(H₂O)₆]³⁺ (yellow).
Nature of metal: 4d > 3d complexes → higher Δ due to greater orbital overlap.
Geometry: Square planar complexes are often more intensely coloured than octahedral.

13

🧲

Magnetic Properties of Coordination Compounds

⊕

Diamagnetic

All electrons paired. Weakly repelled by magnetic fields. Magnetic moment μ = 0 BM.

n = 0 unpaired electrons

↑

Paramagnetic

One or more unpaired electrons. Attracted to magnetic fields. μ > 0 BM.

μ = √[n(n+2)] BM

🔴

Ferromagnetic

Strong permanent magnetism due to domain alignment. Metallic iron, cobalt, nickel. Not typical of coordination compounds.

Bulk property

Spin-Only Magnetic Moment Formula

μ = √[n(n+2)] BM (Bohr Magnetons)

This formula is valid when orbital contribution is negligible (spin-only formula). For first-row transition metals, orbital contribution is often quenched.

Magnetic Moment Values

Unpaired electrons (n)	μ (spin-only) BM	Example
0	0	[Co(CN)₆]³⁻, [Ni(CN)₄]²⁻
1	1.73	[Ti(H₂O)₆]³⁺, [Cu(H₂O)₆]²⁺
2	2.83	[V(H₂O)₆]³⁺, [Ni(NH₃)₆]²⁺
3	3.87	[Cr(H₂O)₆]³⁺, [Co(NH₃)₆]³⁺ (but this is d⁶ low spin, n=0!)
4	4.90	[Fe(H₂O)₆]²⁺, [Co(F)₆]³⁻
5	5.92	[Mn(H₂O)₆]²⁺, [Fe(H₂O)₆]³⁺

High Spin vs Low Spin — Magnetic Implications

High spin complexes (weak field ligands, F⁻, Cl⁻, H₂O, OH⁻): more unpaired electrons → higher magnetic moment → strongly paramagnetic.

Low spin complexes (strong field ligands, CN⁻, CO, en, NH₃): electrons pair up → fewer or zero unpaired electrons → weakly paramagnetic or diamagnetic.

Critical Example

[Fe(CN)₆]⁴⁻ vs [Fe(H₂O)₆]²⁺ — both Fe²⁺ (d⁶)

[Fe(CN)₆]⁴⁻: Low spin, t₂g⁶ eₘ⁰, 0 unpaired, μ = 0 BM, Diamagnetic
[Fe(H₂O)₆]²⁺: High spin, t₂g⁴ eₘ², 4 unpaired, μ = 4.90 BM, Paramagnetic

14

📡

Electronic Spectra of Coordination Compounds

Selection Rules for d–d Transitions

Laporte (Parity) Rule: Transitions between orbitals of the same symmetry (g→g or u→u) are forbidden. d–d transitions are Laporte-forbidden in centrosymmetric complexes → low molar absorptivity (ε = 1–100 L mol⁻¹ cm⁻¹). Tetrahedral complexes (no centre of symmetry) are slightly more intense.
Spin Selection Rule: ΔS = 0 — spin must be conserved. High-spin d⁵ (Mn²⁺, Fe³⁺) complexes are extremely weakly coloured because all d–d transitions are spin-forbidden (changing spin pairing is needed).
Transitions can become partially allowed through vibronic coupling (Jahn-Teller distortions) or spin-orbit coupling — hence complexes are visible despite being formally forbidden.

Spectrochemical Series

Ligands are arranged in increasing order of their ability to split d-orbitals (increasing Δ). This is an empirical series:

Why Strong Field Ligands Cause Higher Δ?

CFT cannot explain this, but Ligand Field Theory explains: strong field ligands are also strong π-acceptors (back-donors like CO, CN⁻) or have strong σ-donor ability. Back-bonding from metal d-orbitals into empty ligand π* orbitals lowers the energy of t₂g further → increases Δ.

Weak field ligands (F⁻, Cl⁻) are π-donors: they raise the energy of t₂g → decrease Δ.

Charge Transfer Spectra

In addition to d–d transitions, intense colour can arise from Charge Transfer (CT) transitions:

LMCT (Ligand to Metal CT): Electron transfers from filled ligand orbital to empty metal orbital. Example: MnO₄⁻ (permanganate, d⁰) — deep purple from LMCT, not d–d. CrO₄²⁻ — yellow.
MLCT (Metal to Ligand CT): Electron transfers from filled metal orbital to empty ligand π* orbital. Common in complexes with π-acceptor ligands (CO, bipyridyl). Very intense (ε up to 10⁵).

15

⚙️

Organometallic Compounds & Metal Carbonyls

Definition of Organometallic Compounds

Compounds containing at least one direct metal–carbon bond. The carbon may be part of an alkyl, aryl, alkenyl, alkynyl, or cyclopentadienyl group.

Examples: Zeise's salt [PtCl₃(η²-C₂H₄)]⁻, Ferrocene [Fe(η⁵-Cp)₂], Grignard reagents RMgX, Methyl lithium CH₃Li, WilkinsOn's catalyst [RhCl(PPh₃)₃].

Hapticity (η) — Binding Mode

Hapticity

The number of contiguous atoms in a ligand bonded to the metal is the hapticity, denoted η (eta).

η¹-C₅H₅ — monohapto cyclopentadienyl (σ-bonded)
η²-C₂H₄ — dihapto ethylene (Zeise's salt)
η⁵-C₅H₅ — pentahapto cyclopentadienyl (ferrocene) — most common
η⁶-C₆H₆ — hexahapto benzene in dibenzenechromium [Cr(η⁶-C₆H₆)₂]

Metal Carbonyls — Structure & Bonding

CO as a Ligand — Synergic Bonding

The metal-CO bond involves two simultaneous interactions:

σ-donation: C lone pair donates into empty metal dz² orbital (C → M)
π-back donation: Filled metal t₂g orbitals donate into empty CO π* antibonding orbital (M → C). This weakens C≡O bond, lowers IR stretching frequency.

This synergic/synergistic bonding stabilizes the metal carbonyl and follows the 18-electron rule.

18-Electron Rule

Effective Atomic Number (EAN) / 18-Electron Rule

Stable organometallic compounds (especially carbonyls) obey the 18-electron rule: the total number of electrons at the metal (metal d-electrons + electrons from all ligands) = 18, mimicking the noble gas configuration.

Total electron count = Metal d-electrons + Ligand electrons = 18

Each CO contributes 2 electrons. Each Cp (η⁵-C₅H₅) contributes 5 electrons. Each Cl contributes 1 or 3 electrons depending on counting method (ionic vs covalent).

Verification: [Fe(CO)₅]

Fe(0) = d⁸ = 8e⁻; 5 × CO = 5 × 2 = 10e⁻; Total = 18 ✓

Ferrocene: Fe(II) = d⁶ = 6e⁻; 2 × Cp⁻ = 2 × 6 = 12e⁻ (ionic); Total = 18 ✓

Types of Metal Carbonyls

Type	Examples	Structure	Electron Count
Mononuclear	Ni(CO)₄, Fe(CO)₅, Cr(CO)₆	Tetrahedral, TBP, Octahedral	All 18e
Dinuclear (metal-metal bond)	Mn₂(CO)₁₀, Co₂(CO)₈	D₄d, bridging CO possible	18e each
Polynuclear	Fe₃(CO)₁₂, Os₃(CO)₁₂	Triangular cluster	18e each
With bridging CO	Co₂(CO)₈, Fe₂(CO)₉	μ₂-CO bridges	18e each

16

🧫

Analytical Chemistry — Coordination Compounds

Qualitative Analysis

Ring test for NO₃⁻: FeSO₄ forms [Fe(NO)(H₂O)₅]²⁺ — brown ring complex at H₂SO₄ interface.
Fe³⁺ detection: SCN⁻ → [Fe(SCN)]²⁺ blood-red complex (highly sensitive).
Fe²⁺ detection: K₃[Fe(CN)₆] → turnbull's blue; or K₄[Fe(CN)₆] + Fe³⁺ → prussian blue.
Cu²⁺ detection: NH₃ excess → [Cu(NH₃)₄]²⁺ deep blue.
Ni²⁺ detection: Dimethylglyoxime (DMG) → bright red chelate [Ni(dmgH)₂] in alkaline medium.
Co²⁺ detection: SCN⁻ in acetone → [Co(SCN)₄]²⁻ blue.

Quantitative Analysis (Complexometry)

EDTA titrations: EDTA (ethylenediaminetetraacetic acid, Y⁴⁻) forms very stable 1:1 chelates with nearly all metal ions. Used for water hardness determination, analysis of Ca²⁺, Mg²⁺, Zn²⁺, Fe³⁺, Cu²⁺.
Metal indicators: Eriochrome Black T (EBT) for Mg²⁺/Ca²⁺; Murexide for Ca²⁺; Xylenol orange for Bi³⁺, Pb²⁺.
Gravimetry: Ni²⁺ precipitated as [Ni(dmgH)₂] and weighed.
Spectrophotometry: Coloured complexes measured at λ_max (Beer-Lambert law: A = εcl).

EDTA — The Universal Chelating Agent

EDTA is a hexadentate ligand: 4 carboxylate O atoms + 2 N atoms = 6 donor atoms.

M^n+ + Y^4- → [MY]^(n-4) logK_f = 10–25 (depending on metal)

Selectivity achieved by controlling pH: at low pH, free Y⁴⁻ is reduced; buffers used to maintain appropriate pH for each metal.

logK_f order: Ca²⁺ (10.7) < Cu²⁺ (18.7) < Fe³⁺ (25.1)

Gravimetric & Colourimetric Reagents

DMG (Dimethylglyoxime): Ni²⁺ and Pd²⁺ — red/yellow precipitate.
Cupferron: Fe³⁺ precipitant in acidic solution.
1,10-Phenanthroline (phen): Fe²⁺ → [Fe(phen)₃]²⁺ orange-red; used in spectrophotometric Fe determination.
Thiocyanate: Fe³⁺, Co²⁺, Mo(VI) detection.
Dithizone: Pb²⁺, Hg²⁺, Zn²⁺ — highly sensitive colorimetric reagent.

17

🌐

Major Applications of Coordination Compounds

💊

Medicine

Cisplatin [Pt(NH₃)₂Cl₂] (cis-) is a first-line anticancer drug. EDTA used in lead/heavy metal poisoning (chelation therapy). [Au(CN)₂]⁻ in rheumatoid arthritis treatment. Vitamin B₁₂ (cyanocobalamin) is a cobalt coordination complex. Carboplatin, oxaliplatin as improved platinum anticancer agents.

🔬

Catalysis

Wilkinson's catalyst [RhCl(PPh₃)₃] for hydrogenation of alkenes. Ziegler-Natta catalyst TiCl₄/AlEt₃ for polymerization (polypropylene). Wacker process: [PdCl₄]²⁻ for ethylene → acetaldehyde. Monsanto process: [RhI₂(CO)₂]⁻ for acetic acid synthesis.

🧬

Biochemistry

Haemoglobin: Fe²⁺ porphyrin complex transports O₂. Chlorophyll: Mg²⁺ porphyrin complex for photosynthesis. Cytochrome c: Fe complex in electron transport. Carbonic anhydrase: Zn²⁺ complex catalyses CO₂ hydration. Myoglobin, ferredoxin, nitrogenase (Mo-Fe complex).

⛏️

Metallurgy & Extraction

Cyanide leaching: Gold and silver extracted as [Au(CN)₂]⁻ and [Ag(CN)₂]⁻. Mond process: Nickel purified via Ni(CO)₄ (formed at 50°C, decomposed at 230°C). Zone refining: Volatile carbonyls used for ultra-pure metals.

📸

Photography

Silver halide grains form image. Fixing: Unexposed AgBr dissolved as [Ag(S₂O₃)₂]³⁻ (thiosulfate complex). Platinum printing uses [Pt(C₂O₄)₂]²⁻. Selenium toning uses selenious acid complexes.

🖌️

Pigments & Dyes

Prussian blue Fe₄[Fe(CN)₆]₃ — deep blue pigment used since 1704. Chrome yellow PbCrO₄. Phthalocyanine pigments (Cu complex) — brilliant blue/green dyes used in printing inks, paints.

⚡

Electroplating

Metals deposited from complex solutions for uniform coatings. Gold electroplating from [Au(CN)₂]⁻. Copper plating from [Cu(EDTA)]²⁻ or cyanide baths. Nickel, zinc, silver plating all use coordination compounds for controlled deposition.

🔋

Solar Energy & Electronics

Grätzel cells (DSSC): Ruthenium bipyridyl complex [Ru(bpy)₂(NCS)₂] as photosensitizer. OLEDs: Ir(ppy)₃ as phosphorescent emitter. Quantum dot sensitizers: Cd/Zn chalcogenide complexes. Molecular wires, switches, and sensors in nanotechnology.

18

📋

Practice Quiz — IIT JEE / NEET Level

Coordination Compounds — Self Test

Q 1 / 12 · Score: 0

Quiz Complete! 🎉

19

⚛️

Electronic Configuration & d-Orbital Filling Visualizer

How to Write Electronic Configuration of Complex Metal Ions

Write the ground-state configuration of the neutral metal atom (use [Ar]/[Kr] core).
Remove electrons from the 4s (or ns) orbital first, then from (n-1)d, to form the cation.
Determine the d-electron count (dⁿ).
Based on the field strength of ligands, fill orbitals according to Hund's rule (high spin) or pair electrons (low spin).

Example: Fe (Z=26) → [Ar] 3d⁶ 4s²
Fe²⁺ → remove 2e from 4s → [Ar] 3d⁶ (d⁶)
Fe³⁺ → remove 3e (2 from 4s, 1 from 3d) → [Ar] 3d⁵ (d⁵)

🔬 d-Orbital Filling Interactive Visualizer

Select Metal Ion

Field Strength

Geometry

Electronic Configurations of Common Metal Ions (First Row Transition Metals)

Element (Z)	Atom Config.	M²⁺ (dⁿ)	M³⁺ (dⁿ)	Oct. HS t₂g eₘ	Oct. LS t₂g eₘ	μ (HS) BM
Ti (22)	[Ar]3d²4s²	Ti²⁺ d²	Ti³⁺ d¹	t₂g¹ (d¹)	t₂g¹ (d¹)	1.73
V (23)	[Ar]3d³4s²	V²⁺ d³	V³⁺ d²	t₂g³ (d³)	t₂g³ (d³)	3.87
Cr (24)	[Ar]3d⁵4s¹	Cr²⁺ d⁴	Cr³⁺ d³	t₂g³eₘ¹ (d⁴)	t₂g⁴ (d⁴)	4.90
Mn (25)	[Ar]3d⁵4s²	Mn²⁺ d⁵	Mn³⁺ d⁴	t₂g³eₘ² (d⁵)	t₂g⁵ (d⁵)	5.92
Fe (26)	[Ar]3d⁶4s²	Fe²⁺ d⁶	Fe³⁺ d⁵	t₂g⁴eₘ² (d⁶)	t₂g⁶ (d⁶)	4.90
Co (27)	[Ar]3d⁷4s²	Co²⁺ d⁷	Co³⁺ d⁶	t₂g⁵eₘ² (d⁷)	t₂g⁶eₘ¹ (d⁷)	3.87
Ni (28)	[Ar]3d⁸4s²	Ni²⁺ d⁸	—	t₂g⁶eₘ² (d⁸)	t₂g⁶eₘ² (d⁸)	2.83
Cu (29)	[Ar]3d¹⁰4s¹	Cu²⁺ d⁹	Cu³⁺ d⁸	t₂g⁶eₘ³ (d⁹)	t₂g⁶eₘ³ (d⁹)	1.73
Zn (30)	[Ar]3d¹⁰4s²	Zn²⁺ d¹⁰	—	t₂g⁶eₘ⁴ (d¹⁰)	t₂g⁶eₘ⁴ (d¹⁰)	0 (dia)

20

⚗️

Jahn–Teller Distortion & Trans Effect

Jahn–Teller Effect

The Jahn–Teller theorem states: any non-linear molecular system with a degenerate electronic ground state will spontaneously distort to remove that degeneracy and lower the overall energy.

When does it occur? Whenever the d-orbital occupancy is unequal in a degenerate set:

Strong J-T distortion: Unequal filling of eₘ set (d⁴ HS, d⁷ LS, d⁹) — axial bonds lengthen/shorten dramatically (tetragonal distortion).
Weak J-T distortion: Unequal filling of t₂g set (d¹, d², d⁴ LS, d⁵ LS, d⁶ HS, d⁷ HS) — smaller geometric effect.

Classic example: Cu²⁺ (d⁹) — [Cu(H₂O)₆]²⁺ has 4 short equatorial bonds and 2 elongated axial bonds. This distortion explains the intense colour and anomalous stability of Cu²⁺ complexes.

d⁹ → eₘ³ (dx²-y²)¹(dz²)² → z-elongation (tetragonal distortion)

Trans Effect

In square planar complexes (especially Pt²⁺, Pd²⁺), certain ligands labilise (weaken) the metal–ligand bond trans to themselves, facilitating substitution at the trans position. This is the trans effect (kinetic phenomenon).

Trans effect series (increasing):

F⁻ ≈ H₂O ≈ OH⁻ < NH₃ < Cl⁻ < Br⁻ < SCN⁻ < I⁻ < NO₂⁻ < CO ≈ CN⁻ ≈ C₂H₄

Strong π-acceptor ligands (CO, CN⁻, alkenes) have the strongest trans effect.

Synthesis of cisplatin: Starting from [PtCl₄]²⁻, adding NH₃ step-by-step using the trans effect of Cl⁻ vs NH₃ selectively gives cis product.

Trans influence (thermodynamic): the weakening of the metal–ligand bond trans to a given ligand in the ground state (measured by bond lengths).

Synthesis Strategy Using Trans Effect — Cisplatin vs Transplatin

Route 1 → Cisplatin (cis-[PtCl₂(NH₃)₂]):

Start: [PtCl₄]²⁻ + NH₃ → [PtCl₃(NH₃)]⁻

Cl⁻ trans effect (moderate) directs 2nd NH₃ to position trans to Cl⁻ (i.e., cis to first NH₃)

→ cis-[PtCl₂(NH₃)₂] ✓

Route 2 → Transplatin (trans-[PtCl₂(NH₃)₂]):

Start: [Pt(NH₃)₄]²⁺ + Cl⁻ → [PtCl(NH₃)₃]⁺

NH₃ has weaker trans effect → 2nd Cl⁻ enters trans to the first Cl⁻

→ trans-[PtCl₂(NH₃)₂] ✓

21

🔢

Effective Atomic Number (EAN) Rule — Extended

Sidgwick's EAN Rule

Proposed by Nevil Vincent Sidgwick, the EAN rule states that stable complexes are formed when the total electron count at the metal (including electrons from ligands) equals the atomic number of the next noble gas.

EAN = (Atomic number of metal) − (Charge on metal) + (Electrons from all ligands)

For first-row transition metals, the noble gas configuration = 36 (Kr). For second-row metals, it is 54 (Xe). For third-row metals, it is 86 (Rn).

Complex	Metal	Z (metal)	d-electrons	Ligand electrons	EAN	Noble gas	Stable?
Ni(CO)₄	Ni(0)	28	10 (d¹⁰)	4×CO = 8	28+8=36	Kr(36)	Yes ✓
Fe(CO)₅	Fe(0)	26	8 (d⁸)	5×CO = 10	26+10=36	Kr(36)	Yes ✓
Cr(CO)₆	Cr(0)	24	6 (d⁶)	6×CO = 12	24+12=36	Kr(36)	Yes ✓
[Fe(η⁵-Cp)₂]	Fe(II)	26	6 (d⁶)	2×Cp⁻=12e	26−2+12=36	Kr(36)	Yes ✓
Mn₂(CO)₁₀	Mn(0) each	25	7 (d⁷)	5×CO+Mn = 10+1	25+11=36	Kr(36)	Yes ✓
[Co(NH₃)₆]³⁺	Co(III)	27	6 (d⁶)	6×NH₃ = 12	27−3+12=36	Kr(36)	Yes ✓
[Ni(CN)₄]²⁻	Ni(II)	28	8 (d⁸)	4×CN⁻ = 8	28−2+8+2=36	Kr(36)	Yes ✓

⚠️

Exceptions: The EAN rule is not universally obeyed. Many ionic/Werner-type complexes (e.g., [Co(NH₃)₆]³⁺ does satisfy it, but [CuCl₄]²⁻ = 34e, not 36) are stable despite not following it. The rule is most reliable for organometallics and metal carbonyls.

22

📐

Formula Quick-Reference Sheet

23

⚡

Important Reactions & Qualitative Tests

Prussian Blue Test (Fe²⁺/Fe³⁺)

Fe³⁺ + K₄[Fe(CN)₆] → KFe[Fe(CN)₆] (Prussian blue ↓)

Fe²⁺ + K₃[Fe(CN)₆] → KFe[Fe(CN)₆] (Turnbull's blue ↓)

Both are the same compound KFe[Fe(CN)₆] — a mixed-valence iron coordination polymer. Blue colour arises from electron delocalization (intervalence charge transfer) between Fe²⁺ and Fe³⁺ centres.

DMG Test for Ni²⁺

Ni²⁺ + 2 DMG → [Ni(dmgH)₂] + 2H⁺ (bright red ↓, pH 5–10)

Dimethylglyoxime (DMG) forms a square planar chelate with Ni²⁺. The complex has intramolecular H-bonds stabilizing the structure. This is one of the most selective and sensitive gravimetric/colorimetric tests in analytical chemistry.

Ring Test for Nitrate

NO₃⁻ + Fe²⁺ + H₂SO₄ → [Fe(NO)]²⁺ (brown ring complex at interface)

[Fe(NO)(H₂O)₅]²⁺ is the brown ring compound — a complex of Fe²⁺ with the NO⁺ ligand (nitrosonium). This is a classic qualitative test for nitrate ion.

Deep Blue Copper Complex

Cu²⁺ + 4 NH₃ → [Cu(NH₃)₄]²⁺ (deep blue, tetraamminecopper)

This reaction is used to detect Cu²⁺ and to dissolve Cu(OH)₂. The complex is square planar, d⁹ → dsp² hybridization. The intense blue colour arises from d–d transitions in the copper(II) tetraamine complex.

Mond Process (Ni purification)

Ni + 4CO →(50°C)→ Ni(CO)₄ →(230°C)→ Ni + 4CO

Crude nickel reacts with CO at 50–60°C to form volatile nickel tetracarbonyl Ni(CO)₄ (a liquid, BP 43°C). When heated to 230°C, it decomposes giving 99.99% pure nickel. The CO is recycled.

Cisplatin — Anticancer Mechanism

cis-[PtCl₂(NH₃)₂] + DNA → Pt crosslinks N7 of adjacent guanines

Inside the cell, Cl⁻ ligands are slowly replaced by H₂O (aquation), forming [Pt(NH₃)₂(H₂O)₂]²⁺. This reacts with guanine N7 atoms on the same DNA strand, creating intrastrand crosslinks that bend DNA and block replication → apoptosis of cancer cells.

24

🎯