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Coordination compounds.pptx..............................

Coordination compounds are chemical complexes where a central metal atom or ion is surrounded by molecules or ions called ligands. These ligands are bonded to the central metal through coordinate covalent bonds, where the ligand donates a pair of electrons. They play a crucial role in various chemical, biological, and industrial processes, with examples including hemoglobin, chlorophyll, and catalysts

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Coordination compounds • Transition metals form large number of complex compounds in which there is a central metal atom bound to number of anions or neutral molecules. • Chlorophyll, haemoglobin, and vitamin B12 are coordination compounds of Magnesium, iron and cobalt respectively.
Werner’s theory of coordination compounds • In coordination compounds metals show two types of linkages(valences) – primary and secondary. • The primary valences are normally ionisable and are satisfied by negative ions. • The secondary valences are non ionisable. These are satisfied by neutral molecules or negative ions. The secondary valence is equal to the coordination number and is fixed for a metal. • The groups bound by secondary linkages to the metal have characteristic spatial arrangements corresponding to different coordination numbers.
Explanation to Werner’s theory • Werner proposed the concept of primary valence and secondary valence for a metal atom. • Binary compounds such as CrCl3, CoCl2, PdCl2 have primary valence of 3, 2, and 2 respectively. • Cobalt(111) chloride reacts with ammonia to form 4 compounds 1. CoCl36NH3(yellow)- gave 3 moles of AgCl with AgNO3 solution. 2. CoCl35NH3(purple)- gave 2 moles of AgCl with AgNO3 solution. 3. CoCl34NH3(yellow)- gave 1 mole of AgCl with AgNO3 solution. 4. CoCl34NH3(violet)- gave 1 moles of AgCl with AgNO3 solution. • This can be explained if 6 groups in all remain bonded to Cobalt atom which cannot be ionised. Werner proposed the term secondary valence for the number of groups directly attached to the metal atom.
Problem • Assign secondary valences to the metals 1. PdCl2.4NH3-2 moles of AgCl is precipitated with excess AgNO3. 2. NiCl2.6 H2O - 2 moles of AgCl is precipitated with excess AgNO3. 3. PtCl4.2HCl - 0 moles of AgCl is precipitated with excess AgNO3. 4. CoCl3.4NH3- 1 moles of AgCl is precipitated with excess AgNO3. 5. PtCl2.2NH3- 0 moles of AgCl is precipitated with excess AgNO3.
Comparison between double salt and a complex • Both double salts and complexes are formed by the combination of two or more stable compounds in stoichiometric ratio. • Eg for double salt- carnallite - KCl.MgCl2.6H2O Mohr’s salt- FeSO4.(NH4)2SO4.6H2O • Example for complex-K4[Fe(CN)6] [Cu(NH3)4]SO4 • Double salts dissociate into simple ions completely when dissolved in water. • The complex part of the complex salt does not dissociate.
Some important terms pertaining to coordination compounds 1. Coordination entity: It constitutes of a central metal atom or ion bonded to a fixed number of ions or molecules. Eg: [Ni(CO)4], [PtCl2(NH3)2], [Fe(CN)6]4-+ 2. Central atom/ion -The atom/ion to which a fixed number of groups are bound in a definite geometrical arrangement. The central atoms/ions are Lewis acids.
3. ligands • The ions or molecules bound to the central atom in the coordination entity are called ligands. • When a ligand is bound to a metal ion through a single donor atom it is said to be unidentate. Eg: Cl- , H2O, NH3. • When a ligand can bind to the metal atom through two donor atoms it is said to be didentate. Eg: NH2CH2CH2NH2(ethane-1,2 diamine) or C2O4 2- (oxalate) • When several donor atoms are present in the same ligand the ligand is said to be polydentate. Eg: EDTA4- • When a di or polydentate ligand uses its two or more donor atoms to bind a single metal atom/ion it is said to be a chelate ligand • Ligand which can ligate through two different atoms is called ambidentate ligand. Eg: NO2 - , SCN-
4. Coordination number • The coordination number of a metal atom/ion in a complex is the number of ligand donor atoms to which the metal is directly bonded. • CN of Pt in [PtCl6]2- is 6 • CN of Ni in [Ni(NH3)4]2+ is 4 • CN of Fe in [Fe(C2O4)3]3- is 6 • CN of Co in [Co(en)3]3+ is 6
5. Coordination sphere • The central atom/ ion and the ligands attached to it are enclosed in a square bracket and is collectively called as coordination sphere. • The ionisable groups written outside the bracket are called counter ions. • For example: In the complex K4[Fe(CN)6], the coordination sphere is [Fe(CN)6]4- and the counter ion is K+ .
6. Coordination polyhedron • The spatial arrangement of the ligand atoms which are directly attached to the central atom/ion is the coordination polyhedron. • The most important coordination polyhedras are octahedral, square planar and tetrahedral.
7. Oxidation number of 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 that are shared with the central atom. • Find the oxidation number of the central metal atom/ion. 1. [Cu(CN)4]3- 2. [Co(NH3)6]3+ 3. K4[Fe(CN)6]
Homoleptic and heteroleptic complexes • Homoleptic complexes: The complexes in which a metal is bound to only one kind of donor groups. Eg: [Co(NH3)6]3+ • Heteroleptic complexes: Complexes in which metal is bound to more than one kind of donor groups. Eg:[Co(NH3)4Cl2]+
Nomenclature of Coordination Compounds • Rules for nomenclature 1. Cation is named first 2. Ligands are named in the alphabetical order before the name of the central metal atom/ion. 3. Names of anionic ligands end with –o. 4. Prefixes mono,di, tri etc are used to indicate the number of individual ligands. When the name of the ligand has a numerical prefix terms bis, tris, and tetrakis are used. 5. Oxidation number of the metal is indicated by Roman numeral in parenthesis. 6. If the complex ion is anion the name of the metal ends with suffix –ate.
Examples: write the IUPAC names of the following complexes. 1. [Cr(NH3)3(H2O)3]Cl3 Ans: triamminetriaquachromium(III) chloride 2. K2[Zn(OH)4] Ans: potassium tetrahydroxidozincate(II) 3. [Ni(CO)4] Ans: tetracarbonylnickel(0) 4. [CoCl2(en)2]+ Ans: dichloridobis(ethane-1,2-diamine)cobalt(III) 5. [Co(NH3)4(H2O)Cl]Cl2 Ans: tetraammineaquachloridocobalt(III)chloride
Rules for writing the formula of coordination compounds • The central atom is listed first • Ligands are then listed in the alphabetical order. • The formula of the entire coordination entity is enclosed in square brackets. When ligands are polyatomic, their formulas are enclosed in parenthesis. • There should be no space between the ligands and the metal within the co- ordination sphere. • When the formula of charged coordination entity is to be written without that of counter ion, the charge is indicated outside the square bracket as a right superscript, with number before sign. • The charge of the cation is balanced by the charge of the anions.
Write the formulas of the following coordination complexes 1. Potassium trioxalatochromate(III) Ans: K3[Cr(C2O4)3] 2. dichloridobis(ethane-1,2-diamine)cobalt(III) chloride Ans: [CoCl2(en)2]Cl 3. pentaamminecarbonatocobalt(III) chloride Ans: [Co(NH3)3(CO3)]Cl 4. Potassium tetracyanidonickelate(II) Ans: K2[Ni(CN)4] 5. amminebromidochloridonitrito-N-platinate(II) Ans: [Pt(NH3)Br Cl NO2]-
Isomerism in coordination compounds • The types of isomerism known among coordination compounds are 1. Stereo isomerism: stereo isomerism can be subdivided into geometrical isomerism and optical isomerism 2. Structural isomerism: Structural isomerism can be subdivided into 1. Linkage isomerism 2. Coordination isomerism 3. Ionisation isomerism 4. Solvate isomerism
Geometrical isomerism • Geometrical isomerism arises in heteroleptic complexes due to the different possible geometrical arrangements of ligands. It is common in coordination complexes with coordination numbers 4 and 6. • Geometrical isomerism in square planar complexes with C.N =4 • In square planar complex with formula [MX2L2] if two identical ligands are arranged adjacent to each other it is a cis isomer, or opposite to each other it is a trans isomer. Geometrical isomers of [Pt(NH3)2Cl2]
Geometrical isomerism(contd) Geometrical isomerism in octahedral complexes with C.N = 6 type: [MX2L4] • Two isomers are possible-cis and trans
Geometrical isomerism(contd) Geometrical isomerism in octahedral complexes with C.N = 6 type: [MX2(L-L)2]
Geometrical isomerism in octahedral complexes with C.N = 6 type: [Ma3b3] • Two types of isomers are possible-facial and meridonial Geometrical isomerism(contd)
Optical isomerism in coordination compounds • Optical isomers are mirror images that cannot be superimposed on another. These are called enantiomers. The two forms of optical isomers are dextro (d) and leavo (l) depending on the direction they rotate the plane of polarised light. Optical isomerism in [Co(en)3]3+ Optical isomerism in [PtCl2(en)2]3+ (only cis form shows optical activity)
Linkage isomerism • Linkage isomerism arises when an ambidentate ligand is present in the coordination complex. In linkage isomers different atoms of the same ambidentate ligand attach to the central metal atom. • Example: [Co(NH3)5(NO2)]Cl2 and [Co(NH3)5(ONO)]Cl2 are linkage isomers
Coordination isomerism • Coordination isomerism arises when both cation and anion are coordination entities. • In coordination isomerism interchange of ligands take place between cationic and anionic entities of different metal ions present in a complex. • Example: [Co(NH3)6][Cr(CN)6] and [Cr(NH3)6][Co(CN)6] are coordination isomers.
Ionisation isomerism • Ionisation isomerism arises when the counter ion in a complex salt is itself a potential ligand. • Ionisations isomers have the same molecular formula but ionises to form different ions in solution. • Example: [Co(NH3)5SO4]Br and [Co(NH3)5Br]SO4 are ionisation isomers.
Solvate isomerism • Solvate isomerism is known as hydrate isomerism when water is the solvent. • In Solvate isomers differ by whether a solvent molecule is directly bonded to the metal ion or is present as solvent molecules in crystal lattice. • Example: [Cr(H2O)6]Cl3 and [Cr(H2O)5Cl] Cl2 .H2O
Bonding in coordination compounds Valance bond theory • According to valence bond theory, the metal atom or ion in the presence of ligands use its atomic orbitals for hybridisation. • Hybridisation results in a set of equivalent orbitals with definite geometry. Eg: octahedral, tetrahedral, square planar etc. • These hybrid orbitals overlap with ligand orbitals which donate electron pairs for bonding. CN hybridisation geometry 4 sp3 tetrahedral 4 dsp2 Square planar 5 sp3 d Trigonal bipiramidal 6 sp3 d2 octahedral 6 d2 sp3 octahedral
Bonding in [Co(NH3)6]3+ Orbitals of Co3+ - 3d 4s 4p d2 sp3 hybrid orbitals of Co3+ - d2 sp3 hybrid orbitals [Co(NH3)6]3+ - 6 pairs of electrons from six NH3 molecules • Geometry - octahedral • Magnetic property -diamagnetic due to the absence of unpaired electrons • The inner d orbitals are used for hybridisation ,so low spin complex ↑↓ ↑ ↑ ↑ ↑ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓
Bonding in [CoF6 ]3- Orbitals of Co3+ - 3d 4s 4p 4d sp3 d2 hybrid orbitals of Co3+ - sp3 d2 hybrid orbitals [CoF6]3- 6 pairs of electrons from ligands • Geometry : octahedral • Magnetic property : paramagnetic due to the presence of unpaired e- • Since outer d orbitals are used for hybridisation high spin complex ↑↓ ↑ ↑ ↑ ↑ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓
Bonding in [NiCl4]2- Orbitals of Ni2+ - 3d 4s 4p sp3 hybrid orbitals of Ni2+ sp3 hybrid orbitals [NiCl4]2- 4 pairs of electrons from 4 Cl- • Geometry : tetrahedral • Magnetic property : paramagnetic due to the presence of unpaired e- ↑↓ ↑↓ ↑↓ ↑ ↑ ↑↓ ↑↓ ↑↓ ↑↓
Bonding in [Ni(CN)4]2- Orbitals of Ni2+ - 3d 4s 4p dsp2 hybrid orbitals of Ni2+ dsp2 hybrid orbitals [Ni(CN)4]2- 4 pairs of electrons from 4 CN- • Geometry : square planar • Magnetic property : diamagnetic due to the absence of unpaired e- ↑↓ ↑↓ ↑↓ ↑↓
Magnetic properties of coordination compounds • The magnetic moment of coordination compounds can be measured by magnetic susceptibility experiments. The results can be used to find the structures of complexes. • For metal ions with d1 , d2 and d3 configuration two d orbitals are available for octahedral hybridisation. • For metals with d4 , d5 and d6 pairing should take place to make two d orbitals available. • If pairing take place 2,1,and 0 unpaired electrons are left. If pairing of electrons take place inner complex is formed. If pairing does not take place outer complex is formed utilising outer d orbitals. • Whether a complex is inner complex or outer complex can be understood by magnetic susceptibility experiment.
Magnetic properties of coordination compounds (contd) Examples: 1. [Mn(CN)6]3- Central metal ion – Mn3+ electronic configuration – 3d4 , From magnetic susceptibility experiments 2 unpaired electrons are present. Suggest that it is an inner complex 2. [MnCl6]3- Central metal ion – Mn3+ electronic configuration – 3d4 , From magnetic susceptibility experiments 4 unpaired electrons are present. Suggest that it is an outer complex
Magnetic properties of coordination compounds (contd) 3. [Fe(CN)6]3- Central metal ion – Fe3+ electronic configuration – 3d5 , From magnetic susceptibility experiments 1 unpaired electron is present. Suggest that it is an inner complex 4. [FeF6]3- Central metal ion – Fe3+ electronic configuration – 3d5 , From magnetic susceptibility experiments 5 unpaired electrons are present. Suggest that it is an outer complex
Magnetic properties of coordination compounds (contd) 5. [Co(ox)3]3- Central metal ion – Co3+ electronic configuration – 3d6 , From magnetic susceptibility experiments 0 unpaired electron is present. Suggest that it is an inner complex 6. [CoF6]3- Central metal ion – Co3+ electronic configuration – 3d6 , From magnetic susceptibility experiments 4 unpaired electrons are present. Suggest that it is an outer complex
Limitations of valence bond theory • It involves many assumptions. • It does not explain the colour exhibited by coordination compounds. • It does not predict whether the 4 coordinate complexes are tetrahedral or square planar. • It does not distinguish between week and strong ligands
Crystal field theory • In crystal field theory the metal- ligand bond is ionic in nature. • Ligands are treated as point charges in case of anions or dipoles in the case of neutral molecules. • The five d orbitals in an isolated gaseous atom is degenerate. • This degeneracy is maintained in a spherically symmetrical field of negative charges around the metal atom. • In the presence of ligands in a particular field, this degeneracy is lost- splitting of d orbitals take place. • The pattern of splitting depends on the nature of the crystal field.
Crystal field splitting in octahedral coordination entities • The splitting of the degenerate d orbitals due to the presence of ligands in a definite geometry is termed as crystal field splitting. • When ligand approach metal ion in an octahedral field, there is repulsion between the electrons in metal d orbitals and electrons of ligands. • and orbitals which point towards the axis experience more repulsion and is raised in energy, where as dxy, dyz and dxz orbitals which are directed between the axis will be lowered in energy.
d- orbital splitting in octahedral field
d- orbital splitting in octahedral field • The crystal field splitting ∆0 depends upon the field produced by the ligand and charge on the metal ion. • Some ligands are able to produce strong fields. Then splitting produced will be large. Weak field ligands produce small splitting. • Ligands can be arranged in a series in the order of increasing field strength as given below I- <Br- < SCN- < Cl- < F- <OH- < C2O4 2- < H2O < NCS- <edta4- < NH3< en< CN- <CO. This series is called spectro chemical series. • When 1, 2 or 3 electrons are present in d orbital they will occupy t2g orbitals which are lower in energy. • The forth electron can go to either t2g or eg orbital. Which of this occurs is decided by crystal field splitting (∆0) and the pairing energy P
d- orbital splitting in octahedral field (contd) • If ∆0 < P, the forth electrons enters eg orbital. Ligands for which ∆0<P, are known as weak field ligands. They form high spin complexes. • If ∆0 > P, the forth electrons enters t2g orbital. Ligands for which ∆0>P, are known as strong field ligands. They form low spin complexes.
Crystal field splitting in tetrahedral coordination entities • In tetrahedral coordination entity formation, the d orbital splitting is inverted and is smaller as compared to octahedral field splitting.
Colour in coordination compounds • Coordination complexes are coloured because they absorb radiation from the visible range of spectrum. • The colour of the complex is complimentary to that of the colour absorbed. • For example: If light corresponding to blue green region is absorbed, the complex appears to be violet. • The energy of the absorbed radiation is used for the excitation of the electron from t2g level to eg level. So according to crystal field theory, the colour of the coordination compounds is due to d-d transition of electrons. • In the absence of ligands, crystal field splitting does not occur and the substance is colourless. For eg: removal of H2O from [Ti(H2O)6]Cl3 on heating makes it colourless. Anhydrous CuSO4 is white but CuSO4.5H2O is blue in colour.
Merits and demerits of Crystal field theory • Crystal field theory successfully explains the formation, structure, colour and magnetic property of coordination complexes. • Crystal field theory assumes ligands as point charges. According to this assumption, anionic ligands should produce more splitting. But anionic ligands are found at the low end of the spectrochemical series which cannot be explained. • It does not take into account the covalent character of bonding.
Bonding in metal carbonyls • The metal –carbon bond in metal carbonyls possess both σ and π character. • The metal- carbon sigma bond is formed by the donation of lone pair of electrons on the carbonyl carbon into the vacant orbital of the metal. • The M-C pi bond is formed by the donation of a pair of electrons from a filled d orbital of metal into the vacant antibonding π* orbital of carbon monoxide.
Stability of coordination compounds • The stability of a complex in solution refers to the degree of association between the two species involved in the state of equilibrium. • Consider the reaction M + 4L == ML4. Larger the stability constant, higher the proportion of ML4. • The instability constant or the dissociation constant is the reciprocal of formation constant.
Importance of coordination compounds • Coordination compounds find use in many quantitative and qualitative analysis. Eg: estimation od hardness of water usinf EDTA. • Extraction of some metals like gold. • Purification of metals Eg: Ni • Coordination compounds are of great importance in biological systems. • Coordination compounds are used as catalysts for many industrial processes.
Question from previous Q.P. (2 three mark questions) 1. Explain the hybridisation, geometry and magnetic property of [Co(NH3)6]3+ ion on the basis of valence bond theory. [Atomic number of cobalt- 27] 2. a. For a given complex [Co(NH3)5NO2]Cl2. Write its IUPAC name and linkage isomer. b. Which set of d-orbitals of a metal atom/ion experience more repulsion in octahedral field created by the ligands? 3. [Ni(CN)4]2- 4. a. Write the cis and trans isomeric structures of [Fe(NH3)2(CN)4]- b. What is the coordination number of Fe in [FeCl2(en)2]Cl
5. [Ni(CN)4]2- 6. a. Write the IUPAC name of K3[Cr(C2O4)3] b. Give the facial and meridonial isomeric structures of [Co(NH3)3(NO2)3] 7. [NiCl4]2- 8.a. Write the IUPAC name of [Co(NH3)4(H2O)Cl]Cl2 b. Explain linkage isomerism with example. 9. [Ni(CN)4]2- 10. a. Explain the crystal field splitting of d- orbitals in octahedral complexes. b. Write the formula of tetraammineaquachloridocobalt(III) chloride
11. [Ni(CN)4]2- 12. a. For the given complex [Co(NH3)5Br]SO4, write the IUPAC name and its ionisation isomer. b. Which set of d- orbitals of metal ion/atom experience more repulsion in octahedral field created by ligand? 13. [Ni(CN)4]2-+ 14. Give the IUPAC name of [Ti(H2O)6]3+ . Draw cis and trans isomers of [Pt(NH3)2Cl2] 15. [CoF6]3- 16. a. Write any two postulates of Werner’s theory of coordination compounds. b . Write the IUPAC name of [Pt(NH3)3(H2O)Cl2]