Alfred Werner, a Swiss chemist and the founder of coordination chemistry, won the Nobel Prize for chemistry in 1913 for his research into the structure of coordination compounds. Prior to him, the concepts of valence bonding and geometry in metal complexes were confusing. He revolutionised the fields of inorganic chemistry and stereochemistry and found applications in many fields such as organic chemistry, analytical chemistry, biochemistry, geochemistry and mineralogy. He proved that stereochemistry is not limited to organic chemistry but is a general phenomenon.
Coordination chemistry is the study of a class of compounds formed by metals. Coordination compound can be explained by the following example:
K4[Fe(CN)6]—Potassium ferrocyanide complex can prepared by an excess of aqueous potassium cyanide being added to aqueous ferrous sulphate, formed as a yellow colour solution. The complex K4[Fe(CN) 6] is called “coordinate compound”. The formation of coordination compound from a metal is called “complexation”. The complex solution contains [Fe(CN) 6]4− and is called “metal complex ion”. It can be isolated as the part [Fe(CN) 6]4− in K4 [Fe(CN) 6] and is called “complex species (or) complex entity”. The formula of the complex species should be represented within brackets—[ ].
Examples:
A simple salt is dissolved in water and produce its constituent simple ions.
Example: FeSO4→Fe+2 + SO42−
However, metal complex is dissolved in water and produces complex ion.
Example: K4[Fe(CN)6] → [Fe(CN)6]4− + 4K+
Example for water soluble complex: [Ag(NH3)2]+(diamino silver(I)
Example for water insoluble complex: [Ni((CH3)2—C=N—OH)] Nickel dimethylglyoxime
Some complexes are soluble in organic solvent.
Example: Bis(acetylacetonate)Cu(II)
Double salt is the compound formed by combining two stable compounds.
Example: Ferric alum Fe2(SO4)3·(NH4)2SO4·24H2O
When double salts are dissolved in water, they will completely dissociate into simple ions. For example, when crystals of ferric alum are dissolved in water, the solution shows properties of NH4+, SO42− and Fe2+ ions.
Coordination compound is formed with metal and ligands by dative bonds. When it dissolves in water, it does not break down to simple ion; only the ionic bond breaks and gives metal ion and complex ion.
Example: K4[Fe(CN)6] When dissolved in water the solution, it consists of K+ and [Fe(CN)6]4−; the complex ion does not break down to Fe+2 and CN−2 ions.
The total number of monodentate ligands attached to the central metal complex is called coordination number; this is equal to number of sigma bonds between the metal and ligand. Coordination number from two to nine is known in complex; four and six are common coordination numbers but three is a rare coordination number.
The number denotes the charge on the central metal atom if all the ligands in the complex were removed along with their electron pairs; it is represented by Roman numerals.
Eg.:
[NiCl4]2−
In the above complex ion if all four chlorine ligands are removed from complex ion then the central atom Ni attain a charge of +2. The oxidation number of this metal in this complex is written as II.
[Cr(NH3)6]3+
In the above complex ion if all six ammonia ligands are removed from complex ion then the central atom Cr attain a charge of +3. The oxidation number of this metal in this complex is written as III.
A ligand is an ion or a molecule capable of donating an electron pair by participating in the formation of a coordinate bond. Depending on the bonding formation, ligands are divided into unidentate, polydentate, bridging, ambident, flexident ligands, etc.
Unidentate ligand is bound by only one atom at a time as a donor and fills only one coordination position of a given cation.
Example: Cl−, Br−, I−, CN−, SCN−, NO3−, NO2−, R3N, pyridine, CO, H2O, etc.
Poly dentate ligand is bonded by more than one atom at a time as a donor and fills more coordination positions of a given cation. Ligands and their bonding capacity given in Table 14.1.
Table 14.1 Ligands and their bonding capacity
Bridging ligand binds simultaneously with two metals by forming a bridge between the two metals and acting as a ligand to both metals.
Ambidentate ligand is capable of bonding through one type of donor atom in one complex but through a different donor atom in another complex. Such a ligand has potentially two ligating atoms. Some ambidentate ligands and their names are given in Table 14.2.
Table 14.2 Ambidentate ligands and their names
A flexidentate ligand can bind to a metal with different numbers of coordinate sites.
Example: EDTA is normally a hexadentate ligand but can act as a pentadentate or tetradentate ligand depends on ligand present at that particular situation; hence, it is called a flexidentate ligand.
Metal complexes are named systematically based on certain rules recommended by the International Union of Pure Applied chemistry (IUPAC).
The cationic complex names begin with the number of ligands, followed by the name of the ligand bonded to the central metal atom (or) ion. The oxidation number of metal is indicated by Roman numerical parenthesis.
The complex name will be written without space between ligand and metal names and the metal name and the parenthesis in a single entity.
In normal salts like BaCl2, the salt is named barium chloride and not barium dichloride because the bivalency of barium is implied. Similarly, [Cr(NH3)6]Cl3 is not named tri chloride; but simply chloride. The first part of the name of the complex Hexaamminechromium(III) automatically indicates that it is of three chlorides.
The name of the anionic complex will be mentioned by the adding of -ate at the end of the metal ion; however, complex positive ions or neural molecules do not have such an ending.
Example: [Fe(CN) 6]-4 hexacyanoferrate—here, the suffix “ate” is attached to the name of the metal.
Anionic complexes having simple cations like Na+, K+, etc., where the complex begins with the name of a cation.
Example: K4[Fe (CN)6] potassium hexacyanoferrate(II)
The complex ion can put in a bracket [ ]. This indicates that the ions are expected to be formed in a solution when the complex is dissolved in solvent. Some common complex anion shown in Table 14.3.
Table 14.3 Names of complex anions containing metal atoms
Negative Ligands
The name of a negative ligand ends in O.
Example:
Neutral Ligand
Neutral ligands are named as such without any special name ending.
Example:
Positive Ligand
Positive ligands end with “ium”.
Example: NH2-NH3+ - hydrazinium
Example: [PtCl2(H2O)(NH2NH3)+] Dichloroaquahydrozoniumplatinum(II)ion
[CrCl3(NO2)3]3− Trichlorotrinitrochromate(III)ion
[Ag(NH3)2]+ - Diamminesilver(I)ion.
However, when the name of the ligand itself has a prefix such as di, tri, tetra, etc., to avoid confusion, prefixes such as (di) = bis, (tri) = tris, (tetra) = tetrakis are used to indicate the number of such ligands present in complex.
[Cu(en)2]SO4 -bis(ethyldiamine)copper(II)sulphate
In isomeric complexes, the ligand is attached through different donor atoms to the metal.
Designating the Ligand Atom in a Polydentate Ligand
A polydentate ligand has more than one donor site, some (or) all of which may be bonded to the metal.
Example:
Nonionic complexes are represented with a single entity.
Example: [Fe(NH3)3(CO)3] - Triamminetricarbonyliron(0).
A complex formed with two or more metal ions per molecule is called a polynuclear complex. The two metal ions in it are bridged by ligands; such bridging ligands are identified by the prefix µ - if there are two or more bridging ligands of the same type, this is indicated by tri -µ, di -µ. If there are two or more bridging ligands of different types, then the prefix µ - is added for each such bridging ligand in the alphabetical order.
Example:
In a complex having a metal-metal bond, the prefix bis or - di is used before the name of metal which involves metal-metal bond. Abbreviations for Polydentate ligands their formulas shown in Table 14.4.
[Br4ReReBr4]-2 bis(tetrabromorhenate)(ReRe) (2−)
Table 14.4 Abbreviation for ligand names
Some important complexes by Scientist names shown in Table 14.5.
Table 14.5 Scientist names for complexes
The various theories like Werner’s theory, Sidgwick’s electronic concept theory, valence bond theory (VBT), crystal field theory (CFT), molecule orbital theory (MOT), etc., have successfully explained the properties of complexes and the bonding between the metal and ligand.
In 1893, Alfred Werner proposed this theory at the age of 26; it is now referred to as Werner’s coordination theory. This theory explains the “primary-secondary valence” of metal complexes.
Example: Four coordinated complexes the four valances are arranged in either square planar or a tetrahedral structure and in six coordinated complexes, six valances are directed towards six corners of the octahedron.
Table 14.6 Werner studied chloroaminecobalt(III)complexes
In complex (A), primary valence is satisfied by three chloride ions and secondary valence is satisfied by 6 NH3 molecules. According to the complex structure, ammonia molecules are tightly bound to cobalt, hence they do not dissociate in solution, but chloride ions are far away from cobalt and are less firmly held by metal. Hence, all the three chloride ions dissociate in the solution, giving 3Cl− ions and [Co (NH3)6]3− with a total of four ions.
Werner’s theory describes the structure of many coordination compounds successfully; however, it does not explain the nature of bonding within the coordination sphere.
According to Sidgwick, the metal and the ligand involve a coordinate bond. Here, the ligand donates the electrons to the central metal atom to form coordination compounds. These ligands are known as donors and the metal ions are the acceptors. The bond between the donor and acceptor is called coordinate, or dative or semi-polar bond. The coordinate bond is generally represented by the “→” arrow, starting from the donor pointing towards the acceptor L→M.
Effective atomic number (EAN) is also known as 18 electron rule or Nobel gas rule. According to Sidgwick, the EAN of the central metal ion is the total number of electrons around the central metal ion, including those gained through coordination by the ligand. The EAN is equal to the atomic number of the next higher inert gas. Examples of complexes which obey EAN rule and which have exception from EAN rule are shown in Table 14.7 and 14.8 respectively.
Table 14.7 Examples of complexes which obey the EAN rule
Table 14.8 Examples of complexes which are exception to the EAN rule
Based on modern principle of bonding and overcome to drawbacks of Sidgwick’s and Werner’s theories, the valence bond theory, crystal field theory and molecule orbital theories are proposed.
In 1935, Linus Pauling and Slater proposed the valence bond theory (VBT). This theory is primarily concerned with the structure and magnetic behaviour of metal complexes. It is also called Pauling theory.
Examples:
A species should have at least one unpaired electron to possess paramagnetic nature. It is attracted by an external magnetic field; theoretically, the paramagnetic moment of an ion can be calculated by its number of unpaired electron using the following spin only formula.
When the species does not contain an unpaired electron, it is diamagnetic.
Depending on the number of d electrons in the metal ion, the VBT can be applied to guess the structure of the complex.
Number of d electrons = Atomic number of central metal atom - Number of electrons lost in ion formation - atomic number of previous inert gas.
Examples:
Ni atomic number = 28
Number of electrons lost in +2 ion formation = 2
Atomic number of previous inert gas = 18
Number of d electrons in Ni = 28 - 2 - 18 = 8
d8 configuration
Sp3 hybridisation - Tetrahedral structure
Two lone pair electrons - Paramagnetic
Observed magnetic moment = 2.83 BM.
dsp2 hybridisation: square planar
No lone pair of electrons: dia magnetic
Observed magnetic moment = 0 BM.
Co-atomic number = 27
Number of electrons lost in +3 ion formation = 3
Atomic number of previous inert gas = 18
Number of d electrons in Co = 27 - 3 - 18 = 6
d6 configuration
d2sp3 hybridisation: Octahedral
No lone pair of electrons: Dia magnetic
Observed magnetic moment = 0 BM.
The crystal field theory was first proposed by J. H Van Vleck and H. Bethe. This theory involves an electrostatic approach to bonding in complexes. It was first applied to ionic type crystalline substance. Therefore, it is called crystal field theory.
The five d-orbitals in isolated gaseous metal ions are degenerated; when a spherically symmetric field of negative charges is placed around the central metal ion, all the five orbitals will be raised in energy as a result of the repulsion between the negative field or negative electrons. Although they still remain degenerated, they will have higher energy than before.
t2g - triply degenerated orbitals are dxy, dyz and dzx.
eg - doubly degenerated orbitals are dz2 and dx2 − y2.
The difference between t2g and eg orbitals is the crystal field stabilisation energy; it is denoted by Δo or 10Dq. The magnitude of Δo depends on the field strength of the ligand as well as the metal ion. It is assumed that the sum of the energies of the five d-orbitals in the free state must be equal to the sum of the energies of five d-orbitals in the octahedral configuration. Splitting occurs in such a way that there is no net change in energy for the fully occupied orbital. If we consider the energy of the free state as zero, then
(1)
∴ (2)
(3)
Multiply equation (3) by two and subtract it from equation from (2).
∴
Since t2g levels are degenerated,
On substituting the value of Edxy, namely −0.4Δo in equation (3)
∴
Since eg orbitals are degenerate
This means that presence of an electron in the t2g level will stabilise the complex to the extent ‘0.4 Δo’. This stabilisation energyis called “crystal filed stabilisation energy”. The presence of electron in eg level will destabilise the complex to the extent of 0.6 Δo.
The metal ion along with its d-orbitals are arranged in the centre of the cube, the ligands occupy alternatively in four corners of the cube. Here, two ligands are in the top face of the cube and other two are in the bottom face but none of the ligand approaches the metal along the Cartesian coordinate axes. Hence, the orbitals along the axes (dz2, dx2 - y2) will be less strongly repelled than those orbitals between the axes (dxz, dxy and dyz). With this, the previously degenerate set of five orbitals is now split into two sets, one at a higher energy and one at lower energy. The higher energy set of orbitals (dxy, dyz and dzx) is labelled as t2 and the lower energy set (dz2 and dx2 - y2) is labelled as e. (The subscript “g” disappears here because the tetrahedron lacks the centre of inversion). The energy separation between the two levels is denoted by Δt (or) 10 Dq.
The crystal field splitting in the tetrahedral field is intrinsically smaller than the octahedral field because there are only four ligands surrounded to the metal, but none of them have a direct effect on the d-orbital. The relationship between the two crystal field separation energies may be represented as Δt = 4/9Δo.
The tetrahedral geometry can occur with the following reasons:
Example:
This is the special case of octahedral symmetry; the removal of two ligands from z-direction completely in octahedral symmetry leads to square planar geometry. The geometry is favoured by the metal ion having a d8 configuration in the presence of a strong field. This combination gives low spin complexes where the first four orbitals are occupied and the high energy dx2 − y2 orbitals are unoccupied.
According to field strength, the single atomic ligands are arranged in Table 14.9 from weak field to strong field. When weak ligand approaches the metal, outer orbitals of the metal are involved in hybridization, that is, high spin complexes (6 coordinated complexes –sp3d2, 4 coordinated complexes –sp3). When strong ligand approaches the metal, inner orbitals of the metal are involved in hybridization, that is, low spin complexes (6 coordinated complexes –d2sp3, 4 coordinated complexes –dsp2).
Table 14.9 Single atomic ligands and their bonding capacity from weak field to strong field
The entries in the table are sorted by field strength directly binding through the stated atom, that is, as a terminal ligand. The ‘strength’ of the ligand may change when the ligand binds in an alternative binding mode as follows:
Molecular orbital theory was proposed by Pauling; it explains mainly magnetic properties, bond strength or bond order, bond length and the covalent character of metal and ligand.
For example, [Co(NH3)6]+3 is a complex in which metal and ligand α-bonding are involved.
This complex contains NH3 ligand which can only participate as a α donor to the metal ion. Now, construct the molecular orbital diagram for octahedral complex.
Let us consider the metal orbitals. The valance orbitals available are 3d, 4s and 4p. A total of nine orbitals are available and these orbitals are suitable to be used in α bonding. This means that the orbitals that we can use should have their lobes pointed along axes. Obviously, 4s orbitals are quite suitable. This called a1g. All 4p orbitals are also suitable. These are labelled as t1u. Now, let us inspect the 3d orbitals; only two of them are suitable. The orbitals dx2 − y2 and dz2 are suitable since they have the lobes along the axes. These are labelled as eg; however, the orbitals dxz, dyz and dxy are not suitable. Since the ligand orbitals cannot overlap with them to give a positive overlap, this cannot be utilised by the metal for α bonding. These remain non-bonding and are labelled as t2g.
Thus, a valance orbital of the metal six is suitable for forming σ-bonding (a1g, t1u, eg); three are non-bonding orbitals (t2g).
Let us assume that by linear combination of the ligand orbitals, we are able to generate six orbitals of symmetry, called ligand group orbital as that of the six metal bonding orbitals. These orbitals on the metal and the ligand combine.
Thus, [CO (NH3)6]3+ [CO has 6 electrons, 6 NH3 ligands have 12 electrons] thus electronic configuration . In contrast, [CO F6]3− has the electronic configuration
MO diagram for octahedral complex
Metal complex ligand
It is not possible to conclude a particular ligand or particular metal complexes of high or low stability. No single factor is expected to account for the relative stability of coordination compounds. They are affected many physical and chemical properties as discussed here.
The complex formation is essential a reaction between cation (metal ion) and ligand. The charge on metal ion is more important to deciding the stability of metal complex. Greater the positive oxidation state of the central metal ion greater will be the attraction for the ligand; hence, greater will be the stability of the complex with same ligand. When metal ion is in a more oxidation state, that complex is more stable when compared to low oxidation metal complex.
Keeping the charge constant as the size of the metal ion decreases, the specific charge per unit surface area increases. Hence, the metal attraction for the ligand increases. Generally, the size of the metal decreases and the stability of metal complex increases.
We compare the stability of complexes formed by similarly charged and sized metal ions belonging to two subgroups of the same periodic family. This stability is explained by outer electronic configuration.
Exmaple: K+ and Cu+ potassium are non-transition elements, whereas copper is a transition metal. Cations of d-block elements are invariably far more stable complex than non-transition metal complexes.
Potassium has inert gas configuration (n - 1) s2p6 in the outer shell, whereas copper has pseudo-inert gas configuration (n - 1) s2p6d10 in the outermost shell. It is a well-known fact that the latter configuration is much poorer in shielding the excess positive charge is located on the nucleus then the formed configuration. Hence effective nuclear charge on copper is more than potassium, that is copper will behave as a greater nuclear charge. Therefore, Cu+ has greater attraction for electrons offered by the ligands. Consequently, copper form more stable complex than potassium
The complexes formed by halide ion have been widely studied for metals the order of stabilities as it follows the sequence F−>Cl−>Br−>I−. However, this order is reversed for a few metals like Pt+2, Cu+, Ag+, Hg+2 and Ti+ where back donation occurs from metal to ligand occur in the donation to the transfer of electron from ligand to the metal. For this back donation from metal to ligand to occur, the ligand must possess a vacant orbital capable of receiving electrons.
Basicity is a measure of electron pair donation; greater the basicity of ligand, greater will be the tendency to donate electron pairs. This means that a more basic ligand will form more stable complex.
The stabilities of complexes are increased by the coordination of the ploy dentate ligand. Coordination of such ligands produces a ring structure with the metal atom forming a part of the ring system. Such ring structure complexes are called chelates. This process of chelate formation is called chelation. The ligand that forms a chelate is called a chelating agent. Due to chelation, extra stability is conferred on the complex. This extra stability is termed as the chelating effect.
Greater the number of chelate rings, greater will be the stability of complex. Some other factors are as follows:
Many methods are used to confirm the formation of complex compounds in a reaction as discussed here.
When a neural covalent complex is formed in a polar medium, it gets precipitated. The formation of the precipitate on the addition of a ligand indicates the formation of a neural covalent complex.
Example:
The altered reactivity of the metal ions in the presence of the ligands is still a useful indication of complex formation.
Example: Fe+3 is not precipitated as Fe(OH)3 by NH4OH in the presence of tartaric acid. This indicates the formation of Fe3+ tartrate complex.
Many water insoluble substances dissolve readily in the presence of a reagent that forms a soluble complex.
Example: Dissolution of AgCl in aqueous ammonia indicates the formation of soluble complex [Ag(NH3)2]Cl.
A colour change of the solution with the addition of another reagent indicates the formation of a new species. In many cases, the colour of the uncomplex metal ion gets altered and in some cases, a totally new colour is produced.
The addition of the ligand to metal in solution complex is formed by the releasing of protons.
Example: Cu+2 ion is treated with salicylaldoxime to form the complex [Cu(sal)2] with the release of protons.
If the complex formation reaction involves a change in the number of ions, then there would be changes in the conductivity of the solution. The change in conductance can be used to detect the complex formation.
The subject of stability of metal complexes is important in understanding the properties of complexes. Many variables associated with the central metal ion and the ligand are greatly complicated. There are two types of stability—thermodynamic stability and kinetic stability.
Thermodynamic Stability
This measures the extent of the complex formed or transformed into another species by metal ligand bond energy stability constant when system reached equilibrium.
Kinetic Stability
This refers to the speed of the reaction with transformation, leading to attainment of equilibrium.
The thermodynamic stability of a complex depends upon the difference in energy between the reactants and the products; greater the reaction energy, greater the thermodynamic stability.
The kinetic stability of the complex depends on the difference in energy between the reactants and the activated complex, that is, activation energy. Greater the activation energy, lesser the reaction rate, implying that the complex is inert.
If the term “stability” is used without any modification, then it refers to thermodynamic stability.
In thermodynamic point of view, metal-to-ligand bond energies, stability constant and the thermodynamic variables derivable from them from being stable or unstable have to be considered.
In kinetic stability, the rates and the mechanism of reactions and also with the energies involved in the formation of the activated complex have to be considered. In the kinetic point of view, it will be more proper to speak of complex as being inert or labile rather than stable or unstable. Very often, these two groups of terms are used in correctly. A stable complex may be inert or labile and an unstable complex may be labile or inert.
Example: CN−ion forms a very stable complex with Ni+2
Ni+2 prefers CN− rather than H2O as a ligand. Thus, [Ni(CN)4]2− is thermodynamically more stable than [Ni(H2O)6]2+. However, when “14C-labelled” CN− (14CN−)is added to the solution, it is almost instantaneously incorporated into complex.
This means that the complex, [Ni(CN)4]2− is kinetically labile. Thus the stability of complex of this does not ensure inertness.
This method is used for solution where only one complex is formed.
The sum of the total concentration C of a complexing agent Cx and metal ion CM is held constant and only their ratio is varied.
A wavelength of light is selected such that where the complex absorbs strongly with metal and ligand but not individually.
A plot of the mole fraction of the ligand in the mixture x versus absorbance gives a triangular curve.
The mole fraction of the ligand in the mixture x = Cx/C
The mole fraction of the metal in the mixture = CM/C
The legs of the triangle are extrapolated until they cross. The mole fraction at the point of this intersection gives the formula of complex, since at this point, for the complex Mxn is
At this point, the ligand and the metal are in the proper relative concentration to give maximum complex formation. Performing the experiment at several different wavelengths at several different values of C would indicate whether there is more than one complex formed in the solution. In such cases, n would not be constant.
The curve may be observed to deviate somewhat from the experiment intersecting lines from the amount of derivation and the stability of constant of the complex can be determined.
For 1:1 complex, the ratio of the true absorbance (A) to the extrapolated absorbance (AExt) is the mole fraction of the complex actually formed.
Where C = Total concentration of the metal (or) ligand, whichever is the limiting concentration at the point considered.
Where K is the stability constant and CM and Cx are the total concentration of metal and ligand respectively.
Coordination compounds are used in several areas of chemistry such as analytical, medicinal, industrial chemistry and agriculture.
Inorganic Qualitative Analysis
Coordination chemistry is used in inorganic qualitative analysis. The formation of metal complex is employed in the separation and identification of some metals.
Separation
The principle of masking is used in separationof some metals from each other on qualitative analysis.
Example: Cu+2 and Cd2+ forms insoluble sulphide in the II group. These two metals are separated by complexing them with CN−. Both form complexes but with a difference.
The reaction of CN− with CuS is an oxidation reduction reaction coupled with complexation. In both complexes, Cu complex is stable and Cd complex is unstable. The marked difference in the stabilities of copper and cadmium cyano complexes is the basis for separating these two metals.
Reductive Complexation
When a ligand reduces a metal of the complexes, the process is called reductive complexation.
Identification of Metals
Complex formation is used to identify several metals in qualitative analysis.
Complexometric Titration
Certain metal ion reacts with certain ligand stoichiometrically and forms stable metal complexes. This reaction can be used to in volumetric estimation of metal ions.
Example: determination of hardness of water.
Hard water is titrated with EDTA ligand; water containing metals Ca and Mg will form complex with EDTA. From the end point, hardness giving substances present in water can be determine.
Complex in Calorimeter
Some of the metals can be calorimetrically estimated by the formation of coloured complex species in solution.
Coordination Compounds in Gravimetry
The formation of coordination compounds serves as an excellent means for gravimetrically estimating certain metals.
Complexes in Separation of Metals
Some chelating agents are used for separating certain metallic mixtures.
[Ans.: Complexation]
[Ans.: Double salt]
[Ans.: Coordination number]
[Ans.: +2]
[Ans.: Unidentate]
[Ans.: Bridging]
[Ans.: Complex anion]
[Ans.: Secondary]
[Ans.: Coordinate bond or dative bond]
[Ans.: 31]
[Ans.: No]
[Ans.: Sidgwick]
[Ans.: Linus Pauling and Slater]
[Ans.: Hybridisation]
[Ans.: dsp2]
[Ans.: Linear]
[Ans.: Valence bond]
[Ans.: Crystal field theory]
[Ans.: Electrostatic attraction]
[Ans.: t2g]
[Ans.: dxy, dyz, dzy, t2g]
[Ans.: dx2−y2 and dz2]
[Ans.: Δ0 (or) 10Dq, crystal field splitting]
[Ans.: t2g]
[Ans.: b]
[Ans.: a]
[Ans.: b]
[Ans.: d]
[Ans.: c]
[Ans.: a]
[Ans.: a]
[Ans.: b]
[Ans.: a]
[Ans.: a]
[Ans.: a]
[Ans.: c]
[Ans.: a]
[Ans.: a]
[Ans.: a]
[Ans.: b]
[Ans.: b]
Ans.: Compound of metals with certain number of species called ligands bound to the metal is called coordination compound.
Ans.: Double salt:
Double salt is formed by the adding of two stable compounds; on dissolution, it breaks down to its simple component ions.
Coordination compounds:
On dissolution, it does not break down into simple ions.
Ans.: The total number of ligands attached to metal complex is called coordination number.
The charge on the complex will be denotes the oxidation state.
Ans.: (i) Monodentate
(ii) Polydentate
(iii) Bridging ligand
(iv) Ambidentate ligand
(v) Flexidentate ligand
Ans.: (i) Hexaaminechromium(III)ion.
(ii) Dichloroaquahydrazoniumplatinum(II)ion.
(iii) Trichlorotrinitrochromate(III)ion.
(iv) Biethylenediaminecopper(II)sulphate.
Ans.: (i) Werner’s theory
(ii) Sidgwick theory
(iii) Valence bond theory
(iv) Crystal field theory
(v) Molecular orbital theory
Ans.: It explains the elements exhibit two types of valence as follows:
Drawbacks:
It is unable to account the 4-coordinated and 6-coordinated complex in coordination chemistry.
It does not explain the nature of bonding between metal and ligand.
Ans.: According to Sidgwick, the total electrons around the metal ion including those gained through the coordination by the ligand is called effective atomic number (EAN).
EAN is equal to the atomic number of the next inert gas atomic number such that metal complexes are stable.
Ans.: In 1935, Pauling and Slater proposed this theory. This theory explains the magnetic properties and structure of the metal complexes.
Ans.: This theory was explained by J.H. Van Vleck and Bethe. It was first applied by crystalline substances and on this basis, it was called crystal field theory.
Merits:
Demerits:
Bond strength chemical properties cannot be explained.
Ans.: (i) Charge of the metal complex: Greater the charge on metal greater will be the stability of complexes.
(ii) Size of the metal ion: Size decreases the stability of metal complexes increases.
Ans.: (i) Formation of precipitate
(ii) Change in chemical behaviour
(iii) Spectral method
(iv) pH method
(v) Conductivity measurements
Q.1 Write a note on the characteristics of coordination compound.
Q.2 Explain the difference between double salt and coordination compound.
Q.3 Explain the following terms:
Q.4 Define ligand. Explain the types of ligands with examples.
Q.5 Describe the nomenclature of coordination compounds in detail.
Q.6 Write a detailed note on any two theories of coordination chemistry?
Q.7 Write a short note on crystal field stabilisation energy and splitting of d-orbital in crystal field theory.
Q.8 Explain the factors affecting coordination compounds.
Q.9 Give a brief explanation on the detection of complex ion formation.
Q.10 Explain stability of coordination compounds.
Q.11 Describe Job’s method.
Q.12 Give the application of coordination complexes in analytical chemistry.
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