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CBSE • Class 12 • Chemistry
Coordination Compounds
Ligands, nomenclature, isomerism, bonding in complexes
Chapter 5
Verified Curriculum Topic
What is Coordination Compounds?
Ligands, nomenclature, isomerism, bonding in complexes
Coordination Compounds matters because it links chemical ideas, reactions, and reasoning patterns that recur throughout the syllabus. At Class 12 level, students are often expected to define terms accurately, explain processes clearly, and connect theory to reactions, observations, or applications.
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Summary
The One Thing
Coordination compounds consist of a central metal atom or ion bonded to surrounding ligands through coordinate bonds. Their structures and properties depend chiefly on ligand type, coordination number, oxidation state, geometry, metal–ligand bonding, and d-orbital splitting.
Reactions, Processes and Experiments
| What happens | Equation or process | What you observe | Type |
|---|---|---|---|
| A central metal atom or ion accepts electron pairs from surrounding ligands. | Coordination compounds are generally represented as [M(L)n]X, where M is the central metal, L represents ligands, n is related to coordination number, and X represents counter ions. | A charged or neutral coordination entity is enclosed in square brackets; counter ions lie outside the brackets. | Formation of a coordination entity |
| Ligands donate lone pairs to the metal ion to form coordinate bonds. | A coordinate bond is a covalent bond in which both shared electrons are donated by the ligand to the metal ion. | The ligand and metal are joined by a metal–ligand coordinate bond. | Coordinate bonding |
| The oxidation state of the central metal is calculated from ligand charges and the overall complex charge. | oxidation state of metal + sum of ligand charges = charge on the complex ion | The calculated oxidation state is written in Roman numerals in the compound’s name. | Oxidation-state calculation |
| Three bidentate ethane-1,2-diamine ligands attach to a cobalt ion through six donor atoms. | [Co(en)3]3+ | Three ligands produce coordination number six because each en ligand is bidentate. | Bidentate coordination; chelate formation |
| A monodentate ligand uses one donor atom to bond to the metal. | Examples include NH3, H2O, Cl−, or CN−. | The number of ligands equals the number of donor atoms contributed by the ligands. | Monodentate coordination |
| A bidentate ligand uses two donor atoms to attach to the same metal ion. | Examples include ethane-1,2-diamine, abbreviated en, or oxalate ion, C2O4^2−. | A single ligand contributes two metal–ligand bonds. | Bidentate coordination |
| A polydentate ligand forms several coordinate bonds with one metal ion. | A polydentate ligand has several donor atoms and can form multiple coordinate bonds with one metal ion. | Several donor atoms from one ligand bind to the same metal, often producing chelate rings. | Polydentate coordination |
| A multidentate ligand attaches through two or more donor atoms to form a ring-shaped complex. | A chelate is a ring-shaped complex formed when a multidentate ligand attaches to the same metal ion through two or more donor atoms. | One ligand forms a ring with the metal ion. | Chelation |
| An ambidentate ligand coordinates through one of two possible donor atoms. | NO2− may coordinate through N or O; SCN− may coordinate through S or N. | The same ligand can produce different attachment modes and therefore linkage isomers. | Ambidentate coordination; linkage isomerism |
| The metal and directly attached ligands form the coordination sphere, while external ions balance the charge. | The central metal ion and ligands directly attached to it are shown within square brackets; an ion outside the brackets is a counter ion. | Counter ions are not directly bonded to the metal and may ionise in solution. | Coordination sphere and counter-ion arrangement |
| Werner’s theory distinguishes oxidation-related valency from coordination-related valency. | Primary valency is related to oxidation state and ionisable ions; secondary valency is related to coordination number and non-ionisable ligands. | Primary-valency ions may ionise, whereas secondary-valency ligands remain within the coordination sphere. | Werner’s coordination theory |
| Coordination compounds are named by stating the cation before the anion, ligands before the metal, and the metal oxidation state in Roman numerals. | Naming order: cation before anion; within a complex, ligands first and metal last. | The name identifies ligand number, ligand type, metal, oxidation state, and whether the complex is anionic. | Coordination-compound nomenclature |
| Anionic ligand names are modified in coordination-compound nomenclature. | chloride → chlorido; cyanide → cyanido; hydroxide → hydroxido; oxide → oxido | Anionic ligand names commonly end in -o. | Ligand nomenclature |
| Neutral ligands receive specific coordination-compound names. | H2O → aqua; NH3 → ammine; CO → carbonyl; NO → nitrosyl | Neutral ligands are identified by their specialised names rather than ordinary molecular names. | Ligand nomenclature |
| Numerical prefixes indicate the number of simple ligands. | mono-, di-, tri-, tetra-, penta-, hexa- | The prefix states how many ligands of that type are present. | Nomenclature |
| Bis-, tris-, and tetrakis- are used for complicated or polydentate ligands. | bis-, tris-, tetrakis- | These prefixes distinguish the number of complex or polydentate ligands. | Nomenclature |
| Anionic complexes use a modified metal name. | Examples include ferrate, cuprate, argentate, and cobaltate. | The metal name usually ends in -ate when the complex ion is anionic. | Nomenclature |
| Hexaamminecobalt(III) chloride is named from its complex cation and counter ions. | [Co(NH3)6]Cl3 is hexaamminecobalt(III) chloride | The complex contains six ammine ligands, cobalt in oxidation state +3, and three chloride counter ions. | Naming example |
| Potassium hexacyanidoferrate(II) is named from its potassium counter ions and anionic complex. | K4[Fe(CN)6] is potassium hexacyanidoferrate(II) | The complex contains six cyanido ligands and iron in oxidation state +2; potassium ions lie outside the coordination sphere. | Naming example |
| Diamminedichloridoplatinum(II) identifies two ammine ligands, two chlorido ligands, and platinum in oxidation state +2. | [Pt(NH3)2Cl2] is diamminedichloridoplatinum(II) | The square-bracketed entity contains four donor atoms directly bonded to platinum. | Naming example |
| Coordination number is determined by the total number of donor atoms directly bonded to the metal. | [Co(en)3]3+ has three ligands but coordination number six because en is bidentate. | Coordination number may be greater than the number of ligand molecules. | Coordination-number determination |
| Two-coordinate complexes commonly have linear geometry. | Coordination number 2: linear | Two donor atoms lie opposite one another around the metal. | Geometry |
| Four-coordinate complexes may be tetrahedral or square planar. | Coordination number 4: tetrahedral or square planar | Donor atoms either occupy a tetrahedral arrangement or lie in one plane. | Geometry |
| Six-coordinate complexes commonly have octahedral geometry. | Coordination number 6: octahedral | Six donor atoms surround the metal ion. | Geometry |
| In an octahedral field, the five d-orbitals split into two energy levels. | The lower-energy set is t2g and the higher-energy set is eg, with splitting energy Δo. | Electrons occupy separated d-orbital energy levels; the arrangement helps determine spin state, colour, and magnetism. | Crystal field splitting |
| In a tetrahedral field, the five d-orbitals split into two energy levels. | The lower-energy set is e and the higher-energy set is t2; Δt = 4/9 Δo. | Tetrahedral splitting is smaller than octahedral splitting. | Crystal field splitting |
| Strong-field ligands produce relatively large d-orbital splitting and may cause electron pairing. | Strong-field examples include CN− and CO. | Electrons may pair, producing a low-spin complex when pairing is energetically favourable. | Strong-field ligand effect |
| Weak-field ligands produce relatively small d-orbital splitting. | Weak-field examples include F−, Cl−, and H2O. | Electrons generally remain unpaired when possible, producing high-spin complexes. | Weak-field ligand effect |
| Ligands are arranged in order of increasing field strength. | I− < Br− < Cl− < F− < OH− < H2O < NH3 < en < CN− < CO | Ligands farther to the right generally cause larger crystal field splitting. | Spectrochemical series |
| Hybridisation descriptions relate metal orbitals to approximate complex geometry. | Octahedral: d2sp3 or sp3d2; tetrahedral: sp3; square planar: dsp2. | The hybridisation description indicates the predicted geometry and, in some cases, inner- or outer-orbital character. | Valence bond theory |
| Inner-orbital complexes use inner d-orbitals in their hybridisation description. | d2sp3 hybridisation is often associated with inner-orbital and low-spin behaviour. | Pairing may occur in inner d-orbitals, leaving fewer unpaired electrons. | Inner-orbital complex |
| Outer-orbital complexes use outer d-orbitals in their hybridisation description. | sp3d2 hybridisation is often associated with outer-orbital and high-spin behaviour. | More unpaired electrons may remain because pairing is not favoured. | Outer-orbital complex |
| The number of unpaired electrons determines magnetic behaviour. | μ = √[n(n + 2)] BM, where n is the number of unpaired electrons and BM means Bohr magneton. | Complexes with unpaired electrons are paramagnetic; a complex with no unpaired electrons is diamagnetic. | Magnetic behaviour |
| Ligands may occupy different relative positions around the metal. | Geometrical isomerism includes cis-trans forms and, in some octahedral complexes, fac-mer forms. | Cis and trans forms differ in ligand positions; fac and mer forms occur in suitable octahedral complexes. | Geometrical isomerism |
| Octahedral complexes with suitable ligand arrangements may form non-superimposable mirror images. | Optical isomerism is commonly observed in [Co(en)3]3+ and in certain cis complexes. | The two forms are non-superimposable mirror images called enantiomers. | Optical isomerism |
| An ambidentate ligand can attach through different donor atoms. | NO2− may coordinate through nitrogen as nitro or through oxygen as nitrito. | Different donor atoms produce linkage isomers. | Linkage isomerism |
| A ligand inside the coordination sphere can exchange with a counter ion outside it. | Ionisation isomers have the same overall composition but produce different ions in solution because a coordinated ligand and counter ion exchange positions. | Different ions are released in solution despite the same overall composition. | Ionisation isomerism |
| Solvent molecules may occupy different positions relative to the coordination sphere. | Solvate or hydrate isomerism is caused by different numbers of solvent molecules inside and outside the coordination sphere. | Isomers differ in the number of solvent molecules coordinated to the metal versus present outside the brackets. | Solvate or hydrate isomerism |
| Ligands may exchange between cationic and anionic complex ions in the same salt. | Coordination isomerism occurs in salts containing both cationic and anionic complexes when ligands exchange between the two metal coordination spheres. | The cationic and anionic complexes have different ligand distributions. | Coordination isomerism |
| Multidentate ligands generally stabilise complexes more effectively than equivalent monodentate ligands. | The chelate effect generally increases the stability of complexes containing multidentate ligands compared with similar complexes containing equivalent monodentate ligands. | Chelated complexes are generally more stable because several coordinate bonds hold the ligand and metal together. | Chelate effect |
| Coordination compounds perform biological, analytical, metallurgical, medical, electroplating, and extraction functions. | Examples include haemoglobin, which contains an iron complex, and chlorophyll, which contains a magnesium complex. | Their significance arises from the properties of their metal centres and ligand arrangements. | Applications |
Key Terms
- Coordination entity: A charged or neutral species containing a central metal atom or ion bonded to surrounding ligands, written inside square brackets.
- Central metal atom or ion: The metal species that accepts electron pairs from ligands; it is commonly a transition-metal ion.
- Ligand: An ion or molecule that donates one or more lone pairs of electrons to the central metal ion to form coordinate bonds.
- Monodentate ligand: A ligand that uses one donor atom to bond with the metal, such as NH3, H2O, Cl−, or CN−.
- Bidentate ligand: A ligand that uses two donor atoms to attach to the same metal ion, such as ethane-1,2-diamine, abbreviated en, or oxalate ion, C2O4^2−.
- Polydentate ligand: A ligand that has several donor atoms and can form multiple coordinate bonds with one metal ion.
- Ambidentate ligand: A ligand that can coordinate through either of two different donor atoms, such as NO2− through N or O and SCN− through S or N.
- Chelate: A ring-shaped complex formed when a multidentate ligand attaches to the same metal ion through two or more donor atoms.
- Coordination number: The total number of donor atoms directly bonded to the central metal ion, not simply the number of ligands.
- Coordination sphere: The central metal ion and ligands directly attached to it, shown within square brackets.
- Counter ion: An ion outside the coordination sphere that balances the charge of the complex ion.
- Oxidation state: The formal charge assigned to the central metal ion after considering the charges of all ligands and the overall charge of the complex.
- Werner's coordination theory: The theory states that metals show primary valency, related to oxidation state and ionisable ions, and secondary valency, related to coordination number and non-ionisable ligands.
- Coordinate bond: A covalent bond in which both shared electrons are donated by the ligand to the metal ion.
- Valence bond theory: A model that explains complex geometry and magnetic behaviour using hybridisation of metal orbitals.
- Crystal field theory: A model that treats ligands as electric fields and explains the splitting of metal d-orbitals in different geometries.
- Octahedral complex: A complex in which six donor atoms surround the metal ion; common hybridisation descriptions include d2sp3 or sp3d2.
- Tetrahedral complex: A complex in which four donor atoms surround the metal ion, commonly described by sp3 hybridisation.
- Square-planar complex: A complex in which four donor atoms lie in one plane around the metal ion, commonly described by dsp2 hybridisation.
- Inner-orbital complex: A complex described using inner d-orbitals in hybridisation, often associated with d2sp3 hybridisation and low-spin behaviour.
- Outer-orbital complex: A complex described using outer d-orbitals in hybridisation, often associated with sp3d2 hybridisation and high-spin behaviour.
- Geometrical isomerism: Isomerism caused by different relative positions of ligands, such as cis and trans forms in square-planar or octahedral complexes.
- Optical isomerism: Isomerism in which two complexes are non-superimposable mirror images called enantiomers.
- Linkage isomerism: Isomerism caused by coordination through different donor atoms of an ambidentate ligand.
- Ionisation isomerism: Isomerism caused by exchange of a ligand inside the coordination sphere with a counter ion outside it.
- Solvate or hydrate isomerism: Isomerism caused by different numbers of solvent molecules inside and outside the coordination sphere.
- Coordination isomerism: Isomerism found in salts containing both cationic and anionic complexes when ligands exchange between the two metal coordination spheres.
- Magnetic moment: The magnetic behaviour of a complex is related to the number of unpaired electrons and can be estimated using μ = √[n(n + 2)] BM.
- Crystal field splitting energy: The energy difference between sets of d-orbitals formed when ligands approach a metal ion; it is represented by Δo for octahedral and Δt for tetrahedral complexes.
- Spectrochemical series: The arrangement of ligands in increasing order of field strength: I− < Br− < Cl− < F− < OH− < H2O < NH3 < en < CN− < CO.
Easily Confused
- Coordination number vs number of ligands: Coordination number counts donor atoms directly bonded to the metal, whereas the number of ligands counts ligand molecules or ions; [Co(en)3]3+ has three ligands but coordination number six.
- Coordination sphere vs counter ion: Species inside square brackets are directly bonded to the metal; ions outside the brackets balance charge and are not directly coordinated.
- Monodentate vs bidentate ligand: A monodentate ligand uses one donor atom, whereas a bidentate ligand uses two donor atoms.
- Bidentate vs polydentate ligand: A bidentate ligand uses two donor atoms; a polydentate ligand uses several donor atoms.
- Ambidentate vs polydentate ligand: An ambidentate ligand can bind through either of two alternative donor atoms, whereas a polydentate ligand binds through several donor atoms at the same time.
- Cis-trans vs fac-mer isomerism: Cis-trans isomerism distinguishes adjacent and opposite ligand positions, while fac-mer isomerism applies to certain octahedral arrangements.
- Linkage vs ionisation isomerism: Linkage isomerism changes the donor atom used by an ambidentate ligand; ionisation isomerism exchanges a coordinated ligand with an external counter ion.
- Ionisation vs solvate or hydrate isomerism: Ionisation isomerism changes the positions of a ligand and counter ion; solvate or hydrate isomerism changes the positions of solvent molecules.
- Inner-orbital vs outer-orbital complex: Inner-orbital complexes commonly use d2sp3 hybridisation and are associated with low spin; outer-orbital complexes commonly use sp3d2 hybridisation and are associated with high spin.
- Strong-field vs weak-field ligand: Strong-field ligands such as CN− and CO produce larger splitting and may cause pairing; weak-field ligands such as F−, Cl−, and H2O produce smaller splitting and generally favour unpaired electrons.
- Paramagnetic vs diamagnetic complex: Paramagnetic complexes contain unpaired electrons; diamagnetic complexes contain no unpaired electrons.
- Δo vs Δt: Δo is octahedral splitting energy, whereas Δt is tetrahedral splitting energy and is approximately 4/9 Δo.
- Valence bond theory vs crystal field theory: Valence bond theory explains geometry and magnetism through hybridisation; crystal field theory explains d-orbital splitting, colour, spin state, and magnetic behaviour.
What Gets Asked
- Calculate the oxidation state of the metal using oxidation state of metal + sum of ligand charges = charge on the complex ion. Marks are lost by ignoring ligand charges or the overall charge of the complex.
- Determine the coordination number and geometry. The key slip is counting ligands rather than donor atoms, especially for bidentate ligands such as en.
- Name coordination compounds such as [Co(NH3)6]Cl3, K4[Fe(CN)6], and [Pt(NH3)2Cl2]. Marks are lost by giving ligands in the wrong order, omitting ligand prefixes, using incorrect anionic ligand names, or failing to state the metal oxidation state.
- Identify the type of isomerism from a structural description. The distinction must be made between cis-trans or fac-mer geometrical isomerism, optical isomerism, linkage isomerism, ionisation isomerism, solvate or hydrate isomerism, and coordination isomerism.
- Predict magnetic behaviour and calculate magnetic moment using μ = √[n(n + 2)] BM. Marks are lost by miscounting unpaired electrons or confusing paramagnetic and diamagnetic behaviour.
- Compare strong-field and weak-field ligands using the spectrochemical series. The specific error is reversing the significance of ligand strength, d-orbital splitting, electron pairing, and high-spin or low-spin behaviour.
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What is a ligand in a coordination compound?
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What is Coordination Compounds in CBSE Class 12 Chemistry?
Ligands, nomenclature, isomerism, bonding in complexes
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