Coordination Compounds PDF

# Chapter-9. Co-ordination Compounds ## Coordination Compounds: (Complex Compounds) - Coordination compounds are the compounds in which the central atom is linked to a number of ions or neutral molecules by co-ordinate bonds. ## Examples of Coordination Compounds in Biological Systems: - Chloroph...

# Chapter-9. Co-ordination Compounds ## Coordination Compounds: (Complex Compounds) - Coordination compounds are the compounds in which the central atom is linked to a number of ions or neutral molecules by co-ordinate bonds. ## Examples of Coordination Compounds in Biological Systems: - Chlorophyll, haemoglobin and vitamin B12 are coordination compounds of Mg, Fe and Co respectively. ## Differentiate Between a Double Salt And a Complex Compound: 1. **Double Salts** - Formed by the combination of two or more stable compounds in stoichiometric ratio. - They are stable in a solid state, but dissociate into constituent ions completely when dissolved in water. - **Potash alum:** $K_2SO_4.Al_2(SO_4)_3.24H_2O$ dissolves in water as $K^+, SO_4^{2-}, Al^{3+}$ and $H_2O$. - **Mohr salt:** $FeSO_4(NH_4)_2SO_4. 6H_2O$ dissolves as $Fe^{2+}, 2NH_4^+, 2SO_4^{2-}$ and $H_2O$. - Double salts are ionic compounds and do not contain any co-ordinate bond. - Properties of double salts are same as those of the constituent compounds. - A double salt loses its identity in solutions. 2. **Complexes** - Are also formed by the combination of two or more stable compounds in stoichiometric ratio. - Complex compounds do not dissociate when dissolved in $H_2O$. - $K_4[Fe(CN)_6]$ does not dissociate to $4K^+$ and $[Fe(CN)_6]^{4-}$ in solution. - The bond between the ligand and the central metal atom/ion in a complex is a coordinate bond. - The properties of complex compounds are different from its constituents. ## Nomenclature of Complex Compounds - **Rules:** Refer to Text book. - **Ligands:** | Ligand | Name | Charge | Density | Ligand | Name | Charge | Density | |---|---|---|---|---|---|---|---| | $CN$ | Cyanido | -1 | 1 | $H$ | Hydrido | -1 | 1 | | $Cl$ | Chloro | -1 | 1 | $OH$ | Hydroxido | -1 | 1 | | $Br$ | Bromo | -1 | 1 | $SO_3^{2-}$ | Sulphito | -2 | 1 | | $I$ | Iodo | -1 | 1 | $SO_4^{2-}$ | Sulphato | -2 | 1 | | $NO_3$ | Nitrato | -1 | 1 | $CO_3^{2-}$ | Carbonato | -2 | 1 | | $CH_3COO$ | Acetato | -1 | 1 | $COO^{2-}$ | Oxalato | -2 | 2 | | $NO_2$ | Nitrato-N | -1 | 1 | $(CO_2)_2^{4-}$ | Ethane-1,2-diamine | -2 | 2 | | $ONO$ | Nitrito-O | -1 | 1 | $CH_2-NH_2$ | | | | | $SCN$ | Thiocyanato-S | -1 | 1 | $CH_2-NH_2$ | | | | | $NCS$ | Thiocyanato-N | -1 | 1 | $(en)$ | Methanamine | 0 | 1 | | $H_2O$ | Aqua | 0 | 1 | $CH_3NH_2$ | Methanamine | 0 | 1 | | $NH_3$ | Ammine | 0 | 1 | $CH_2-NH_2$ | | | | | $CO$ | Carbonyl | 0 | 1 | $CH_2-N-CH_2COO$ | | | | | $NO$ | Nitrosyl | 0 | 1 | $CH_2-N-CH_2COO$ | | | | | $(C_6H_5)_3P$ | Triphenyl phosphine | 0 | 1 | $CH_2-N-CH_2COO$ | | | | | $C_5H_5N$ | Pyridine | 0 | 1 | $CH_2-N-CH_2COO$ | | | | | $(Py)$ | | | | ($EDTA$) | Ethylene diamine - tetraacetato | -4 | 6 | ## Definitions 1. **Coordination entity:** The ions or molecules bound to the central atom/ion in the coordination entity are called ligands. 2. **Coordination sphere:** The sphere in which donar atoms/ligating atoms are classified into unidentate and polydentate. 3. **Polyhedron:** Three-dimensional shape formed by the ligands and central metal. 4. **Central metal atom/ion:** The central metal atom or ion in a coordination complex. 5. **Ligands:** The ions or molecules bound to the central atom/ion in the coordination entity are called ligands. Ligands are classified into unidentate, didentate and polydentate. They are donar atoms/ligating atoms. 6. **Unidentate ligands:** When a ligand is bound to the central metal atom/ion through a single donar atom, the ligand is said to be uninidentate ligand. A unidentate ligand forms one co-ordinate bond. **e.g:** $Cl, NH_3$ (donar atom is $N$), $H_2O$ (donar atom is $O$). 7. **Didentate ligands:** When a ligand is bound to the central metal atom/ion through two donar atoms (ligating atoms), the ligand is said to be didentate ligand. Didentate ligand forms two co-ordinate bonds. **Density:** 2 **example:** $NH_2-CH_2-CH_2-NH_2$ (Ethane-1,2-diamine). Two N atoms serve as donar atoms. 8. **Polydentate ligands:** When a ligand is bound to the central metal atom/ion through many donar atoms (more than 2), the ligand is said to be polydentate. **example:** - $N(CH_2CH_2NH_2)_3$ - $EDTA$ (Ethylenediamine tetraacetate ion). - Density: 6 - Charge: -4 - $EDTA$ 4- is a hexadentate ligand. It can bind through four oxygen atoms and two N atoms to the central atom/ion. 9. **Denticity:** The number of ligating atoms/donar atoms present in a ligand is called its denticity. 10. **Chelate ligand and chelate:** When a didentate or a polydentate ligand uses two or more donar atoms simultaneously to bind to a single metal ion to form a ring-like structure is called a chelate ligand. The resulting complex is called chelate. - Chelate complexes are more stable than the similar complexes containing unidentate ligands. - **example:** $[Cu(NH_2CH_2CH_2NH_2)_2]^{2+}$ 11. **Ambidentate ligands:** A ligand which has two different donar atoms but forms only one co-ordinate bond at a time with the central atom in a complex is called ambidentate ligand. - **example:** 1. $NO_2$ and $ONO$ - $NO_2$ (nitrito-N) forms coordinate bond through 'N'. - $ONO$ (nitrito-O) forms coordinate bond through 'O'. - $M\leftarrow N=O$ and $M\leftarrow O-N=O$. 2. $SCN$ and $NCS$ - $SCN$ (thiocyanato-S) - Ligate through 'S'. - $NCS$ (thiocyanato-N) - Ligate through 'N'. - $M\leftarrow SCN$ and $M \leftarrow NCS$. 12. **Differentiate between ambidentate and didentate ligands:** A didentate ligand has two donar atoms and it forms two co-ordinate bonds with the central atom simultaneously. **e.g**: $(C_2O_4)^{2-}$ and $NH_2-CH_2-CH_2NH_2$ - Though an ambidentate ligand has two donar atoms, it can form only one coordinate bond at a time. - **Refer back.** ## Coordination Number - It is the total number of ligand donar atoms bonded in a complex. It is when a metal is directly bonded in a complex. - The coordination number of a central atom/ion is determined by counting the total number of $\sigma$-bonds (co-ordinate bonds) present around it. 'Pi' bonds are not counted for this purpose. ## Definitions - **Counter ion:** Ref. 248 - **Oxidation number:** Ref. 248 - **Homoleptic & Heteroleptic complexes:** Ref. Page 248. ## Questions - Which of the following species can act as ligands? - $NH_3, Cl, H_2O, CH_4, C_2H_6, NH_4^+$. - Classify the following complexes as homoleptic or heteroleptic: - 1. $[Co(NH_3)_6]^{3+}$ - 2. $[Co(NH_3)_4Cl_2]^+$ - 3. $[Co(en)_3]^{2+}$ ## Nomenclature - **Cationic Complex:** $[Fe(H_2O)_6] Cl_3$ → Hexaqua iron(II) chloride - **Neutral Complex:** $[Ni(CO)_4]$ → Tetracarbonyl Nickel(0) - **Anionic Complex:** $K_4[Fe(CN)_6]$ → Potassium hexacyanidoferrate (II) ## Questions 1. $[Co(NH_3)_6]Br_3$ 2. $[Co(H_2O)_3(CO_3)]Cl$ 3. $[Co(en)_2(ox)]NO_3$ or $[Co(NH_2-CH_2-CH_2-NH_2)_2(C_2O_4)]NO_3$ 4. $[PtI(NH_2CH_3)(NO_2)_2](NO_3)_3$ 5. $[CoBr(NO_2)_2(NO_3)_3]ClO_4$ 6. $[CoBr(NO_2)_2(NO_3)_3]Br$ 7. $[Cr(C_5H_5N)(Ph_3P)_3]SO_4Cl$ or $[Cr(C_5H_5N)(C_6H_5)_3P)_3]SO_4Cl$ 8. $[Co(NH_2-CH_2-CH_2-NH_2)_3]_2(SO_4)_3$ or $[Co(en)_3]_2(SO_4)_3$ 9. $[Zn(H_2O)_4]^{2+}$ 10. $[Co(H_2O)(CN)_2(en)_2]^+$ ## Isomerism in Co-ordination Compounds - Isomers are compounds with the same molecular formula but different structural arrangement. This phenomenon is called isomerism. - Isomers are classified into two: **Structural Isomers** and **Stereo Isomers**. - **Stereo isomers** are further classified into: **geometrical isomers** and **optical isomers**. - **Structural isomers** are classified into **ionisation isomers**, **co-ordination isomers**, **linkage isomer** and **solvate** isomerism. ## Stereo Isomers - Stereo isomers have the same chemical formula and stereo chemical bonds but they have different spatial arrangement. ## Geometrical Isomerism (cis-trans isomerism) - This type of isomerism arises in heteroleptic complexes. When two identical ligands occupy adjacent positions in a co-ordination polyhedra, the isomer is called **cis**. - When 2-identical groups occupy opposite positions, the isomer is called **trans**. - Geometrical isomerism is possible in square planar complexes (CN = 4) and octahedral complexes (CN= 6). - **Geometrical Isomerism in Square Planar Complexes** - Square planar complexes having unidentate ligands of the type $[MX_2L_2], [MX_2LA]$ and $MX_4L_2$ show geometrical isomerism. Where 'M' is the central atom and X, L, A and B are ligands. - **examples:** $[Co(NH_3)_2Cl_2], [Co(NH_3)_2ClBr], [Pt(NH_3)_2Cl_2]$ etc. - **Square planar complexes of the type $[MABXL]$ shows three isomers - two cis and one trans.** - **Square planar complexes of the type $[MX_2L_2]$ (e.g. $[Co(NH_3)_2Cl_2]$ and $[MX_3L]$ (e.g. $[Co(NH_3)_3Cl]$ do not show geometrical isomerism. Why?** - **Square planar complexes containing unsymmetrical didentate ligands show geometrical isomerism.** - **Geometrical Isomerism in Octahedral Complexes.** - Unidentate ligands of the formula $[MX_2L_4]$ show geometrical isomerism. **e.g**: $[Co(NH_3)_4Cl_2]^+$ - **Symmetrical didentate ligands** - Complexes of formula $[MX_2(L-L)_2]$ show geometrical isomerism. $L-L$ is a symmetrical didentate ligand. Like en ($H_2N-CH_2-CH_2-NH_2$), ox ($COO-COO$)- - **examples:** $[Co(en)_2Cl_2]^+, [NiCl_2(C_2O_4)_2]^-$ - **another type of geometrical isomerism occurs in octahedral complexes of the type $[MX_3L_3]$. If three donar atoms of the same ligands occupy adjacent positions at the corners of an octahedral face, we have the facial (fac) isomer. When positions are around the meridian of the octahedron, we get a meridional (mer) isomer.** ## Optical Isomerism - Optically active compounds rotate the place of polarised light. Optical isomers are isomers which bear optical mirror image relationship, but they are non superimposable on each other. They are also called enantiomers. Enantiomers possess the property of chirality. Optical isomers possess the same physical and chemical props, but one of the isomers rotate the plane of polarised light to right and the other one to the left. When the rotation is to the left, the isomer is called leavo rotatory ($l$ or -) and when the rotation is to the right, the isomer is called dextro rotatory ($d$ or +). - A 1:1 equilibrium mixture of $d$ and $l$ isomers give a racemic mixture with zero rotation. - **example:** **Complexes with three didentate ligands show optical activity** $e.g.$ $[Co(en)_3]^{3+}, (Cr(ox)_3)^{3-}$. - **Complexes with two didentate ligands and two monodentate ligands** $e.g.$ $[CoCl_2(en)_2]^+, [CrCl_2(ox)_2]^{3-} $. - These compounds show geometrical isomerism. The trans isomer is optically inactive. But cis isomer is. Therefore it is superimposable on its mirror image. Therefore it is optically active. It is not superimposable on its mirror image. Therefore it is optically active. ## Coordination Isomerism - This type of isomerism is shown by complex complexes where cation and anion are complex and there is an interchange of ligands between cationic and anionic entities taking place. - **e.g:** - $[Co(NH_3)_6][Co(CN)_6]$ and $[Co(NH_3)_6][Co(ox)_3]$ - $[Cu(NH_3)_4](PtCl_4)$ and $[Pt(NH_3)_4](CuCl_4)$ - $[Co(NH_3)_6][Co(SCN)_4]$ and $[Co(NH_3)_4(SCN)][Co(NH_3)_2(SCN)_4]$. ## Solvate Isomerism - This is a special type of coordination isomerism. Compounds which have the same composition but differ in the number of solvent molecules present as ligands and as free solvent molecules outside the coordination sphere (in the crystal lattice) are called solvate isomers. If $H_2O$ is the solvent, these isomers are called hydrate isomers. - **e.g**: $CrCl_3.6H_2O$ exists in these three isomer forms. - $[Cr(H_2O)_6]Cl_3$ (violet), - $[Co(H_2O)_5Cl]Cl_2.H_2O$ (pale green), - $[Cr(H_2O_4)Cl_2]Cl.2H_2O$ (dark green). - These three isomers can be identified by the addition of excess of $AgNO_3$ to their aq. solutions which precipitates chloride ions as $AgCl$ in the molar ratio 3:2:1 respectively. ## Bonding in Coordination Compounds - **Werner's theory:** This is the first theory which explains the structure of coordination compounds. - **Postulates of Werner's theory:** - **primary valencies:** Metals show two types of linkages (valencies) - primary and secondary. - **Primary valencies:** are normally ionisable and are satisfied by negative ions. - **Secondary valencies:** are non ionisable. The secondary valency is satisfied by neutral molecules or ions. The secondary valency number is fixed for a metal and is equal to the coordination number. - **The ions/groups bound by secondary linkages to the metal have definite spatial arrangement corresponding to different coordination numbers. This gives different geometrical to the complexes such as tetrahedral, octahedral and square planar geometry.** - **In modern theory, such spatial arrangements in co-ordination polyhedra. The species within the square bracket are coordination entities or complexes and the ions outside the square bracket are called counter ions.** - **[Do the questions from Page 245 & 246]** ## Limitations of Werner's Theory - His theory could not answer the following questions: - Why only certain metals form complexes? - Why the bonds in coordination compounds have directional properties? ## Why Coordination Compounds Have Characteristic Magnetic and Optical Properties? - To answer these questions, many theories have been put forwarded: - **Valence Bond Theory (VBT):** Explains the chemical bonding in coordination compounds. - **Crystal Field Theory (CFT):** - **Ligand Field Theory (LFT):** - **Molecular Orbital Theory (MO):** - **Valence Bond Theory (by Pauling):** ## Valence Bond Theory - **Postulates:** - Metal - ligand bond arises due to the donation of electron pairs by ligands to the central atom/ion. (Coordinate bond) - In order to accommodate these electrons, the central metal atom(ion) should possess requisite numbers of vacant orbitals of equal energy. Therefore, the metal atom(ion) uses its $(n-1)d$, ns, np or ns, np, nd orbitals for hybridization to yield a set of equivalent orbitals of definite geometry such as tetrahedral, octahedral, square planar etc. | C.N | Hybridization | Shape | |---|---|---| | 4 | $sp^3$ | Tetrahedral | | 4 | $dsp^2$ | Square planar | | 5 | $sp^3d$ / $dsp^3$ | Trigonal bipyramidal | | 6 | $sp^3d^2$ | Octahedral | | 6 | $d^2sp^3$ | Octahedral | - It is possible to predict the geometry of a complex from its magnetic behaviour. - In octahedral complexes, if inner d orbitals ($(n-1)d$ orbitals) are used, the hybridization is $dsp^2$ and the complex is termed as inner orbital complex or low spin complex (as spin paired) - If outer d orbitals (nd orbitals) are used, the hybridization is $sp^3d^2$ and the complex is termed as outer orbital complex or high spin complex or spin free complex. - **Magnetic props of coordination compds (by VBT theory)** - The magnetic moment of coordination compounds can be measured by magnetic susceptibility experiment. The results can be used to obtain information about the structure adopted by complexes. - **For metal ions with upto three electrons to the 'd' orbitals, the magnetic behaviour of free ions and their coordination entities is similar.** - **e.g:** $Ti^{3+}= [Ar] 3d^1$, $V^{3+} =[Ar]3d^2$ - $Cr^{3+} = [Ar] 3d^3$. - However the metal ions with $d^4$, $d^5$ and $d^6$ electrons, the magnetic behaviours of free ions and magnetic behaviour of the coordination entities need not be the same. ## Crystal Field Theory (CFT) - It is an electrostatic model, which considers metal-ligand bond to be ionic (arising purely from electrostatic interactions between the metal ion and the ligand. Ligands are considered as point charges in the case of anions and dipoles in the case of neutral molecules. The five $d$ orbitals in isolated gaseous atom/ion have the same energy (degenerate). This degeneracy is maintained if a spherically symmetrical negative charge surrounds the metal atom/ion. When the negative field is due to ligands (either anions or dipolar molecules like $NH_3$ and $H_2O$), in a complex, the degeneracy of $d$ orbitals is lifted. It results in splitting of d orbitals. The pattern of splitting depends on the nature of the crystal field. - **Crystal field splitting in octahedral complexes:** - Let us assume that six ligands are positioned symmetrically along the cartesian axes, with the metal atom at the origin. As the ligands approach the central metal ion, first there is an increase in energy of d-orbital (but degeneracy is maintained). Afterwards, the orbitals lying along the axes ($d_{z^2}$ and $d_{x^2-y^2}$) get repelled more strongly than the orbitals which are lying between the axes ($d_{xy}, d_{yz}, d_{xz}$) relative to the average energy of spherical crystal field. Thus degenerate $d$ orbitals split into two sets: the lower energy orbital set, $t_{2g}$ and higher energy orbital set $e_g$. This splitting of the degenerate levels due to the presence of ligands in a defined geometry is termed as crystal field splitting and the energy difference between $t_{2g}$ and $e_g$ is called crystal field splitting denoted by △o (△o for octahedral). - **Crystal field splitting energy $D_o$** is defined as the energy difference between the two sets of d-orbital energy levels (i.e. $t_{2g}$ and $e_g$). - **The energy of the two $e_g$ orbitals will increase by 3/5 $D_o$ and that of three $t_{2g}$ orbitals will decrease by 2/5 $D_o$.** ## Crystal Field Splitting in Tetrahedral Complexes - The crystal field splitting $D_o$ depends upon the field produced by the ligand and the charge on the metal ion. - For strong field ligands, splitting will be large and weak field ligands result in small splitting. - Ligands can be arranged in a series in the increasing order of field strengths. Such a series is called spectrochemical series. - $I < Br < SCN < Cl < S^{2-} < F^- < OH < ox^2- < H_2O < NH_3 < en < CN < CO$ ## Arrangement of e in Octahedral Coordination Entities - **Electronic configuration in complexes:** 1. **$d^1$ complexes:** Electronic Configuration - $t_{2g} \uparrow$ $e_g$ $ \downarrow \uparrow$. 2. **$d^2$ complexes:** Elec. config.: $t_{2g} \uparrow \uparrow$ $e_g$ $ \uparrow \downarrow \uparrow$ 3. **$d^3$ complexes:** Elec. config.: $t_{2g} \uparrow \uparrow \uparrow$ $e_g$ $ \uparrow \downarrow \uparrow$ 4. **$d^4$ complexes:** Here two possibilities occur. (i) The fourth electron may enter the $e_g$ orbital. (ii) The fourth electron may enter an $t_{2g}$ orbital. The actual configuration depends on the relative values of △o and P, where △o is crystal field splitting energy and P is the energy required for pairing. - **In weak field, △o<P, fourth election goes to one of the $e_g$ orbitals, giving the configuration $t_{2g}^3 e_g^1$. Thus high spin configuration is obtained.** - **In strong field, Doo>P, It becomes more energetically favourable for the fourth electron to occupy $t_{2g}$ orbitals with configuration $t_{2g}^4 e_g^0$. Which produce this effect is known as strong field ligands and form low spin complexes.** 5. **$d^5$ complexes:** - $△o<P$, high spin → Weak field configuration $t_{2g}^3 e_g^2$ - $△o>P$, low spin → Strong field ligand configuration $t_{2g}^5e_g^0$ 6. **$d^6$ Complexes:** - △o<P, high spin complexes - weak filed 🠒 config. $t_{2g}^4 e_g^2$ - △o>P, low spin - strong field 🠒 $t_{2g}^6e_g^0$ 7. **$d^7$ complexes:** - △o<P, high spin 🠒 weak filed → configuration $t_{2g}^5 e_g^2$ - △o>P, low spin 🠒 strong field 🠒 $t_{2g}^6e_g^1$ 8. **$d^8$**- Both fields $t_{2g}^6 e_g^2$ 9. **$d^9$** - Both fields $t_{2g}^6 e_g^3$ 10. **$d^{10}$** - Both fields $t_{2g}^6 e_g^4$ - Calculations show that $d^4$ to $d^7$ configuration entities are more stable for strong field as compared to weak field. ## Crystal Field Splitting in Tetrahedral Complexes - **Energy:** - $e_g$ $d_{x^2-y^2}, d_{z^2}$ - $t_{2g}$ $d_{xy}, d_{yz}, d_{xz}$ - **(d orbital to free ion spherical d orbital)** - In tetrahedral complexes, octahedral field splitting is smaller as compared to same metal ligand. For the same metal, same ligand and splitting energy is not large enough. Therefore orbital configuration are rarely observed. ## Colour of Transition Metal Complexes - One of the achievements of crystal field theory is its ability to explain the colours of coordination species. We know that colours of a substance is due to the absorption of light at a specific wavelength in the visible region and the transmission or reflection of the rest of the wavelengths (complementary colours). - In the case of transition metal complexes, the energy difference between $e_g$ set of orbitals and $t_{2g}$ set of orbitals is not very high. When visible light falls on them, $e_g$ get excited from lower set of orbitals to higher set of orbitals. The $e_g$ goes from $t_{2g}$ orbitals to $e_g$ orbitals. The energy needed for this excitation is absorbed from visible region and $e_g$ orbitals to $t_{2g}$ orbitals and the complementary colour (left over wavelengths) is reflected. The complex having $e_g$ specific wavelength gets excited from $t_{2g}$ level to $e_g$ level ($t_{2g}^n e_g^0$ → $t_{2g}^{n-1} e_g^1$) on the absorption of wavelengths from yellow-green region and appears as violet complementary colour. ## Bonding in Metal Carbonyls - Homoleptic carbonyls are formed by most of the transition metals e.g. $Ni(CO)_4$, $Cr(CO)_6$, $Fe(CO)_5$ etc. - The metal - carbon bond in metal carbonyl possess both $σ$ and **π** characters. The $M-C$ σ- bond is formed by the donation of lone pair of electrons on the carbonyl carbon into the vacant orbital of the metal. The $M-C$ **π** bond is formed by the donation of a pair of electrons from filled $d$-orbital of the metal to the vacant **π** orbital $C-O$ bond. - Thus $M-C$ bond possess both **σ** and **π** characters. - This synergic effect strengthens the bond between metal and $CO$. ## Shapes of Various Carbonyls - **$Ni(CO)_4$:** Tetra carbonyl nickel(0) - $hyb - sp^3$ - shape - tetrahedral - **$Fe(CO)_5$:** Pentacarbonyl iron (CO). - $dsp^3$ - Trigonal bipyramidal. - **$Cr(CO)_6$:** Hexa carbonyl. Chromium (0) - $dsp^3$ - Octahedral. - $Mn_2(CO)_{10}$: Decacarbonyl dimanganese(0) - **$Co_2(CO)_8$:** Octacarbonyldicobalt(o) has Co-Co bond bridged by two $CO$ (carbonyl) groups. ## Stability of Complexes - The stability of a coordination compound in solutions refers to the degree of association between the metal and ligand at equilibrium. The stability can be quantitatively expressed in terms of stability constant (β). The larger the value of stability constant, more stable the complex is. - $e.g: Cu + 4NH_3 \rightleftharpoons [Cu(NH_3)_4]^{2+}$ - $β$ = $[Cu(NH_3)_4]^{2+} / [Cu^{2+}][NH_3]^4$ - **Instability constant** or **dissociation constant** is the reciprocal of stability constant. - **Q: Calculate the overall complex dissociation equilibrium constant of $Eu(NH_3)_7^{2+}$, if its stability constant (β) is 21 × 10^13.** - **Ans: Dissociation constant = 1/β = 1/ (21 × 10^13)** ## Factors Affecting the Stability of a Complex 1. **Charge of the metal:** (directly proportional) 2. **Strength of the ligand:** (" ") 3. **Chelation:** (" ") ## Which One is More Stable in Each of the Following? 1. $K_4[Fe(CN)_6]$ & $K_3[Fe(CN)_6]$ 2. $[Fe(H_2O)_6]^{3+}$ & $[Fe(NH_3)_6]^{3+}$ 3. $[Fe(NH_3)_6]^{3+}$ & $[Fe(C_2O_4)_3]^{3-}$ ## Application of Coordination Compounds (Importance of Coordination Compounds) 1. **In quantitative and qualitative analysis:** - The presence of $Ni^{2+}$ can be qualitatively determined by adding dimethyl glyoxime in presence of $NH_4OH$ to the salt solution. A red ppt of Nickel dimethyl glyoxime complex gets obtained. - Hardness of water can be quantitatively estimated by complex titration with $Na_2EDTA$. - The $Ca^{2+}$ and $Mg^{2+}$ form stable complexes with $EDTA$. - The selective estimation of these ions can be done due to the difference in the stability constants of $Ca$ and $Mg$ complexes. 2. **Extraction of metals.** - **For example:** - $4Au + 8CN + 2H_2O + O_2 \rightarrow 4[Au(CN)_2]^- + 4OH$ - $2[Au(CN)_2]^- + Zn \rightarrow [Zn(CN)_4]^{2-} + 2Au$ - Firstly gold forms a complex with $CN$ in presence of $H_2O$ and O_2. Gold can be liberated from this complex by the addition of $Zn$. ## Applications of Coordination Compounds - **Purification of metals:** - **e.g.:** Impure $Ni$ is converted to $[Ni(CO)_4]$ which is decomposed to yield pure $Ni$. - $Ni+4CO \rightleftharpoons [Ni(CO)_4]$ - **Co-ordination compounds in the biological field:** - **example:** Chlorophyll is a coordination compound of $Mg$, haemoglobin, the red pigment of blood which is the oxygen carrier, is a coordination compound of $Fe$, Vitamin $B_{12}$ (cyanocobalamine), the anti pernicious anemia factor, is a coordination compound of cobalt. - **Enzymes:** Like carbonic anhydrase and carboxy peptidase A are coordination compounds. - **Coordination compounds as catalysts:** - **e.g.:** Wilkinson catalyst $ [RhCl(Ph_3P)_3]$ is used for the hydrogenation of alkene. - **In electroplating:** - Articles can be electroplated with $Ag$ and $Au$ smoothly and evenly from solutions of complexes $[Ag(CN)_2]^-$ and $[Au(CN)_2]^-$ than from a solution of simple metal ions. - **In photography:** - In black and white photography, the developed film is fixed by washing AgBr with hypo solution, which converts the undecomposed $AgBr$ to $[Ag(S_2O_3)_2]^{3-}$. - **In medicine:** - Chelates therapy: The excess of $Cu$ and $Fe$ in plants and animals are removed by the chelating (good D-penicillamine and des ferrioxine B) via the formation of complexes. EDTA is used in the treatment of $Pb$ poisoning. $Cr$ plated and related compounds effectively inhibit the growth of tumours. # Chap 9 - Complexes ## 9.1: Ref. Notes - Page 11 ## 9.2: - $FeSO_4$ does not form any complex with $(NH_4)_2SO_4$. - Instead it forms a double salt $FeSO_4(NH_4)_2SO_4. 6H_2O$ (Mohr salt) which dissociate completely. Thus it gives test for $Fe^{2+}$. - In the second case, a complex $[Cu(NH_3)_4]SO_4$ is formed. The complex ion $[Cu(NH_3)_4]^{2+}$ does not dissociate to give $Cu^{2+}$. Therefore it does not give the test for $Cu^{2+}$. ## 9.3: Ref. Text - Pages 240-241. ## 9.4: Ref. Text ## 9.5: 1. $[Co(H_2O)(CN)(en)_2]^{2+} \Rightarrow x+0-1+0+2 = +3, x=+3$ 2. $[CoBr_2(en)_2]^{+} \Rightarrow x- 2 + 0 = +1, x=+3$ 3. $[