Coordination Compounds PDF

## Coordination Compounds ### Introduction Compounds in which direct coordinate bonds are present between a metal/metal ion and a ligand. * Don't loose their identity even in the aqueous medium. * Example: `NH3 AgNH3` ### Ligands and Coordination Compounds #### Ligands: * Atoms / group of atoms...

## Coordination Compounds ### Introduction Compounds in which direct coordinate bonds are present between a metal/metal ion and a ligand. * Don't loose their identity even in the aqueous medium. * Example: `NH3 AgNH3` ### Ligands and Coordination Compounds #### Ligands: * Atoms / group of atoms which donate an electron pair to the metal or metal ion. * All ligands are Lewis bases but not all Lewis bases are ligands. * `m+ <-----> ligand (L.B.)` * `H+ <-----> lewis base ligand (X)` #### Representation: - `[M (Ligands)n]` #### Explanation: - `M` = Metal - `m` = Natural tendency of metal - `Ligands` = Ligands bonded to the central metal atom. - `n` = Coordination Number. ### Coordination Number * The number of coordinate bonds through which a central metal atom is joined with ligands. * Example: * `[Pt(NH3)2(Cl2)]` * CN = 4, OS = +2 ### Note: | Metal | CN | |:-------|:--------| | Ag | 2 | | Cu+ | 2, 4 | | Pt+2 | 4 | | Ni+2 | 4,6 | | Co+2 | 4,6 | | Pd+2 | 4,6 | | Al+3 | 6 | | Fe+3 | 6 | | Co+3 | 6 | For a fixed oxidation state, a metal can have only a fixed coordination number. ### Types of Complex Compounds: 1. **Positive:** - Complexes: Cation - Normal Ion: Anion - Example: `[Ag(NH3)2]+Cl` 2. **Anionic:** - Example: `K4[Fe(CN)6] 4K+ + [Fe(CN)6]4-` 3. **Neutral:** - Example: `[Ni(CO)]4 / [Pt(NH3)2Cl2]` 4. **Complex Cation:** - Example: `[Pt(NH3)4] [PtCl4]` ### Types of Ligands * **On basis of charge:** - **Negative Ligands**: * `F- , NO3- , O2-` * `Cl- , SO4^-2, O2 - ` * `Br , SO3^-2, O2^-2` * `I- , HSO3-` - **Positive Ligands**: * `NO+ , NH2-NH2` - **Neutral Ligands**: * `NH3 , NH2-NH2` * `Me-NH2 , NH2-NH` * `NH2 , N^-3 ` * `H2O , CH3-CH2-NH2` * `CH2-CH2 , (en)` * **Denticity:** * The number of coordinate bonds a ligand can form at a time with a central metal atom. * **On basis of Denticity:** - **Monodentate:** Can form only 1 coordinate bond with a central metal atom at a time. * Example: `H2O, NH3, Me-NH2, CO , etc.` - **Bidentate:** Can form 2 coordinate bonds with a central metal atom at a time. * Example: * `CH2-NH2` * `CH2-NH2` * `(en) Ethylenediammine` * `Me-CN-CH2` * `Me-CH-CH2` * `(bn) Butylene Diamine` * `CH2 - NH2` * `CH2-NH2` * `(pm) Propylene Diamine` - **Tridentate:** Can form 3 coordinate bonds with a central metal atom at a time. * Example: * `CH2-NH2` * `CH2-COO` * `CH2-NH2` - **Tetradentate**: Can form 4 coordinate bonds with a central metal atom at a time. - **Pentadentate:** Can form 5 coordinate bonds with a central metal atom at a time. * Example: * `[DMG]-` * `Ni2+ + DMG- = [Ni(DMG)2]` - **Hexadentate:** Can form 6 coordinate bonds with a central metal atom at a time. * Example: * `(EDTA)4-` * `Ethylenediamine tetra acetate` * `CH2` * `CH2COO` * `CH2NH2` * `CH2COO` * **Ambidentate ligand:** Ligands having 2 different donor sites but only one site is used for the donation of electrons. - Example: - `M=C=S-M` - `(thiocyanide)` - `(NCS)` - `NEC-SO-M` - `(isothiocyanide)` - `(SCN)` - `M=N=O` - `(Nitrito-N)` - `O=N=O-M` - `(Nitrato-O)` * **Flexidentate ligand:** - Ligands behaving as mono/bi dentate ligands * Example: * `[Pt(NH3)4CO3]2+` - Bidentate * `[Pt(NH3)4CO3]2+` - Monodentate ### Naming of Coordination Compounds 1. The cationic part is named first, then the anionic part. 2. Counter ions are written as in their usual names, mentioning their number. 3. In the complex part, the name of the ligand is mentioned first, then the metal. 4. If more than one ligand is used, use prefixes like di, tri, tetra, etc. for the ligand. If the ligand has two, three, or four or more repeating units use bis, tris, tetrakis, etc. 5. In the case of negative ligands ending in `e`, replace the `e` with `o`. - Example: - *Chloride:* Chloro / Chlorido - *Cyanide:* Cyano / Cyanido 6. The metal name is written just after the ligands, followed by its oxidation state. 7. If the complex is charged, the metal name contains `ate`. - *Example: - `Pb` - Plumbate - `Fe` - Ferrate - `Ag` - Argentate - `Au` - Aurrate 8. Example: - `K4[Fe(CN)6]` - Potassium hexacyanoferrate(II) - `[Ni(CO)4]` - Tetracarbonyl nickel(0) - `[Ag(NH3)2]Cl` - Diamminesilver(I) chloride - `[Pt(NH3)4Cl2]` - Tetrammineplatinum(II) chloride - `[CO(NH3)4(en)]+3` - Tetrammine(ethylenediamine)cobalt(II) ion ### Naming Of Bridging Complexes * Bridging ligands are the ligands attached to more than one central metal atom. * The prefix `M-` is used before their naming in complexes. * Example: `[Co(NH3)4(NH2)2CoCl4]` * Tetrammine cobalt(III) - diamido - tetrachlorido cobalt(II) * `2x + 0 - 4 - 2 = 2 - 3` ### Isomerism * **Stereo Isomerism:** * **Geometrical Isomers:** Based on the spatial arrangement of atoms or groups around the central metal ion. * **Optical Isomers:** Chiral molecules that are non-superimposable mirror images of each other. * **Structural Isomerism:** Based on the arrangement of atoms within the coordination sphere. * **Ionization Isomerism:** The exchange of atoms or ions between the coordination sphere and the counterion. * Example: `[Pt(NH3)4Cl2]Br2` and `[Pt(NH3)4Br2]Cl2` * **Hydrate Isomerism:** The difference in the number of water molecules coordinated to the central metal ion. * Example: * `[Cr(H2O)6Cl3]` is different from `[Cr(H2O)5Cl]Cl2.H2O` * `[Cr(H2O)4Cl2]Cl.2H2O` * There are isomers, so there are a total of n = n(H2O) isomers. * **Linkage Isomerism:** Occurs in complexes containing ambidentate ligands. Example: * `[Co(NH3)5NO2]Cl2` is yellow * `[Co(NH3)5ONO]Cl2` is red * **Spectroscopy:** * `E.T -> hv` is a spectroscopy technique for studying the electronic energy levels and molecular structure of chemicals. * `VIBGYOR` are the different colors of visible light, and each color has a corresponding wavelength and energy level. * Violet has the highest energy, and red has the lowest energy. * **Coordination Isomerism:** Both the cationic and anionic complexes have the same metal ions and the same types of ligands, but they are distributed between the cation and anion differently. * Oxidation State and coordination number are important factors to consider when determining if a complex can undergo a certain type of isomerism. * **Coordination Position Isomerism:** Present mainly in bridging complexes in which different ligands exchange their position. * **Ligand Isomerism:** A ligand itself may exhibit different isomeric forms and forms different complexes accordingly. * Example: - `[Pt(PM)(NH3)2]Cl2` - Propylene diammine is a 5-membered ring. - `[Pt(TN)(NH3)2]Cl2` - Trimethylene diammine is a 6-membered ring. * **Polymerisation Isomerism:** Complexes have the same empirical formula but they vary in the number of metal ions and ligands and ultimately affect the structure and composition of the complex. For example: * `[Pt(NH3)2Cl2](1:2:2)` and `[Pt(NH3)4][PtCl4](2:4:4)` * `[Pt(NH3)4][PtCl4]` can also exist in 1:2:2 ratio. ### Stereo Isomerism * **Tetrahedral:**(CN = 4) * Example: * `[M(AB)3]` * `[M(AA)2]` * `[M(AB)2]` * `[M(AA)Bc]` * Tetrahedral complexes - All geometrical isomers are optically inactive, except when a chiral ligand is also present. * **Square Planar:** * Example: For one bidentate ligand: * `[MAA2]` * `[MA2B2]` * `[MA2BC]` * `[MABCD]` * Geometric isomers * `cis` - same side * `trans` - opposite side * Example: * `[M(en)2]` ### Bidentate * `cis` and `trans` isomers can be formed only when the ligands on both sides are different. * Bidentate complexes with tetrahedral geometry does not show any optical isomerism. * Square planar complexes do not show optical isomerism for `[M(AA)2]` or `[M(AA)B]` but show optical isomerism for `[M(AB)2]` and `[M(AB)BC]`. ### Metal Carbonyls * Metal carbonyls are compounds in which metal is bonded to carbon monoxide (CO). * CO acts as ligands. * Back bonding occurs between the metal and CO. * Metal orbitals donate electrons to the empty π\* orbitals of CO, which increases the bond strength of the metal-carbon bond. * The stability of metal carbonyls is due to the synergistic effect, which is the combination of sigma bonding and π back bonding. * `[Mn(CO)5]` - Metal carbonyls like Mn(CO)5 can exist as both monomers as well as dimers and trimers in solution. * `Mn(CO)5` - can exist in two forms: * `Mn(CO)5` (monomer) * `[Mn(CO)5]2 ` (dimer) * `[Mn(CO)5]3 `(trimer) - The monomer is often more stable. - The stretching frequency of the C-O bond is directly proportional to the bond strength which is directly proportional to the bond order of the C-O bond. ### Non - Classical Ligands: π Acceptor Ligands * **Non-classical ligands** are mostly π acceptors that accept electron pairs from metal and donate electrons to metal. * **π Acceptor Ligands:** * Ligands that not only donate an electron pair to metal, but also accept electron pairs in their vacant orbitals. * Example: `CO`, `CNO`, `NO2` * Other Examples: `PR3`, `P(Ph)3`, `C2H4` ### The 18 - Electron Rule (EAN) * According to Sidgwick and Wig, a complex will be stable if the effective atomic number (EAN) of the central metal atom is nearest to the noble gas. * This theory is not valid because many complexes do not follow the 18-electron rule but they are known to exist as stable compounds. ### EAN Calculation * EAN = Atomic Number of Metal - Oxidation State + Number of electrons accepted by the metal. * Example: * `K3[Fe(CN)6]` - 26 - 3 + 2(6) = 35. * `K3[Fe(CN)6]` The EAN of Fe is 35 which is close to the nearest noble gas, Krypton (Kr) * This means that the complex is relatively stable. ### Werner's Bond Theory * According to Werner's bond theory, when a metal ion forms a complex compound it undergoes coordinate covalent bond formation. * During coordinate covalent bond formation, the metal ion accepts the electron pair from the ligand. * The metal ion provides a vacant orbital for the bond formation. * The arrangement of the vacant orbitals depends on * Ligand nature * Metal nature * Environment * Chelation * The metal ion is an electron pair acceptor (Lewis acid), and ligands are electron pair donors (Lewis base). ### Valence Bond Theory (VBT) * The VBT explains the formation of coordinate covalent bonds in coordination complexes. * The central metal atom uses its vacant d-orbitals to accept electron pairs donated by ligands. This results in the formation of coordinate covalent bonds. * To determine the hybridisation, consider the following factors: * Ligand nature (strong field ligand or weak field ligand). * Metal nature (3d, 4d or 5d). * Environment of the coordination complex. * The chelating effect of ligands. * In coordination complexes, the arrangement of the d-orbitals changes, leading to the formation of inner orbital complexes and outer orbital complexes. * Outer orbital complexes are formed when the d-orbitals of the metal atom are already filled, and the bonding involves the higher energy outer orbitals. * Inner orbital complexes are formed when the d-orbitals of the metal atom are hybridized with the lower energy vacant orbitals. * **Spin Only Magnetic Moment:** This is used to calculate the spin only magnetic moment of the coordination complex. * `µ = √(n)(n+2)` * Where `n` is the number of unpaired electrons. * This method is based on the assumption that the electrons in the d-orbitals are responsible for the magnetic moment. * **Example:** * `[Fe(CN)6]^3-` * Fe: `[Ar] 4s^2 3d^6` * Fe^3+: `[Ar] 3d^5` * The Fe^3+ ion is a strong field ligand. The electrons are paired. * `[Fe(CN)6]^3-` - Inner orbital complex. * The hybridization is `d^2sp^3` * The magnetic moment can be calculated using the spin only formula: * `µ = √ (1)(1+2) = √3 = 1.73 BM` * **Example:** * `[Ni(en)3]^2+` * `Ni [Ar} 4s^2 3d^8` * `Ni2+ [Ar} 3d^8` * `(en) is a strong field ligand.` - `[Ni(en)3]^2+` - Inner orbital complex. - Hybridization: `d^2sp^3` * Magnetic moment: `µ= √ (2)(2+2) = √8 =2.83 BM` * **Example:** * `[Cr(H2O)6]^3+` * `Cr [Ar} 4s^1 3d^5` * `Cr^3+ = [Ar} 3d^3` * `(H2O)` is a weak field ligand * `[Cr(H2O)6]^3+` - Outer orbital complex. * Hybridization: `sp^3d^2` * The magnetic moment can be calculated using the spin only formula. * `µ= √ (3)(3+2) = √15 = 3.87 BM` ### Crystal Field Theory (CFT) * The CFT describes the splitting of d-orbitals of transition metal ions in coordination complexes, resulting in different electronic energy levels and a stability change. * According to CFT: * The bonding is primarily electrostatic interactions between the positively charged metal ion and the negatively charged ligands. * The ligands are not bound by covalent bonds, but by electrostatic interactions. * This theory is not based on hybridization, but rather on the interaction of the metal ion electronic configuration with the electric field generated by the ligands. * **Octahedral complexes** are formed when six ligands approach the metal ion in such a way that the metal ion occupies a position at the center of an octahedron. * **Octahedral complexes** can be visualized as a cube in which a metal ion is present at the body center (BCC) and ligands are approaching from the cube faces (FCC). #### Crystal Field Splitting * When ligands approach the metal ion, the d-orbitals of the metal ion split into two sets with different energy levels. * This energy difference is called crystal field splitting energy (CFSE). * The energy levels are represented as: * `e_g (dx^2-y^2 and dz^2)` * `t_{2g} (d_{xy}, d_{xz}, and d_{yz})` * `Δ_o` refers to the energy difference between `e_g` and `t_{2g}` orbitals. * The magnitude of `Δ_o` is dependent on a few factors: * Nature of the metal ion. * Nature of ligands. * Geometry of the complex. * **High spin complexes:** The complexes in which the metal ion maximizes unpaired electrons. This is generally observed when `Δ_o` is less than pairing energy. * **Low spin complexes:** The complexes in which the metal ion minimizes unpaired electrons, and the electrons are paired in the lower energy `t_{2g}` orbitals. This is generally observed when `Δ_o` is greater than pairing energy. * **Example:** * `[Fe(CN)6]^3-` * Fe^3+ `= 3d^5` * CN- is a strong field ligand. * `Δ_o` is greater than the pairing energy, resulting in a low spin complex with a `d^2sp^3` hybridization and one unpaired electron. * **Example:** * `[Fe(H2O)6]^3+` * Fe^3+ `= 3d^5` * H2O is a weak field ligand. * `Δ_o` is less than the pairing energy, resulting in a high spin complex with a `sp^3d^2` hybridization and five unpaired electrons. ### Crystal Field Stabilization Energy (CFSE) * The CFSE is the additional stability gained by a coordination complex due to the splitting of the d-orbitals in the presence of a crystal field. * The CFSE is a measure of the energy difference between the ground electronic state of the complex and the free ion state. * This value is usually expressed in terms of `Δ_o` and the number of electrons in `t_{2g}` and `e_g` orbitals. * **CFSE = 0.4n*(t_{2g}) - 0.6n(e_g) + P(n_p)` * `n` is the number of electrons in the orbitals. * `P` is the pairing energy, expressed in terms of `Δ_o`. * If `P = 0` for low spin complex, * If `P = 0.6 Δ_o` for high spin complexes * If a complex has a higher CFSE, it is more stable. ### Splitting In Tetrahedral Complexes * In tetrahedral complexes, the ligands are located at the four corners of a tetrahedron. * The d-orbitals split into two sets of different energies: * `e(d_{xy}, d_{xz}, d_{yz})` * `t_2(d_{x^2-y^2}, dz^2)` * The energy difference between these orbitals is denoted by `Δ_t`. * In tetrahedral complexes, `Δ_t` is smaller than `Δ_o` for the same metal ion and ligands because the ligands interact less strongly with the metal ion in a tetrahedral field compared to an octahedral field. * The CFSE for a tetrahedral complex can be calculated using the equation: * CFSE = `- 0.6n(e) + 0.4n(t_2)` + `nP_o` * `n` is the number of d-electrons in the orbitals. * `P_o` is the pairing energy (expressed in terms of `Δ_t`). ### Splitting In Square Planar Complexes * For square planar complexes, the d-orbitals split into three sets with different energies: * `d_{x^2-y^2}` (highest energy) * `d_{xy}, d_{xz}, d_{yz}` (intermediate energy) * `dz^2` (lowest energy) * `Δ_s` is denoted for the difference in energy levels * **Square Planar complexes** are less common than octahedral and tetrahedral complexes because the ligands are located in a plane perpendicular to the axis of the complex. ### Applications of CFSE * **Color of coordination compounds:** * The color of a coordination compound is due to the absorption of specific wavelengths of visible light. * When visible light falls on the coordination compound, electrons in the lower energy d-orbitals get excited to the higher energy levels through. * The compound absorbs a certain wavelength of visible light, which corresponds to the energy difference between the two d-orbitals, and the complementary color is observed. The wavelength of light absorbed. * **d-d transition:** This type of transition involves the excitation of an electron from a lower energy d-orbital to a higher energy d-orbital. * **Charge transfer spectra:** This occurs by promoting an electron from a ligand orbital to a metal orbital or vice versa, resulting in a different absorbance spectrum. * **Stability of Coordination Complexes:** CFSE is a crucial factor in determining the stability of coordination complexes. Complexes with higher CFSE values are more stable. * A complex with a greater number of d-electrons in the lower energy orbitals will have higher stability and form a more stable complex. * As the coordination number increases, the stability of the complex increases but has its limitations and reaches a certain point where its stability decreases. * **Stepwise Formation Constant:** This constant is a measure of the stability of a complex as ligands are added to the metal ion one at a time. The higher the formation constant, the more stable the complex is. * **Overall Formation Constant:** This constant is a measure of the overall stability of a complex when all the ligands are added to the metal ion at once. * **Chelation Effect:** The increased stability of a complex when a chelating ligand is present. * Ligands have a strong effect on the coordination number. * Chelating ligands form more stable complexes compared to unidentate ligands. * The formation constant and the overall stability of the coordination complex increase when the denticity of the ligand increases or when the ligand forms more rings. * **Metal Carbonyls:** * Metal carbonyls are coordination complexes in which one or more CO molecules bind to a metal atom. * Metal carbonyls exhibit a combination of sigma bonding and π back bonding. * The back bonding contributes to the unique properties of metal carbonyls including: * **High stability** : Back bonding strengthens the M-CO bond. * **Low melting and boiling points** : This is explained by the weakening of the M-CO bond due to back bonding. * **Effective Atomic Number (EAN):** The EAN theory is a simple way to predict the stability of coordination complexes. According to this rule, a coordination complex is considered stable when it achieves an electron configuration similar to that of a noble gas. Based on this concept, the stability of metal carbonyls is explained by the fact that they can achieve a noble gas electron configuration. * **Crystal Field Splitting Energy and Magnetic Properties:** The Crystal Field Splitting Energy (CFSE) is a crucial factor in determining the magnetic properties of coordination complexes. * **Weak Field Ligands:** Weak field ligands produce a small crystal field splitting energy and lead to high spin complexes with a greater number of unpaired electrons and a higher magnetic moment. * **Strong Field Ligands:** Strong field ligands produce a larger crystal field splitting energy which leads to low spin complexes with a lower number of unpaired electrons and a lower magnetic moment.