NCERT Solutions, Question Answer and Mind Map for Class 12 Chemistry Chapter 9, “Coordination Compounds,” is a study material package designed to help students understand the properties and characteristics of coordination compounds, their structure, and the nomenclature and isomerism associated with them.
NCERT Solutions provide detailed explanations and answers to the questions presented in the chapter. The solutions cover all the topics in the chapter, including the Werner’s theory of coordination compounds, the structure of coordination compounds, nomenclature of coordination compounds, and isomerism in coordination compounds. They also provide tips on how to answer different types of questions, including short answer, long answer, and multiple-choice questions.
The question-answer section of the chapter covers a wide range of topics, from the classification of ligands and the determination of oxidation states to the structure of coordination compounds and the different types of isomerism. It also includes questions on the stability and reactivity of coordination compounds.
The mind map provides a visual representation of the key topics covered in the chapter, allowing students to understand the connections between different concepts and ideas. The mind map covers the different types of ligands, the structure and nomenclature of coordination compounds, and the various types of isomerism.
NCERT Solution / Notes Class 12 Chemistry Chapter 9 Coordination Compounds with Mind Map PDF Download
Coordination compounds are a special class of compounds in which the central metal atom is surrounded by ions or molecules beyond their normal valency. These are also referred to as complex compounds. These compounds are called Coordination compounds.
Many biologically important compounds are coordination compounds in which complicated organic species are bound to metal ions. Common examples are haemoglobin which is a coordination compound of iron, chlorophyll which is a coordination compound of magnesium, vitamin B12 which is a coordination compound of cobalt etc.
The coordination compounds contain a central metal atom or ion surrounded by a number of oppositely charged ions or neutral molecules more than its normal valency.
For example: When aqueous ammonia is added to green solution of nickel chloride, NiCl2, the colour changes to purple. The Ni2+ ions almost disappear from the solution. The solution on evaporation, yields purple crystals corresponding to the formula [Ni(NH3)6]Cl2. Such a compound is called coordination (or complex) compound. The properties of the complex compound are completely different from those of Ni2+ ions or ammonia molecules.
When the compound is dissolved in water, it ionises to give a new species (Ni(NH3)6]2+. Such an ion is called complex ion.
[Ni(NH3)6]Cl2 → [Ni(NH3)6]2+ + 2Cl–
At this stage, it may be noted that the species in the square brackets does not ionise. It remains as single entity. It is known as complex entity.
Double salts are addition or molecular compounds which are formed by two apparently saturated compounds but they lose their identity when dissolved in water.
The common double salts are:
Mohr’s salt : FeSO4. (NH4)2SO4.6H2O
Potash alum : K2SO4.Al2(SO4)3.24H2O
Carnallite : KCl.MgCl2.6H2O
For example: Mohr’s salt dissolves in water and gives the characteristic properties of Fe2+, NH4+, and SO42- ions. Thus, double salts are stable in solid state but break up into constituents when dissolved in water.
FeSO4.(NH4)2SO4 → Fe2+ + (aq) + 2NH4+ (aq) + 2SO42-(aq)
On the other hand, the coordination compounds retain their identities in the solid state as well as when dissolved in water or any other solvent. Their properties are completely different from the constituents (metal and ions or molecules).
For example, [Ni(NH3)6]Cl2 does not show the properties of NiCl2 or ammonia. Similarly, complex ion such as [Fe(CN)6]4- of K4[Fe(CN)6] does not dissociate into Fe2+ and CN¯ ions.
Werner’s Coordination Theory
Alfred Werner a Swiss chemist, in 1892 prepared a large number of coordination compounds and studied their physical, chemical and isomeric behaviour by simple experimental techniques. He isolated cobalt compounds from the reaction of cobalt chloride and ammonia.
The earlier studies of cobalt complexes were precipitation reactions, conductance measurements and isomeric behaviour.
(1) Precipitation Studies
The number of ions furnished by a complex in a solution can be determined by precipitation reactions.
(a) The number of Cl‾ ions in a solution of various amines were determined by the treatment with silver nitrate solution. From the amount of white precipitate of AgCl formed per mole of the compound, the number of Cl‾ ions can be calculated.
(b) When the compound CoCl3.6NH3 is treated with excess of AgNO3 , 3 mol of AgCl are obtained from 1 mol of the compound i.e. all the three Cl¯ ions are precipitated.
(c) When the compound CoCl3.5NH3 is treated with excess of AgNO3, 2 mol of AgCl are obtained i.e., only two Cl¯ ions are precipitated. This means that the compound CoCl3.5NH3 has three ionizable chloride ions whereas in the compound CoCl3.5NH3 only two chlorine atoms are ionizable as Cl¯ ions.
CoCl3.6NH3 → 3 AgCl ( corresponding to 3 Cl¯ ions)
CoCl3.5NH3 → 2 AgCl ( corresponding to 2 Cl¯ ions)
Similarly, the number of chloride ions precipitated in the case of the compounds CoCl3.4NH3 and CoCl3.3NH3 have been found to be 1 and none.
(2) Conductance measurements
The measurement of molar conductances (^m) of solutions of coordination compounds helps to estimate the number of ions furnished by the compound in solution.
By comparing the molar conductance of the compounds with those of some known electrolytes, Werner was able to predict the number of ions present in the solution.
For example: The complex CoCl3.6NH3 behaved as 1:3 electrolyte, CoCl3.5NH3 as 1:2 electrolyte, CoCl3.4NH3 as 1:1 electrolyte.
(3) Isomers of compounds
Werner attempted to assign structures of different coordination compounds by comparison of the number of known isomers and the number of theoretically possible structures.
Postulates of Werner’s Coordination Theory
(1) In co-ordination compounds, metal atoms exhibit two types of valencies namely, the primary valency and the secondary valency. The primary valency is ionizable whereas the secondary valency is non ionizable. The primary valency corresponds to oxidation state and the secondary valency corresponds to coordination number.
(2) Every metal atom has a fixed number of secondary valencies i.e., it he fixed coordination number.
(3) The metal atom tends to satisfy both its primary as well as secondary valencies. Primary valencies are satisfied by negative ions whereas secondary valencies are satisfied either by negative ions or by neutral molecules. In certain cases, a negative ion may satisfy both types of valencies.
(4) The secondary valencies are always directed towards the fixed position in space and this leads to definite geometry of the coordination compound. Secondary valencies have characteristic spatial arrangements corresponding to different coordination numbers. Spatial arrangements are called coordination polyhedra.
For example: If a metal ion has six secondary valencies, these are arranged octahedrally around the central metal ion. If the metal ion has four secondary valencies, these are arranged in either tetrahedral or square planar arrangement around the central metal ion. The secondary valencies, thus, determine the stereochemistry of the complex.
Thus, a metal atom exhibits primary valencies in the formation of its salts (e.g., CoCl3 , AgNO3) while the metal atom exhibits its secondary valencies in the formation of its complex ions
e.g.: [Co(NH3)6]3+ , [Ag(NH3)2]+
Structures of Coordination Compounds on the Basis of Werner’s Theory
(1) CoCl3.6NH3 : Cobalt has primary valency (oxidation state) of three and secondary valency (coordination number) six. Secondary valencies are represented by thick lines (—) and primary valencies are shown by dotted lines (….). In the complex, all the 6 secondary valencies are occupied by six NH3 molecules. The Cl¯ ions are bonded to Co by three primary valencies. These chloride ions are ionisable and therefore can be precipitated on the addition of silver nitrate. The central metal ion and the neutral molecules or ions (ligands) satisfying secondary valencies are written in a square bracket while writing the formula of the complex compound.
Thus, the coordination compound may be formulated as [Co(NH3)6]Cl3 .
The primary valencies are ionizable and therefore, all the chloride ions would get precipitated on the addition of silver nitrate.
The species within the square brackets are also called coordination entities (or complexes). The ions outside the square brackets are called counter ions. Thus, in the coordination compound [Co(NH3)6Cl]3, [Co(NH3)6]3+ represents coordination entity and 3Cl¯ ions represent counter ions
The ionisation of the coordination compound is written as:
[Co(NH3)6Cl]3 ⇔ [Co(NH3)6]3+ + 3Cl¯
(2) CoCl3.5NH3 : In this compound, the coordination number of cobalt is 6 but now five positions are occupied by NH3 molecules and the sixth position by one of the chloride ions.This chloride ion has dual character as it satisfies secondary as well as a primary valency as indicated by a full line as well as a dotted line.The two Cl¯ ions satisfy the remaining two primary valencies of cobalt. This satisfies 6 secondary and 3 primary valencies of cobalt. However, on ionisation, only two Cl¯ ions will be precipitated because one Cl- ion which also satisfied secondary valency, will not be precipitated.
Thus, the coordination compound may be formulated as [CoCl(NH3)5]Cl2 which has [CoCl(NH3)5]2+ complex entity and 2Cl‾ ions as counter ions. The ionisation of the coordination compound may be written as:
[CoCl(NH3)5]Cl2 ⇔ [CoCl(NH3)5]2+ + 2Cl‾
(3) CoCl3.4NH3 : In the compound CoCl3.4NH3 , two chloride ions exhibit dual character of satisfying both primary and secondary valencies. It will give precipitate with silver nitrate corresponding to only one Cl‾ ion and the number of ions in this case is 2.
It may be formulated as
[CoCl2(NH3)4]Cl⇔ [CoCl2(NH3)4]+ + Cl‾
(4) CoCl3.NH3 : In the compound CoCl3.3NH3 ,three chloride ions satisfy primary and secondary valencies. All the chloride ions are non-ionisable and will not be precipitated by the addition of AgNO3. Therefore, the coordination compound behaves as neutral non- conducting molecule.
It may be formulated as [CoCl3(NH3)3] and does not ionise.
[CoCl3(NH3)3]⇔ Does not ionize
Differences between coordination compound and double bond
|Coordination compound||Double salt|
|A coordination compound contains a central metal atom or ion surrounded by several oppositely charged ions or neutral molecules. These ions ormolecules re-bonded to the metal atom or ion by a coordinate bond.||When two salts in stoichiometric ratio are crystallised together from their saturated solution, they are called double salts.|
|Example: K4[Fe(CN)6]||Example: FeSO4.(NH4)2SO4.6H2O(Mohr’s salt)|
|They do not dissociate into simple ions when dissolved in water.||They dissociate into simple ions when dissolved in water.|
- Coordination entity: A coordination entity constitutes a central metal atom or ion bonded to a fixed number of ions or molecules. Example: In K4[Fe(CN)6], [Fe(CN)6]4− represents a coordination entity.
- Central atom or ion: In a coordination entity, the atom/ion to which a fixed number of ions/groups are bound in a definite geometrical arrangement is called the central atom or ion. Example: In K4[Fe(CN)6], Fe²+ is the central metal ion.
- Ligands: A molecule, ion or group which is bonded to the metal atom or ion in a complex or coordination compound by a coordinate bond is called a ligand. It may be neutral, positively or negatively charged. Examples: H2O, CN−, NO+ etc.
- Donor atom: An atom of the ligand attached directly to the metal is called the donor atom. Example: In the complex K4[Fe(CN)6], carbon is a donor atom.
- Coordination number: The coordination number (CN) of a metal ion in a complex can be defined as the number of ligand donor atoms to which the metal is directly bonded. Example: In the complex K4[Fe(CN)6], the coordination number of Fe is 6.
- Coordination sphere: The central atom/ion and the ligands attached to it are enclosed in square bracket and is collectively termed the coordination sphere. Example: In the complex K4[Fe(CN)6], [Fe(CN)6]4− is the coordination sphere.
- Counter ions: The ions present outside the coordination sphere are called counter ions. Example: In the complex K4[Fe(CN)6], K+ is the counter ion.
- Coordination polyhedron: The spatial arrangement of the ligand atoms which are directly attached to the central atom/ion defines a coordination polyhedron about the central atom. The most common coordination polyhedra are octahedral, square planar and tetrahedral. Examples: [PtCl4]2− is square planar, Ni(CO)4 is tetrahedral and [Cu(NH3)6]3+ is octahedral.
- Charge on the complex ion: The charge on the complex ion is equal to the algebraic sum of the charges on all the ligands coordinated to the central metal ion.
- Denticity: The number of ligating (linking) atoms present in a ligand is called denticity.
- Unidentate ligands: The ligands whose only donor atom is bonded to a metal atom are called unidentate ligands. Examples: H2O, NH3, CO, CN−
- Didentate ligands: The ligands which contain two donor atoms or ions through which they are bonded to the metal ion. Example: Ethylene diamine (H2NCH2CH2NH2) has two
- Polydentate ligand: When several donor atoms are present in a single ligand, the ligand is called a polydentate ligand. Example: In N(CH2CH2NH2)3, the ligand is said to be polydentate. Ethylenediaminetetraacetate ion (EDTA4–) is an important hexadentate ligand. It can bind through two nitrogen and four oxygen atoms to a central metal ion.
- Chelate: An inorganic metal complex in which there is a close ring of atoms caused by attachment of a ligand to a metal atom at two points. An example is the complex ion formed between ethylene diamine and cupric ion [Cu(NH2CH2NH2)2]2+.
- Ambidentate ligands: Ligands which can ligate (link) through two different atoms present in it are called ambidentate ligands. Examples: NO − and SCN−. NO − can link through N as well as O, while SCN− can link through S as well as N.
- Werner’s coordination theory: Werner was able to explain the nature of bonding in complexes. The postulates of Werner’s theory are
- Metal shows two kinds of valencies—primary valence and secondary valence.
|Primary valence||Secondary valence|
|This valence is normally ionisable.||This valence is non-ionisable.|
|It is equal to the positive charge on the central metal atom.||The secondary valency equals to the number of ligand atoms coordinated to the metal. It is also called the coordination number of the metal.|
|These valencies are satisfied by negatively charged ions.||It is commonly satisfied by neutral and negatively charged, sometimes by positively charged ligands.|
|Example: In CrCl3, the primary valency is three. It is equal to theoxidation state of the central metal ion.|
- The ions/groups bound by secondary linkages to the metal have characteristic spatial arrangements corresponding to different coordination numbers.
- The most common geometrical shapes in coordination compounds are octahedral, square planar and tetrahedral.
- Oxidation number of the central atom: The oxidation number of the central atom in a complex is defined as the charge it would carry if all the ligands are removed along with the electron pairs which are shared with the central atom.
- Homoleptic complexes: Those complexes in which metal or ion is coordinately bonded to only one kind of donor atom. Example: [Co(NH3)6]3+
- Heteroleptic complexes: Those complexes in which metal or ion is coordinately bonded to more than one kind of donor atom. Example: [CoCl2(NH3)4]+, [Co(NH3)5Br]2+
- Isomers: Two or more compounds which have the same chemical formula but different arrangement of atoms are called isomers.
- Types of isomerism
- Structural isomerism
- Linkage isomerism
- Solvate isomerism or hydrate isomerism
- Ionisation isomerism
- Coordination isomerism
- Geometrical isomerism
- Optical isomerism
- Structural isomerism
- Structural isomerism: This type of isomerism arises due to the difference in structures of coordination compounds. Structural isomerism, or constitutional isomerism, is a form of isomerism in which molecules with the same molecular formula have atoms bonded together in different orders.
- Ionisation isomerism: This form of isomerism arises when the counter ion in a complex salt is itself a potential ligand and can displace a ligand which can then become the counter ion.
Examples: [Co(NH3)5Br] SO4 and [Co(NH3)5 SO4] Br
- Solvate isomerism: It is isomerism in which the solvent is involved as the ligand. If the solvent is water, then it is called hydrate isomerism. Example: [Cr(H2O)6]Cl3 and [CrCl2(H2O)4] Cl2.2H2O
- Linkage isomerism: Linkage isomerism arises in a coordination compound containing an ambidentate ligand. In the isomerism, a ligand can form linkage with metal through different atoms.
Examples: [Co(NH3)5ONO]Cl2 and [Co(NH3)5NO2]Cl2
- Coordination isomerism: This type of isomerism arises from the interchange of ligands between cationic and anionic entities of different metal ions present in a complex. Examples: [Co(NH3)6][Cr(C2O4)3] and [Cr(NH3)6][Co(C2O4)3]
- Stereoisomerism: This type of isomerism arises because of different spatial arrangement.
- Geometrical isomerism: It arises in heteroleptic complexes due to different possible geometrical arrangements of ligands.
- Optical isomerism: Optical isomers are those isomers which are non-superimposable mirror images.
Isomerism in Coordination Compounds
Isomerism: Two or more compounds having the same molecular formula but different arrangement of atoms are called isomers and the phenomenon is called isomerism. Because of different arrangement of atoms, isomers differ in one or more physical or chemical properties.
Isomers can be broadly classified into two major categories :
(A) Structural isomers
1. Ionisation isomerism
2. Hydrate isomerism
3. Coordination isomerism
4. Linkage isomerism
1. Geometrical isomerism
2. Optical isomerism
(A) Structural isomerism
The isomers which have same molecular formula but different structural arrangement of atoms or groups of atoms around the central metal ion are called structural isomers.
(1) Ionisation isomerism
The compounds which have same molecular formula but give different ions in solution are called ionisation isomers. In this type of isomerism, the difference arises from the interchange of groups within or outside the coordination entity. This type of isomerism occurs when the counter ion in a coordination compound is itself a potential ligand.
For example: There are two isomers of the compound of the formula Co(NH3)5BrSO4
(a) One of these is red-violet and forms a precipitate with BaCl2 indicating that sulphate ion is outside the coordination entity.
(b) The second one is red and does not form precipitate with BaCl2 but forms a precipitate of AgBr with silver nitrate indicating that bromide ion is outside the coordination entity.
The structures of the two compounds and their mode of ionisation are :
Other compounds showing this type of isomerism are:
(i) [CoCl2(NH3)4]NO2 and [CoCl(NO2)(NH3)4]Cl
(ii) [Co(NO3)(NH3)]SO4 and [Co(SO4)(NH3)5]NO3
(iii) [PtCl2(NH3)4]Br2 and [PtBr2(NH3)4]Cl2
(iv) [CoCl(NO2)(NH3)4]Cl and [CoCl2(NH3)4]NO2
(2) Solvate or Hydrate isomerism
The compounds which have the same molecular formula but differ by whether or not a solvent molecule is directly bonded to the metal ion or merely present as free solvent molecules in the crystal lattice are called solvate isomers.
It is also known as hydrate isomerism where water is involved as a solvent.Thus, hydrate isomers differ in the number of water molecules present as ligands or as molecules of hydration.
In type of isomerism water molecules may occur inside and outside the coordination sphere as a coordinated group or a water of hydration.
For example, there are three isomers having the molecular formula CrCl3⋅6H2O.
These are :
(CrCl3(H2O)3], [CrCl(H2O)5]Cl2⋅H2O and [CrCl2(H2O)4]Cl⋅2H2O
(i) [Cr(H2O)6]Cl3 : It does not lose water when treated with conc. H2SO4 and three chloride ions are precipitated with AgNO3.
It loses one water molecule when treated with conc. H2SO4 and 2Cl¯ ions are precipitated with AgNO3
(iii) (CrCl2(H2O)4]Cl.2H2O : It loses two water molecules on treatment with conc. H2SO4 dark green and one Cl¯ ion is precipitated with AgNO3.
Similarly, the following two isomers are hydrate isomers :
[CoCl(en)2(H2O)]Cl2 and [CoCl2(en)2]Cl.H2O
[CoCl(H2O)(NH3)4]Cl2 and [CoCl2(NH3)4]Cl.H2O
[CrCl2(C2H5N)2(H2O)2]Cl and [CrCl3(C2H5N5)2H2O].H2O
(3) Coordination isomerism
The type of isomerism occurs in compounds containing both cationic and anionic entities and the isomers differ in the distribution of ligands in the coordination entity of cationic and anionic parts. This type of isomerism arises from the interchange of ligands between cationic and anionic entities of different metal ions present in the complex.
The examples are:
(i) [Co(NH3)6][Cr(CN)6] and [Cr(NH3)6] [Co(CN)6]
(ii) [Cu(NH3)4[PtCl4] and [Pt(NH3)4][CuCl4]
This type of isomerism is also shown by compounds in which the metal ion is the same in both cationic and anionic complexes.
For example :
(i) [Cr(NH3)6][Cr(CN)6] and [Cr(CN)2(NH3)4][Cr(CN)4(NH3)2]
(ii) [(Pt(NH3)4] [PtCl4] and [PtCl(NH3)3][PtCl3(NH3)]
(4) Linkage isomerism
The compounds which have the same molecular formula but differ in the mode of attachment of a ligand to the metal atom or ion are called linkage isomers.
For example: In NO2¯ ion, the nitrogen atom as well as the oxygen atom can donate their lone pairs. This gives rise to isomerism.
If nitrogen donates its lone pair, one particular compound will be formed.
If oxygen donates its lone pair, a different compound (although having the same molecular formula) is obtained. If the bonding is through N, the ligand is named as nitrito-N(or nitro) and if it is through O, it is named as nitrito-O (or nitrito).
NO2¯ nitrito-N (or nitro)
ONO¯ nitrito-O (or nitrito)
For example: Jorgensen discovered such behaviour in the complex [(Co(NH3)5(NO2)]Cl. He prepared two different pentaamminecobalt(II) chloride each containing the NO2 group in the complex ion. These are:
The unidentate ligands which can bind to the central atom through two donor atoms are also called ambidentate ligands.
Other examples of ligands are:
CN Cyano (through C)
NC Isocyano (through N)
SCN Thiocyanato (through S)
NCS Isothiocyanato (through N)
Stereoisomers are those isomers which have the same position of atoms or groups but they differ in the spatial arrangements around the central atom. Two types of isomerism viz., geometrical isomerism and optical isomerism.
(1) Geometrical isomerism
Geometrical isomerism arises in heteroleptic complexes due to ligands occupying different positions around the central ion. The ligands occupy positions either adjacent to one another or opposite to one another. These are referred to as cis- form (ligands occupy adjacent positions) and trans- form (ligands occupy opposite positions). This type of isomerism is, therefore, also referred to as cis-trans isomerism.
(a) Geometrical isomerism in complexes of coordination number 4
The complexes having coordination number 4 adopt tetrahedral or square planar geometry. The geometrical isomerism is not possible in tetrahedral complexes. This is because in tetrahedral geometry all the positions adjacent to one another in these complexes.
However, square planar complexes show geometrical isomerism.
- Square planar complexes of the type MA2X2 , MA2XY, MABX2, MABXY can exist as geometrical isomers (Here A and B are neutral ligands such as H2O, NH3, CO, NO, C5H5N whereas X and Y are anionic ligands such as Cl¯, NO2‾, CN¯, SCN¯ etc.)
(ii) [PtCl(C5H5N)2 (NH3)] exists in cis and trans form as:
(iii) Square planar complexes of the type MABCD show three isomers. The structures of these isomers can be written by fixing the position of one ligand (say A) and placing the other ligands B, C and D trans to it.
The complex [Pt(NO2)(py) (NH2OH)(NH3)]+ exists in three geometrical isomers as represented below:
(iv) Geometrical isomerism cannot occur in complexes of the type MA4 , MA3B or MAB3 because all possible spatial arrangements for any of these complexes will be exactly equivalent.
(v) The square planar complexes containing unsymmetrical bidentate ligands such as [M(AB)2] also show geometrical isomerism. For example, the complex [Pt(gly)2] where gly = NH2CH2COO¯ exists in cis and trans form :
(vi) Geometrical isomerism is also shown by bridged binuclear complexes of the type M2A2X4. For example: the complex [PtCl2 P(C2H5)3]2 exhibits geometrical isomers as :
IUPAC Nomenclature of Coordination Compounds
Rules for Writing Formula
The formula of a compound is a shorthand method used to provide basic information about the constitution of a compound in a concise and convenient manner.
(1) The formula of the cation whether simple or complex is written first followed by that of the anion.
(2) The coordination entity is written in square brackets.
(3) The sequence of symbols within the coordination entity is : first the symbol of the central metal atom followed by ligands in alphabetical order.
The ligands in the coordination entity are arranged as
(a) The different ligands are arranged alphabetically according to the first symbol of their formulae. For example, H2O , NH3, NO3¯, SO42- and OH¯ etc. are cited at H, N, N, O, S and O.
(b) When the two ligands have same defining atom, the ligand with fewer such atoms is cited first followed by the ligand having more atoms. For example, NH3 precedes N2
(c) If the numbers of defining atoms are equal, subsequent symbol decides the sequence.
For example NH2¯ precedes NO2¯ because H comes before O.
(d) Polydentate ligands are also listed alphabetically. In case of abbreviated ligand, the first letter of the abbreviation is used to determine the position of the ligand in alphabetical order.
(e) The formula for the co-ordination entity, whether charged or not, is enclosed in square brackets. Polyatomic ligands are enclosed in parentheses (), but all ligands are written without any separation in between.
(f) There should be no space between the representations of ionic species within the formula.
(g) Sometimes abbreviations are used for formulae of the ligands. These abbreviations should be in lower case and enclosed in parenthesis.
For example, py is used for pyridine and en is used for ethane-1, 2-diamine or ethylene diamine.
(h) The number of cations or anions to be written in the formula is calculated on the basis that total positive charge must be equal to total negative charge.
(i) When the formula of the charged coordination entity is written without the formula of the counter ion, the charge is indicated outside the square brackets as a right superscript with the number before the sign (+ or -).
[Ni(NH3)6]2+, [Co(CN)6]3- , [CoCl(NH3)5]2+, etc.
Some common examples are :
[CoCl2(NH3)4]Cl Cl is cited first than NH3
[Co(H2O)2(NH3)4]Cl2 H is cited before N
[Pt Cl2(C5H5N)NH3] Among neutral ligands C5H5N and NH3, C5H5N precedes NH3
Rules for naming the coordination compounds
(1) Order of naming ions
In ionic complexes, the cation is named first and then the anion.
NaCl : sodium chloride
Non-ionic complexes are given a one word name.
(2) Naming the coordination entity
In naming the coordination entity, the ligands are named first and then the central metal ion.
(3) Names of ligands
The names of anionic ligands (organic or inorganic) end in o-. In general, if the anionic ligand name ends in -ide, -ite or -ate, the final ‘e’ is replaced by ’o’ giving-ido,-ito and-ato respectively. For inorganic anionic ligands containing numerical prefixes such as triphosphate enclosing marks () are added. The names of positive ligands end in -ium. The neutral ligands are named as such.
For example :
(i) Negative ligands end in -o :
- F- fluoride
- Cl- chlorido
- Br¯ bromido
- NO2¯ nitrito-N
- ONO¯ nitrito-O
- SO42¯ sulphato
- OH- hydroxo
- C2O42‾ oxalato
- CN– cyano
- CH3COO¯ acetato
- SCN- thiocyanato
- O22– peroxo
- O2- oxo
- N3– nitrido
- P3- phosphido
- N3¯ azido
- NCS¯ isothiocyanato
- H¯ hydrido
- CO32– carbonato
- NO3¯ nitrato
- NH2¯ amido
- NH2- imido
- ClO3¯ Chlorato
- S2O32– thiosulphato
(ii) positive ligands end in -ium
- NH2NH3+ hydrazinium
- NO2+ nitronium
(iii) neutral ligands are named as such
- NH2CH2CH2NH2 ethane-1,2-diamine or ethylenediamine
- C6H5N pyridine (py)
- (C6H5)3P triphenylphosphine
- PH3 phosphine
- CS thiocarbonyl
- H2NCSNH2 thiourea (tu)
However, there are a few exceptions in naming neutral ligands. For example,
- H2O aqua
- NH3 ammine
- NO nitrosyl
- CO carbonyl
(4) Order of naming ligands
When more than one type of ligands are present, they are named in alphabetical order of preference without any consideration of charge.
For example : In the complex [CoCl(NO2)(NH3)4]+, the ligands are named in the order :
ammine, chlorido and nitrito-N. Similarly, in the complex K3[Fe(CN)5NO], the ligands are named as cyano and nitrosyl.
(5) Numerical prefixes to indicate number of ligands
When more than one ligands of a particular kind are present in the complex, the di-, tri-, tetra-, penta-, hexa-, etc. are used to indicate their number: two, three, four, five and six respectively.
When the name of the ligand, includes the numerical prefix (di, tri, tetra), the prefixes bis, tris, tetrakis are used for two, three, four ligands, respectively. Such ligands are called complexligands.
For example: to indicate two simple ligands such as chloro, bromo, ammine, oxalato, etc., we use the prefix di but to indicate two complex ligands such as ethylenediamine we use the prefix bis (ethylenediamine) or bis (1,2-ethanediamine).
The name of the complex ligand is given in brackets.
For example :
[Co(en)3]3+ tris(ethane-1,2,-diamine)cobalt(III) ion
[NiCl2(PPh3)2] dichloridobis (triphenylphosphine)nickel(II)
(6) Ending of names
When the complex is anionic, the name of the central metal atom ends in -ate. For cationic and neutral complexes, the name of the metal is written without any characteristic ending.
For example, the cationic complex [Co(NH3)6]Cl3 is named without characteristic ending of the name of the metal as :
[Co(NH3)6]Cl3 hexaamminecobalt (III) chloride
The coordination compound, K[PtCl5(NH3)] which contains the anionic complex [PtCl5(NH3)]¯ is named with ending of the name of the metal as -ate.
K[PtCl5(NH3)] potassium amminepentachloridoplatinate (IV)
Similarly, the anionic complex Ca2[Fe(CN)6] is named as calcium hexacyanoferrate (II).
[Co(SCN)4]2– tetrathiocyanatocobaltate(II) ion.
For anionic complexes the Latin names of certain metals are commonly used.
For example: ferrate for Fe, cuperate for Cu, argentate for Ag, aurate for Au, stannate for
However, if the complex is cationic or neutral the name of the metal is given as such e.g., iron for Fe, silver for Ag, gold for Au, copper for Cu, etc.
K3[Fe(CN)6] Potassium hexacyanoferrate (III)
[Fe(CO)5] Pentacarbonyliron (0)
(7) Oxidation state of central metal ion
The oxidation state of the central metal ion is designated by a Roman numeral (such as II, III, IV) in the brackets at the end of the name of the complex without a gap between the two. Let us discuss some examples:
In this case the ligands are chloro and ammine. The complex is cation and chloride is anion. The oxidation state of cobalt is III as
x+ 5(0) – 1- 2 = 0
x = +3.
Its name is pentaamminechloridocobalt(III) chloride
In this case, the ligands are cyano. The complex is anionic. The oxidation state of iron is +3 as 3 (+1) +x + 6(-1) = 0
The name of the complex is potassium hexacyanoferrate(III). If the complex containing the central ion, Fe3+ is anionic, the Latin name of metal is used i.e. ferrate.
(iii) [Co(H2NCH2CH2NH2)3] (SO4)3
In this case, ligands are ethane-1, 2-diamine (or ethylenediamine). The complex is cationic. The oxidation state of cobalt is +3 as
[x + 3 x 0]2 + 3 (-2) = 0 or
2x = +6 or
x = +3
The name of the complex is tris (ethane-1, 2-diamine) cobalt(II) sulphate.
tris is used because the ligand is complex ligand.
(iv) [Ag(NH3)2] [Ag(CN)2]
Both cation and anion are complexes. The oxidation state of silver in both cationic and anionic complexes is +1. The name of the complex is diamminesilver (I) dicyanoargentate (I)
(8) Point of attachment
When a ligand can coordinate through more than one atom, then the point of attachment of the ligand is indicated by putting the symbol of the atom through which coordination occurs after the name of the ligand. For example: NO2¯ can coordinate through -N or -O. If it coordinates through N, it is called nitrito -N (or simply as nitro). If it coordinates through O, -ONO, it is called nitrito -O or simply as nitrito.
NO2¯ (through N)
-ONO¯ (through O)
Nitrito – N
Nitrito – O
SCN¯ (through S) thiocyanato
NCS (through N) isothiocyanato
[Co(ONO)(NH3)5]SO4 pentaamminenitrito-O-cobalt(III) sulphate
(9) Naming geometrical isomers
Geometrical isomers are named by the use of the terms cis-to designate adjacent positions and trans- to designate opposite positions.
For example: in square planar complexes shown below
(a) the positions 1, 4 ; 1, 2 ; 2, 3 and 3, 4 are cis-
(b) 1,3 and 2, 4 are trans
In naming these complexes, cis or trans is written before the names of these compounds. Similarly, for octahedral complexes, cis- and trans are used in isomerism in coordination compounds.
(10) Naming optical isomers
Optically active compounds are designated by the symbols (+) or d -for dextrorotatory and (-) or l- for laevorotatory.
For example : d- K3[Cr(C2O4)3] Potassium (+) trioxalatochromate (III)
- Valence bond theory:
According to this theory, the metal atom or ion under the influence of ligands can use its (n − 1)d, ns, np or ns, np or nd orbitals for hybridization to yield a set of equivalent orbitals of definite geometry such as octahedral, tetrahedral and square planar. These hybridised orbitals are allowed to overlap with ligand orbitals which can donate electron pairs for bonding.
|Coordinationnumber||Type of hybridisation||Distribution of hybrid orbitals in space|
|6||sp3d2 (nd orbitals are involved; outer orbital complex or high-spin or spin- free complex)||Octahedral|
|7||d2sp3 [(n − 1)d orbitals are involved; inner orbital complex or low-spin or spin- paired complex]||Octahedral|
- Magnetic properties of coordination compounds: A coordination compound is paramagnetic in nature if it has unpaired electrons and diamagnetic if all the electrons in the coordination compound are paired.
Magnetic moment μ=nn+2 where n is the number of unpaired electrons.
- Crystal Field Theory: It assumes the ligands to be point charges and there is an electrostatic force of attraction between ligands and the metal atom or ion. It is a theoretical assumption.
- Crystal field splitting in octahedral coordination complexes
- Crystal field splitting in tetrahedral coordination complexes
- For the same metal, the same ligands and metal–ligand distances, the difference in energy between eg and t2g level is
- Metal carbonyls: Metal carbonyls are homoleptic complexes in which carbon monoxide (CO) acts as the ligand. Example: Ni(CO)4
The metal–carbon bond in metal carbonyls possesses both σ and characters. The metal–carbon bond in metal carbonyls possess both s and p characters. The M–C σ bond is formed by the donation of a lone pair of electrons from the carbonyl carbon into a vacant orbital of the metal. The M–C bond is formed by the donation of a pair of electrons from a filled d orbital of metal into the vacant anti-bonding * orbital of carbon monoxide. The metal to ligand bonding creates a synergic effect which strengthens the bond between CO and the metal.
- Chapter 1 The Solid State
- Chapter 2 Solutions
- Chapter 3 Electrochemistry
- Chapter 4 Chemical Kinetics
- Chapter 5 Surface Chemistry
- Chapter 6 General Principles and Processes of Isolation of Elements
- Chapter 7 The p Block Elements
- Chapter 8 The d and f Block Elements
- Chapter 9 Coordination Compounds
- Chapter 10 Haloalkanes and Haloarenes
- Chapter 11 Alcohols Phenols and Ethers
- Chapter 12 Aldehydes Ketones and Carboxylic Acids
- Chapter 13 Amines
- Chapter 14 Biomolecules
- Chapter 15 Polymers
- Chapter 16 Chemistry in Everyday Life