CBSE Class 12 Chemistry Revision Notes Chapter 9
Class 12 Chemistry Chapter 9 Notes: Coordination Compounds
Coordination compounds are the backbone of advanced inorganic and bioinorganic Chemistry and are extensively used in the chemical and pharma industries. Key topics covered in Coordination compounds Class 12 Chemistry Chapter 9 are Werner’s theory, ligands, magnetic properties of coordination compound, coordination, entity crystal field theory, colour in coordination compounds, bonding, and stability constant of coordination compounds.
Extramarks is one of the leading online learning platforms in India, where students from CBSE boards refer to our study materials and NCERT Solutions. Our Chemistry subject matter experts have prepared NCERT Solutions for Class 12 Chemistry Chapter 9 detailed study notes that encourage the students to learn ‘how’ and ‘why’ and go beyond the books to learn the concepts, frame your own examples to avoid rote learning traps. Students can register on the Extramarks website and get access to Class 12 Chemistry Chapter 9 Notes.
The notes are provided in a systematic and organised chapter-wise format. Students will find easy-to-understand explanations for all the topics in our chapter notes. They are explained, with examples, diagrams, pictorial representations and flowcharts. With the help of our Class 12 Chemistry Chapter 9 Notes, students who want to learn more and nail the competition with these notes.
Key Topics Covered in Class 12 Chemistry Chapter 9 Notes
A molecule or a cation binds to the metal atom to form a coordination complex. Such an ion or molecule is called a ligand.
- These ligands function as electron-pair donors. Unidentate ligands are ligands that have one donor atom.
- Bidentate ligands are ligands that have two donor atoms.
- Ligands which have three donor atoms are called Tridentate ligands. Hexadentate ligands are ligands that have six donor atoms.
Coordination compounds are essential to recognise that life would not have been possible without the +nce of chlorophyll (Magnesium – complex) in plants and haemoglobin (Fe- complex) in human blood. Studying these compounds will broaden our understanding of chemical bonding and the physical properties of coordination compounds, such as magnetic properties.
Coordination compounds also find applications in electroplating, textiles dyeing and medicinal Chemistry. Good grades are crucial in Class 12 because you get admission into a prestigious university of your choice and also do well in their entrance exam. .
The Chemistry Class 12 Chapter 9 Notes is an outline of the chapter coordination compounds from the NCERT textbook. Coordination compounds are those chemical compounds that are composed of an array of anions or neutral molecules bound to a central metal atom via coordinate covalent bonds.
Coordination compounds are also called coordination complexes. The molecules or ions bound to the central atom are called ligands (also known as complexing agents). Besides Class 12 Chemistry Chapter 9 Notes, students are recommended to refer to a repository of resources such as NCERT Solutions and CBSE Revision Notes for a more detailed explanation, all available at the Extramarks’ website.
Definition of some Important Terms Involving Coordination Compounds:
Here are some of the definitions of coordination compounds:
- i) Coordination Entity:
Generally, a chemical compound in which the central atom or ion (or the coordination centre) is bound to a set number of atoms, molecules, or ions is a coordination entity.
Various examples of such coordination entities include [CoCl3(NH3)3] and [Fe(CN)6]4-.
- ii) Central Atoms and Central Ions:
As discussed previously, the central atoms and anions are atoms and ions attached to a set number of atoms, molecules, or ions.
In coordination complexes, the central atoms or ions are Lewis acids and can therefore appear as electron-pair acceptors.
The atoms, molecules, or ions bound to the coordination centre or the central atom/ion are ligands.
These ligands may either be simple ions or molecules (like Cl– or NH3 ammonia) or in the form of relative macromolecules, such as ethane-1,2-diamine (NH2-CH2-CH2-NH2).
- iv) Coordination Number:
The coordination number of the central atom or ion in the coordination compound is mentioned in the overall number of sigma bonds. The ligands here are forced to stay at the coordination centre.
For example, the coordination complex given by [Ni(NH3)4]2+, yields the coordination number of nickel as 4.
- v) Coordination Sphere:
The non-ionisable part of a complex compound consists of a central transition metal ion surrounded by neighbouring atoms or groups enclosed in square brackets.
The coordination centre, the ligands combined in the coordination centre, and the net charge of the chemical compound develop the coordination sphere when written together.
This coordination sphere is generally accompanied by a counter ion (the ionisable groups attached to charged coordination complexes).
For. Example: [Co(NH3)6]C/3 – coordination sphere.
- vi) Coordination Polyhedron:
The geometric shape created by the attachment of the ligands to the coordination centre is known as the coordination polyhedron.
Such kinds of spatial arrangements in coordination compounds include tetrahedral and square planar shapes.
For example, [Co(NH3)6]3+ is octahedral, [Ni(CO)4] is tetrahedral and [PtCl4]2– is square planar.
vii) Oxidation Number:
The oxidation number of the central metal atom can be calculated by finding the charge combined with it when all the electron pairs donated by the ligands are removed from it.
For Example, [Co (NH3)6] Cl3 in this, the oxidation state of NH3 is 0 and for Cl is -1, and the oxidation state of Co is taken as x, so,
(x-3) x 1 = 0 or x = +3
Another Example is [Cu (CN)4]3- in this, oxidation state of Cu is taken as x, CN as -1 then,
x + (-1 X 4) = -3
or x = +1
For Example, the platinum atom’s oxidation number in the complex [PtCl6]2- is +4.
viii) Homoleptic and Heteroleptic Complexes:
When the coordination centre atom is bound to only one type of electron pair donating ligand group, the coordination complex is known as a homoleptic complex, for Example: [Cu(CN)4]3-.
When the central coordination atom is bound to various ligands, the coordination compound is called a heteroleptic complex, which is [Co(NH3)4Cl2]+.
Students may refer to a repository of resources such as NCERT Solutions and CBSE Revision Notes and Class 12 Chemistry Chapter 9 Notes to get a more in-depth understanding about the definitions, formulae, theories and get a better hold on the chapter .
General Properties of Coordination Compounds:
The general properties of coordination compounds as covered in our Class 12 Chemistry Chapter 9 Notes:
- The coordination compounds developed by the transition elements are coloured due to unpaired electrons, which absorb light during the electronic transitions.
For example, the complexes containing Iron(II) can show green and pale green colours, but the coordination compounds containing iron Fe(III) have a brown or yellowish-brown.
- When the coordination centre is a metal ion, the corresponding coordination complexes are magnetic because of unpaired electrons.
- Coordination compounds show different chemical reactivity. It can be a part of inner-sphere electron transfers and outer-sphere electron transfer reactions.
- Complex compounds with specific ligands can aid in transforming molecules in a catalytic or a stoichiometric manner.
Double Salts and Coordination Complex:
As described in our Class 12 Chemistry Chapter 9 Notes, double salts are completely ionised in aqueous solutions, and every single ion in the solution gives the confirmatory test.
For example, Potash Alum is double sulphate. It is K2SO4.Al2(SO4)3.
24H2O on Ionisation gives:
K+, SO42− and Al+3 ions respond to the corresponding tests.
Coordinate complexes are incompletely ionisable in the aqueous solutions. These give a complexation which does not show complete Ionisation.
For example, Potassium Ferrocyanide. [K4Fe(CN)6] It ionises to give K+ and [Fe(CN)6]−4 [ferro cyanide ions].
Types of Coordination Complexes:
- Cationic complexes: In this coordination sphere is a cation. Example: [Co(NH3)6]Cl3
- Anionic complexes: In this coordination sphere is Anion. Example: K4[Fe(CH)6].
- Neutral Complexes: In this coordination sphere are neither cation nor anions. Example: [Ni(CO)4]
- Homoleptic complex: The complex consists of a similar type of ligands. Example: K4[Fe(CN)6]
- Heteroleptic complexes: These consist of different types of ligands. Example: [Co(NH3)5Cl]SO4
- Mononuclear complexes: In this coordination sphere has a single transition metal ion. Example: K4[Fe(CN)6] and
- Polynuclear complexes: This coordination sphere has more than one transition metal ion.
IUPAC Nomenclature of Coordination Compounds:
Coordination compounds are Formulated and adequately named according to the IUPAC system (International union of pure and applied chemistry). It is essential for writing systematic names and formulas, mainly when dealing with isomers.
It is elementary to get information about the constitution of compounds if we understand the formula. The standard rules followed in coordination compounds’ nomenclature are summarised below. For a more elaborative description of them students can refer to our Class 12 Chemistry Chapter 9 Notes.
- The central atom is listed first.
- The ligands names are always written before the central metal ion in naming complex coordination complexes.
- Suppose the coordination centre is bound to more than one ligand. In that case, the names of the ligands are listed in an alphabetical order which is not affected by the numerical prefixes(charge) that applies to the ligands.
- The prefixes that give some insight into the number of ligands, when several monodentate ligands are present in the coordination compound, are of the type like di-, tri-, tetra-, etc.
- If many polydentate ligands are attached to the central metal ion, the prefixes are bis-, tris-, etc.
- The names of the anions present in coordination compounds must end with the letter ‘o’, which generally replaces the letter ‘e’. So, the sulphate anion must be written as ‘sulfate’, and the chloride anion must be ‘chloride’.
- The following neutral ligands are allotted specific names in coordination compounds: NH3 (ammine), H2O (aqua), CO (carbonyl), and NO (nitrosyl).
- After the ligands are named, then the name of the central metal atom is written. The suffix’- ate’ is applied if the complex has an anionic charge.
- When writing the name of the central metal atom in an anionic complex, priority is stated to the Latin name of the metal atom if it exists (except for mercury).
- The oxidation state of the central metal atoms/ions is specified with the help of Roman numerals, which are enclosed in a set of parentheses.
- If a counter ion accompanies the coordination compound, the cationic entity must be written before the anionic entity.
As for ionic compounds, the naming is completed by writing the first cation name and then naming the Anion. For example, NaCl is ionic and written as sodium chloride. The first part is cationic, and the other is anionic.
Coordination compounds consist of two parts:
While naming, the name of the cationic part is written first, followed by the anionic part. There are certain sets of rules followed while writing the name of the cation.
Besides Class 12 Chemistry Chapter 9 Notes, students are recommended to refer to a repository of resources such as NCERT Solutions, CBSE important questions, etc. for a more comprehensive understanding and also learning about the question patterns.
Naming the Cationic Species in which the Coordination Sphere is Positively Charged:
- The name of the metal cation is written first. The coordination sphere is cationic with a positive charge.
For Example, [Co (NH3)6] Cl3 for naming the coordination entity is done first and later the counter ion. So in a given complex, the name of [Co (NH3)6] is written first, then the counter ion Cl.
- In the case of a coordination entity, the ligands’ name is written first and then the central metal atom.
For Example, in this [Co (NH3)6], the name of NH3 is written first rather than for Co.
- You can write homoleptic ligands in any manner, whereas heterolytic ligands are written alphabetically.
For example, in complex [Co (NH3)5Cl], the name of ligands is written alphabetically for ammine first and then chloro.
Rules for writing the name of the Ligands as given in our Class 12 Chemistry Chapter 9 Notes:
- When the ligand names end with -ate or -its, then e is replaced by o
As for oxalate, it is written as oxalate, sulphite as sulfite
- If the ligand name ends with -ide , then ide is replaced by -o.
For example, chloride becomes chloro.
- Neutral ligands are named as such; for water, it is aqua; for ammonia, it is ammine.
- If more than one ligands are present, then alphabetical order is followed, and di, tri, and tetra are prefixed before the ligand’s name.
For example, in complex [Co (NH3)6] ligand name will be hexamine, that is, Hexa for 6.
- Polydentate ligands include numerical prefixes –like di is replaced by bi, tri is replaced by tris, tetra replaced by tetra etc.
For example, [CoCl2(en)2] Cl in this, the ligands are dichloro bis ethylenediamine
In case of the Central Metal Atom:
The Class 12 Chemistry Chapter 9 Notes states, the oxidation state of the central metal cation is written in a numeral after the name of the central metal atom.
For example, in complex [Co (NH3)6] Cl3, it is hExamine cobalt (III) that is cobalt (III) is a central metal atom with its Oxidation state in numeral
- If a complex is cationic, then usually, the name is used. An example will be the same as given above.
- If the complex is anionic, then the metal atom name ends with -ate
For example, K3[Fe(CN)6] in this complex, the name is written as potassium hexacyanoferrate(III); which is, in this example, a coordination entity and is an anionic complex, and the name of the central metal cation is written with -ate followed by oxidation state in the numeral.
Note: The typical central metal atom name is used if the complex is neutral.
Some Commonly used Ligands are:
– Bromo( Br–), floro( F–), oxo, hydroxo( OH–), cyano (CN–), carbonato (CO32), –acetate( CH3COO–), amine( NH3), aqua (H2O), nitrosyl( NO), carbonyl (CO), dioxygen( O2), dinitrogen( N2), pyridine(C5H5N), ethylenediamine( H2NCH2CH2NH2).
Some examples of IUPAC naming will make the rules clear:
- [Cr (NH3)(H2O)3]Cl3
- [Ag(NH3)2][Ag(CN)2] is named as
- [Pt (NH3)2Cl (NO2)]
The surrounding atoms, anions and molecules around the central transition metal ion are called Ligands. They act as Lewis bases and donate electron pairs to transition metal ions, forming a dative bond between ligands and the transition metal cation. Hence these compounds are coordination complexes.
Different Types of Ligands as explained in Class 12 Chemistry Chapter 9 Notes:
Depending on the nature of the bond between the ligand and the central atom, ligands are classified as follows:
- Anionic ligands: CN–, Br–, Cl–
- Cationic ligands: NO+
- Neutral ligands: C.O., H2O, NH3
Ligands can be further classified as:
- Unidentate Ligands:
Unidentate ligands are ligands that can bind to the coordination centre. Ammonia(NH3) is an excellent example of a unidentate ligand. Some common unidentate are Cl–, H2O etc.
- Bidentate Ligands:
Ligands that can bind to the central atom via two separate donor atoms, such as ethane-1,2-diamine, are bidentate.
Oxalate ions is a bidentate as they can bond through two atoms to the central atom in a coordination compound and Ethane-1, 2-diamine:
Structure of EDTA(ethylene diamine tetraacetate ion)
Source: NCERT textbook
- Polydentate Ligands:
Some ligands have several donor atoms which can combine to the coordination centre. These ligands usually are called polydentate ligands.
A polydentate ligand [ EDTA4- ion (ethylene diamine tetraacetate ion)] can bind to the coordination centre via its four oxygen (O)atoms and two nitrogen(N) atoms.
- Chelate Ligand:
When a polydentate ligand binds itself to the same central metal atom through 2 or more donor atoms, it is called a chelate ligand. The atoms which ligate to the metal ion determine the denticity of such ligands.
- Ambidentate Ligand:
Some ligands can attach to the central atom via the atoms of two different elements called ambidentate ligands.
For example, the thiocyanate (SCN– )ion can bind to a ligand via the nitrogen or sulphur atoms. Such ligands are known as ambidentate ligands.
To understand the above phenomena, visit the Extramarks website for more study materials such as Class 12 Chemistry Chapter 9 Notes, CBSE Important Questions, CBSE Extra Questions, etc.that can help them understand the important points easily and make them revise quickly.
Isomerism in Coordination Compounds:
Two or more than two compounds with the same chemical formula but a different arrangement of atoms are known as isomers. Its classification is solely based on the difference in the atom arrangement.
Coordination compounds predominantly show two isomerism types:
- Structural isomerism
Each of which can be further subdivided.
As explained in our Class 12 Chemistry Chapter 9 Notes, coordination compounds with the same chemical formula and chemical bonds but different spatial arrangements are called stereoisomers. These isomers are further divided into optical isomerism and geometrical isomerism.
- Optical Isomerism in Coordination Compounds: As per the definition compiled in CBSE Solutions Class 12 Chemistry Chapter 9, Optical Isomers are mirror images which cannot be superimposed on one another. These are called enantiomers. The molecules or ions that can be non-superimposable mirror images are chiral.
These isomers are of two forms:
- i) The isomer that rotates plane-polarised light towards a clockwise direction is the Dextro or ‘d’ or ‘+’ isomer.
- ii) The isomer that rotates plane-polarised light in an anticlockwise direction is the leave isomer or ‘l’, ‘-‘ isomer.
The racemic mixture of ‘d’ and ‘l’ isomer is the equimolar mixture.
Example of Optical Isomerism:
- Geometrical Isomerism: Geometrical isomerism is observed in heteroleptic complexes (complexes with more than one type of ligands) due to various possible geometric dispositions of the ligands. This behaviour is commonly observed in coordination compounds with the coordination numbers 4 and 6.
Geometrical isomerism of complexes having coordination number as 4:
- As ligands are in different directions, ML4 tetrahedral complexes do not show cis-trans isomerism.
- MABCD has three geometrical isomers. 2-cis and 1-trans.
- MA2B2 complex shows cis and trans isomers.
- ML6 octahedral complex never shows geometrical isomerism.
- MA2B4 complex shows cis-trans isomerism.
Another type of Geometrical isomerism that occurs in octahedral compounds in MA3B3 like [Co(NH3)3(NO3)3].
It can be of two types:
- Facial: In this, three donor atoms of the identical ligands occupy adjacent positions at the corners of an octahedral face
- Meridional: In this, three donor atoms of the identical ligands occupy positions around the meridian of an octahedron.
Structural isomerism is shown by the coordination compounds having the same chemical formula but a different arrangement of atoms.
These are further divided into four types:
- i) Linkage Isomerism:
Linkage isomerism is shown by coordination compounds having ambidentate ligands. It happens in the case of ambidentate ligands, where they differ in the point of attachment.
For example, in the case of CN (cyano) and the case of NC (isocyanate)
[Co(NH3)5No]SO4 and [Co(NH3)5ONO]SO4
In coordination isomerism, the exchange of ligands between cationic and anionic entities having different metal ions present in coordination compounds occurs.
For example, [Co(NH3)6][Cr(CN)6] and [Cr(NH3)6][Co(CN)6] are coordination isomers.
Ionisation isomerism is formed when the counter ion in a complex salt, a potential ligand, replaces the ligand.
For example: [Co(NH3)5(SO4)]Br and [Co(NH3)5Br]SO4.
Solvate isomerism is a particular case of ionization isomerism in which compounds differ depending on the number of solvents (water molecules) directly bonded to the metal ion.
Bonding in Coordination Compounds:
Alfred Werner explained the first theory in 1892.
He performed different experiments to show that the surrounding atoms exist around the central atom. He conducted a precipitate study.
When CoCl3.6NH3 was precipitated with an AgNO3 solution, it gave three moles of AgCl. This complex shows that 3 Chloride ions are not directly attached to cobalt; that’s why it was precipitated with silver nitrate solution. This gave him a trick about primary and secondary valencies.
To understand more about the above Bonding in Coordination Compounds, visit the Extramarks website to access our Class 12 Chemistry Chapter 9 Notes.
As per the theory compiled in CBSE Solutions Class 12 Chemistry Chapter 9, Alfred Werner, in 1898, proposed a theory to understand the structure of coordination compounds.
By mixing AgNO3(silver nitrate) with CoCl3·6NH3, all three chloride ions got converted to AgCl (silver chloride). But, when AgNO3 was mixed with CoCl3·5NH3, two moles of AgCl were formed.
Further, on mixing CoCl3·4NH3 with AgNO3, one mole of AgCl was formed. Based on this observation, the following theory was postulated:
Postulates of Werner’s Theory:
- The central metal atom in the coordination compound exhibits two valency types: primary and secondary linkages or valencies.
- Primary linkages are ionisable and are satisfied by the negative ions.
- Secondary linkages are non-ionisable. These are satisfied by negative ions. Also, the secondary valence is fixed for any metal with an equal coordination number.
- The ions bound by the secondary linkages to the metal exhibit spatial characteristics. The arrangements correspond to different coordination numbers.
Difference between Primary and Secondary Valency in Coordination Compounds:
- Werner’s Theory
Primary valency: These are ionisable and satisfied by charged ions. Primary valency does not support the structure of the complex. It can function as a secondary valence too.
Secondary valency: These are Non-ionisable and are satisfied by ligands. Secondary valency helps in structure. However, it can not work as a primary valency
CoCl3.6NH3 is a Werner complex.
Such spatial arrangement is termed coordination polyhedra. The terms inside are coordination complexes, and the ions outside are counter-ions. Dark lines show ionisable parts, and light lines show non-ionisable parts of a werner’s complex.
Limitations of Werner’s Theory:
- It fails to explain coordination compounds’ magnetic behaviour, colour, and optical properties.
- It failed to explain why all elements do not form coordination compounds.
- This theory does not explain the directional properties of bonds in coordination compounds.
- It does not understand the stability of the complex and the nature of complexes.
- This theory could not distinguish between weak and strong ligands.
He was able to explain various facts about coordination compounds. Still, He failed to explain why only certain elements participate in coordinate bonds and why the coordination entity has a unique geometry. Due to these reasons, another theory was proposed, ‘The Valence Bond Theory.
Students can refer to our Class 12 Chemistry Chapter 9 Notes for more information about Werner’s Theory.
Valence Bond Theory:
Under CBSE and with the help of Class 12 Chemistry Chapter 9 Notes, students learn about VBT theory. Pauling gave it in 1931 when he proposed the idea of donating lone pairs to the central metal atom.
Bonding in coordination compounds occurs due to the overlap of the orbital of a ligand with a vacant orbital of the central metal atom. All the vacant d orbitals have the same energy, but the degeneration of the d orbital breaks when ligand approaches.
Hybridisation is considered while drawing polyhedral. Metal ions in the presence of ligands can use their (n-1)d ns np or ns np Nd. If the inner d orbital is used, then the complex is regarded as an inner orbital complex, and if the outer d orbital is used, then the complex is an outer orbital complex. The ligands conclude which orbitals out of these to be used, and accordingly, the geometry is decided.
The number of orbitals and types of hybridisations:
|Types of Hybridisation||Distribution of Hybrid Orbitals in Space|
For example: To find the shape using valence bond theory (VBT) following ways should be used for any coordination compounds::
- Remove the electrons from the metal and create the ion.
- Rearrange metal electrons if necessary.
- Overlapping of hybrid orbitals of metal ions with ligands.
Let us consider an example, [Co(NH3)6]3+. In this complex, the central metal atom ‘Co’ has an atomic no. of 27. The electronic configuration of Co = Ar183d74s2Co3+= (Ar)183d6
Example: [Fe(Co)5]: (inner orbital complex and diamagnetic)
Example: in [CoF6]3-…… (outer orbital complex and paramagnetic)
Image source: NCERT textbook
Drawbacks of Valence Bond Theory:
- This theory couldn’t explain why some complexes of metal oxidation state form inner orbital complexes, while in a few other complexes, the same metal atom ion in the same state forms the outer orbital complex.
- The magnetic behaviour explained wasn’t satisfactory.
- The colour exhibited by coordination compounds could not be explained by this theory.
- The theory couldn’t distinguish between strong and weak ligands.
Crystal Field Theory:
It was given by Hans Bethe and John van Vleck who states that anionic ligands are considered negative point charges and neutral ligands as point dipoles.
- It assumes the central metal cation and ligands as point charges. A central metal atom has a positive charge when a complex is formed.
- Ligands –have a negative charge.
- This theory considers the electrostatic interaction between the central metal atom and ligand. There is no intermixing of atomic orbitals with metal orbitals or no insertion of electrons.
- When a complex is created, the central metal atom is surrounded by oppositely charged ligands.
- No hybridisation takes place.
- The ligand molecule must approach the central metal atom to form a bond.
- The d orbital of the central metal atom degenerates in the absence of an external magnetic field. Still, in most complexes, this degeneration is lost because the field produced by the ligand is asymmetrical and all d orbitals are not equally affected by the ligand field.
- Loss of degeneracy leads to the splitting of d orbitals and their energies called crystal field splitting(CFS).
- The d orbital splits into two sets: The axial set is dxy,dyz, dzx and the non-axial set is dx2-y2,dz2.
- Repulsive forces form between electrons of metal and lone pair ligands; electron energy fluctuates or changes.
The Crystal Field Theory has been explained in detail with illustration by our Chemistry subject matter experts in our Class 12 Chemistry Chapter 9 Notes. CBSE students across India have already registered on Extramark’s official website and are getting the best learning experience. These notes are written in a systematic and organised manner making it convenient for students to remember everything clearly.
Crystal Field Splitting in Octahedral Coordination Entities:
Metal atoms or ions surround six ligands. As an example, orbitals are more directional toward orbitals. They show greater repulsion and higher energy. T2 g orbitals lie away from ligands. Thus they have lesser repulsion and lower energy.
In an octahedral complex, ligands have to approach the central metal atom along the coordination axis. The d orbitals whose lobes lie along the axis will experience more repulsion during the approach.
Due to this, their energy will enhance, and the other non-axial set will suffer less repulsion. In conclusion, the non-axial will have less energy as compared to the axial set (Army greater than t2g)
Crystal Field Splitting In Tetrahedral Entities:
The ligands have to approach the central metal atom between the coordination axis. The d orbitals with the lobes lying along the axis will experience less repulsion during the approach. Their energy will increase, and the non-axial set will suffer more repulsion.
As a result, the non-axial will have more energy than the axial set (t2g more significant than the example). The coordination number and number of ligands is 4. The T2g set has higher energy than example, as t2g orbitals are closer to the direction of approach of ligands.
For all the Complexes:
If the ligands are strong, the CFSE is more; therefore, pairing will occur, and for weak ligands, the CFSE is less.
Limitations of CFT:
Assuming ligands as point charges, it considers anionic ligands as the excellent source of the splitting effect. The anionic ligands are found at the low-end level of the spectrochemical series.
Effective Atomic Number Rule:
Sidgwick proposed an effective atomic number rule. The total number of electrons passed by the central transition metal ion after the donation of electrons by the ligand is an adequate atomic number.
A complex is stable if the adequate atomic number equals the atomic number of the nearest inert gas.
Example: Find out the effective atomic number of the following complexes?
Number of electrons in Fe2+ = 24
Number of electrons by Six CN = 2×6 = 12
Total number of electrons possessed by Fe2+ = 24 + 12
Therefore, the effective atomic number = 36.
Number of electrons in Co+3 = 24
Number of electrons by Six NH3 = 2×6 = 12
Total number of electrons possessed by Co+3 = 24 + 12
Therefore, the effective atomic number = 36.
Students can get more examples of effective atomic number rules from our Class 12 Chemistry Chapter 9 Notes. Along with chapter notes, students should refer to a repository of resources
such as NCERT Extra Questions, CBSE Revision Notes, and NCERT Solutions for Class 12 Chemistry chapters to solve exercises and practice questions and clarify their doubts to master the topic.
Magnetic Properties of Complexes:
- The complex in which the central transition metal ion has unpaired electrons is paramagnetic.
- The complex in which the central transition metal ion has no unpaired electrons is called diamagnetic.
- The magnetic moment of a complex is determined by the spin only formula.
M = √[n(n+2)] BM
BM = Bohr Magneton
The magnetic moment of complex compounds relies upon:
- Type of Hybridisation.
- The oxidation state of the central transition metal ion.
- The number of unpaired electrons.
As explained in our Class 12 Chemistry Chapter 9 Notes, the chemical series of increasing order of field strength of ligands is called the Spectrochemical series. These spectrochemical series are I– < Br– < SCN– < Cl– < S-2 < F– < OH– < C2O4-2 < H2O < NCS− < (EDTA)-4 < NH3 < en < CN < CO
Stability of Complexes:
A complex is formed in various steps. Each step is reversible, and the equilibrium constant is known as the stepwise formation constant. Let us observe the formation of complex ML4
M + L ⇄ ML1 in the presence of K1
ML1 + L ⇄ ML2 in the presence of K2
ML2 + L ⇄ ML3 in the presence of K3
ML3 + L ⇄ ML4 in the presence of K4
The net formation constant or stability constant, β equals K1 × K2 × K3 × K4, and 1/β is equal to the instability constant.
Factors Affecting Stability of Complexes:
The small size and high nuclear charge of central transition metal ions affect the stability of complexes. Crystal field stabilisation energy (CFSE) should be more. Complexes containing chelating ligands are more stable. Octahedral complexes are commonly more stable than tetrahedral.
To understand more about the factors affecting the stability of complexes, students are suggested to visit the Extramarks’ website and get access to our Class 12 Chemistry Chapter 9 Notes for thorough understanding of the topic.
Colour of Complexes:
Complexes in which the central transition metal ion contains unpaired electrons show colour. It is – d’ transition. The colour of complexes depends upon:
- Several unpaired electrons in transition metal ions.
- Nature of ligands.
- The oxidation state of the central transition metal ion.
- The light wavelength absorbed and emitted.
- The proportion of ligands in the coordination sphere.
For Example: [Ni(H2O)6]2+ + en(aq)→[Ni(H2O)4en]+2 – Green change into Pale blue
Bonding in Metal Complexes: [Metal Carbonyls]
Organometallic compounds: These are compounds in which metal, metalloid, or nonmetal is directly linked to a carbon atom of a hydrocarbon.
For Example (C2H5)2Zn etc
Please Note that metal cyanides and carbides are not organometallic compounds as the carbon atom is not directly joined to metal.
Types of organometallic compounds:
- Sigma organometallic compounds
- pi organometallic compounds
Sigma Organo-metallic Compounds:
These are obtained by bonding nonmetals with metalloid elements with carbon.
For example, RMgX, (CH3)3Al etc.
Pi Organo- metallic compound:
These compounds are created mainly by transition elements. A regular sigma bond is formed through the pi-electron cloud of an organic molecule—for example, ferrocene, Zeise’s salt etc.
Sigma and pi organo-metallic compound:
Transition metal carbonyls form these compounds.
For example, Ni(CO)4, Fe(CO)5 etc.
Complexes where carbon monoxide acts as ligands are known as metal carbonyls.
Example: [Ni(CO)4] Tetracarbonyl Nickel (0) and [Fe(CO)5] Penta Carbonyl Iron (0). With the use of these complexes, an a′σ’ bond is created by the overlapping of the vacant ‘d’ orbital of metal ion and a filled orbital of C-atom (carbon).
A π bond is mainly formed by the lateral overlapping of filled inner orbitals of metal ions and the vacant carbon atom. Thus synergic bonding exists in metal carbonyls.
Shapes of these structures are shown below:
- The shape of [Ni(Co)4], [Fe(Co)5]
Considering Bonding in Metal Carbonyls:
As per Class 12 Chemistry Chapter 9 Notes, these compounds possess both s and p characters. The Metal-Carbon sigma bond is formed by donating a lone pair of electrons on the carbonyl carbon into a vacant orbital of metal. The Metal-Carbon pi bond is formed by donating a pair of electrons from the filled d orbital of metal into the vacant antibonding pi orbital of carbon monoxide ion. The metal to ligand bonding develops a synergic effect which strengthens the bond between CO and metal, as shown below:
Example of synergetic bonding interactions in a carbonyl complex
Applications of Coordination Compounds:
The characteristic properties of coordination compounds discussed in the previous subsection are beneficial in various processes and industries. Some of these applications of coordination compounds are mentioned below. In case you want to study in detail, an elaborate explanation of these applications has been covered in our Class 12 Chemistry Chapter 9 Notes.
- It is used to detect and estimate metal ions in qualitative and quantitative analysis. For example, EDTA, DMG, alpha nitroso, beta naphthol, cupron etc., give colour reactions.
- The colour of the coordination compounds formed with transition metals is extensively used in industries to colour materials. It is used in the dye as well as the pigment industries.
- Some complex compounds with cyanide as a ligand are used in electroplating. These compounds are also very much utilised in photography.
- Coordination complexes are very useful in extracting metals from their ores. For example, nickel and cobalt can be extracted from their ores by hydrometallurgical processes involving ions of coordination compounds.
- The hardness of water is determined by titration of calcium ions and magnesium ions with Na2EDTA, and their estimation can be concluded due to the difference in their stability constants.
- Extraction of metals like silver and gold makes use of complex compounds. For example, gold forms.
- Complex[Au(CN)2]– in an aqueous solution and can be separated into metallic form by adding zinc.
Applications In Biology:
As explained in the theory notes of our Class 12 Chemistry Chapter 9 Notes, pigments are responsible for photosynthesis in biological systems, and chlorophyll is a complex of Magnesium.
- Haemoglobin consists of Heme complex-ion, a tetrapyrrole Porphyrin ring structure with a central Fe2+ ion.
- Vitamin B12 consists of a tetrapyrrole porphyrin ring complex with a central Co+3 ion, and its coordination number is 6.
Applications In Laboratory:
- Ni+2 ion is estimated using a complexing agent Dimethylglyoxime (DMG). The hardness of water is calculated using complexes of Ca++ and Mg++ with EDTA.
- Complexes, like rhodium complex,[(Ph3P)3RhCl], a Wilkinson catalyst, are used to hydrogenate alkenes.
- In Medicine: Cisplatin is widely used in the treatment of cancer.
- In Photography: Developing the film involves complex formation.
- In Metallurgy: In the extraction of gold and silver by the Mcarthur Forest Process involves a complex of cyanide ions.
- Articles can be electroplated with Ag and Au much more smoothly and evenly from a solution of complexes, [Ag(CN)2]– and [Ag(CN)2]– than from a solution of metal ions.
- The chelate theory is used in the medicinal industry. For Example, EDTA is used in the treatment of lead poisoning. Some complexes of platinum can inhibit the growth of tumours. For example. Cis platin[PtCl2(NH3)2] and related compounds.
- The developed film in black & white photography is fixed by washing with a hypo solution, which dissolves the undecomposed AgBr to form complexes [Ag(S2O3)2]3-.
CBSE Solutions Class 12 Chemistry Chapter 9: Exercise & Solutions
Extramarks NCERT solutions and Chapter Notes also contain a lot of exercises and solutions for students to practise. These exercise and answer solutions are based on the latest NCERT books. Every exercise is compiled to add more value to your learning experience with the Extramarks. Students may refer to a repository of resources such as Revision Notes, past year questions papers, and essential questions and learn more about NCERT solutions on Extramarks. These solutions are comprehensive and provide detailed explanations to ensure that the students enjoy the learning process and can handle any tricky question with ease.
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- Chapter 9: Exercise 9.1 Solutions
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CBSE Exemplar Class 12 Chemistry Chapter 9
All solutions and problems are provided to help students prepare for their final examinations. These Exemplar questions are a little more complex, and they cover all the concepts stated in each and every chapter of the Class 12 Chemistry. .
Students will understand all the concepts covered in their syllabus by practising Chemistry CBSE Exemplar for Class 12.
Exemplars provide the best solutions to all those challenging and tricky questions that students face in their exams. To match the ideas taught in class and provide the most outstanding revision materials or worksheets for students, the subject matter experts strictly follow updated 2021-22 CBSE guidelines.
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Key Features of Class 12 Chemistry Chapter 9 Notes
The coordination compound is one of the important l topics for tests and competitive examinations like IIT, JEE and NEET. Hence while studying Chemistry, these concise notes and solutions provided by Extramarks will prove useful for the students. Class 12 Chemistry Chapter 9 Notes can help students understand the topics covered under the chapter easily and quickly.
Following are the key features of Class 12 Chemistry Chapter 9 Notes:
- . Comprehensive and detailed notes written in a simple and easy language.
- These notes have been prepared by experienced subject matter experts while adhering to the latest CBSE guidelines.
- These notes include a thorough explanation of each and every topic covered in the chapter.
- It can be used to study right after school for assignments and tests or as revision notes before examinations.
- These notes are equally helpful for board examinations and various competitive examinations like NEET, JEE Main, Advanced .
- During exams you won’t need any other help or assistance if you are using Extramarks resource which has solutions for all your queries. Don’t forget to pick the right study material to step up your academic performance and stay ahead of the pack.
FAQs (Frequently Asked Questions)
1. Can the Chapter 9 Chemistry Class 12 Notes be used as Revision Notes?
Absolutely! Chemistry Chapter 9 Class 12 Notes can be used as Revision Notes as it gives all the information students need to understand the chapter. It allows students to study for the examinations without being stressed and anxious. These notes are the concise, authentic study material which is the ultimate tool for students to speed up their preparation and improve their grades.
2. How important are the CBSE Solutions for Class 12 Chemistry Chapter 9?
The Chemistry Class 12 Chapter 9 CBSE Solutions help students understand all the concepts and give a detailed explanation of , different equations, theories and complex formulas, definitions etc . It allows students to study for the examinations without depending on teachers or parents. . Students may refer to CBSE Solutions and Class 12 Chemistry Chapter 9 Notes on the Extramarks website to check the plethora of resources available and make the most of it.
3. Are all the chapter topics covered under Class 12 Chemistry Chapter 9 Notes?
Yes, Extramarks Class 12 Chemistry Chapter 9 Notes cover all topics and provide students with the key information to quickly understand and revise the chapter.
4. Give two examples of strong-field and weak field ligands?
NH3 and H2O are weak whereas CO and CN are strong field ligands.