Coordination Chemistry
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TOPIC : 1. Introduction to complex, coordination chemistry. 2. Ligands explained (with eg) 3. Classification of complexes (with eg) 4. EAN effective atomic number (with eg)
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Transition Metal Electronic Spectra
The document provides information about electronic spectra and terms for carbon p electrons and transition metal d electron configurations. It discusses:
1) Possible terms that arise from carbon's 2p electrons, including 1D2, 3P2, 3P1, 3P0 and 1S0 terms. Hund's rules are used to determine the ground state term.
2) Microstate tables that list all possible combinations of orbital angular momentum (L) and spin (S) for electron configurations.
3) Tanabe-Sugano diagrams that show the splitting of d electron terms in an octahedral ligand field and allow determination of transition energies.
4) Charge transfer transitions that can occur from the
Orgel diagram
This document discusses the electronic structure and spectra of metal complexes. It begins by introducing ligand field theory and how the d orbitals of the metal ion split into different energy levels depending on the geometry and ligand field strength. Orgel diagrams are used to illustrate the splitting patterns for different d electron configurations from d1 to d10. Selection rules for electronic transitions are described. Tanabe-Sugano diagrams show how transition energies vary with ligand field strength. Methods for determining the ligand field splitting parameter (Δo) from experimental spectra are also outlined, along with examples of different types of spectra observed.
Chemical bonding part 2
VSEPR theory describes the electron-group arrangements and molecular shapes that result from electron-pair repulsions around a central atom. The theory states that valence electron groups will adopt an arrangement that minimizes repulsions between these groups. This results in five basic molecular geometries - linear, trigonal planar, tetrahedral, trigonal bipyramidal, and octahedral. Factors such as double bonds, lone pairs, and differing atomic sizes can cause deviations from ideal bond angles predicted by VSEPR theory.
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This document discusses atomic structure and electron configuration. It begins by explaining Slater's rules for calculating effective nuclear charge. It then provides examples of applying Slater's rules to determine electron shielding and effective nuclear charge. The document also covers electron configurations, term symbols, Hund's rules, and periodic trends in atomic size, ionization energy, and metallic character across periods and groups. It defines concepts like ionization potential, electron affinity, and electronegativity scales. In summary, the document provides an in-depth overview of theoretical atomic structure concepts.
Orgel diagrams; D and F/P Orgel Diagrams
Orgel diagrams depict the splitting of energy levels in transition metal complexes. They show the splitting of d electron configurations into terms based on whether the complex has an octahedral or tetrahedral ligand field. There are two main types of Orgel diagrams: D diagrams for d1, d4, d6, d9 complexes and F/P diagrams for d2, d3, d7, d8 complexes. The diagrams qualitatively show the possible electronic transitions between terms based on the complex's geometry and electron configuration. Orgel diagrams are useful for understanding the optical, magnetic, and spectral properties of transition metal complexes.
How To Write Electron Configurations
1. Determine the total number of electrons by finding the element's atomic number on the periodic table.
2. Fill the lowest energy electron shells first, following the order of filling: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, etc.
3. Write the electron configuration by specifying the energy level, subshell, and number of electrons in each subshell using superscripts. The valence electrons are those in the highest occupied shell.
IBDP HL atomic structure
This document provides information about electron configurations and the filling of orbitals in atoms. It includes trends in ionization energies from the periodic table, evidence for electron sublevels, and rules for filling orbitals according to the Aufbau principle, Hund's rule, and the Pauli exclusion principle. The document poses practice problems in predicting ionization energies and writing electron configurations for various elements.
2012 Orbital Hybrization, Sigma and Pi Bonds
Carbon forms four equal hybrid orbitals through sp3 hybridization to allow methane to adopt its tetrahedral electron geometry. This involves one s orbital and three p orbitals combining to form four new hybrid orbitals. Sigma bonds are formed by the head-on overlap of hybrid orbitals. Sp2 hybridization with one s and two p orbitals results in trigonal planar geometries like ethene. Pi bonds in double and triple bonds involve overlap of unhybridized p orbitals above and below the sigma bond.
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The document provides information about electronic spectra and terms for carbon p electrons and transition metal d electron configurations. It discusses:
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2) Microstate tables that list all possible combinations of orbital angular momentum (L) and spin (S) for electron configurations.
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Orgel diagram
This document discusses the electronic structure and spectra of metal complexes. It begins by introducing ligand field theory and how the d orbitals of the metal ion split into different energy levels depending on the geometry and ligand field strength. Orgel diagrams are used to illustrate the splitting patterns for different d electron configurations from d1 to d10. Selection rules for electronic transitions are described. Tanabe-Sugano diagrams show how transition energies vary with ligand field strength. Methods for determining the ligand field splitting parameter (Δo) from experimental spectra are also outlined, along with examples of different types of spectra observed.
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VSEPR theory describes the electron-group arrangements and molecular shapes that result from electron-pair repulsions around a central atom. The theory states that valence electron groups will adopt an arrangement that minimizes repulsions between these groups. This results in five basic molecular geometries - linear, trigonal planar, tetrahedral, trigonal bipyramidal, and octahedral. Factors such as double bonds, lone pairs, and differing atomic sizes can cause deviations from ideal bond angles predicted by VSEPR theory.
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This document discusses atomic structure and electron configuration. It begins by explaining Slater's rules for calculating effective nuclear charge. It then provides examples of applying Slater's rules to determine electron shielding and effective nuclear charge. The document also covers electron configurations, term symbols, Hund's rules, and periodic trends in atomic size, ionization energy, and metallic character across periods and groups. It defines concepts like ionization potential, electron affinity, and electronegativity scales. In summary, the document provides an in-depth overview of theoretical atomic structure concepts.
Orgel diagrams; D and F/P Orgel Diagrams
Orgel diagrams depict the splitting of energy levels in transition metal complexes. They show the splitting of d electron configurations into terms based on whether the complex has an octahedral or tetrahedral ligand field. There are two main types of Orgel diagrams: D diagrams for d1, d4, d6, d9 complexes and F/P diagrams for d2, d3, d7, d8 complexes. The diagrams qualitatively show the possible electronic transitions between terms based on the complex's geometry and electron configuration. Orgel diagrams are useful for understanding the optical, magnetic, and spectral properties of transition metal complexes.
How To Write Electron Configurations
1. Determine the total number of electrons by finding the element's atomic number on the periodic table.
2. Fill the lowest energy electron shells first, following the order of filling: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, etc.
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The spectrochemical series arranges ligands in order of their crystal field splitting parameter (Δ), which indicates their ability to repel electrons in a metal-ligand complex. Strong field ligands like cyanide cause large Δ and greater splitting of d-orbital energies. In an octahedral complex, strong field ligands create a large energy gap between the lower-energy t2g and higher-energy eg orbitals, forcing electrons into the lower t2g orbitals and producing a complex with low spin. Weak field ligands like halides cause small Δ and less splitting, allowing electrons to fill orbitals normally and producing a high-spin complex. The type of ligand affects the splitting of orbitals and spin state in transition
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1. Carbon atoms can form more than 4 bonds through hybridization of orbitals. In methane, carbon's 2s and 2p orbitals hybridize to form 4 equivalent sp3 hybrid orbitals, allowing carbon to form 4 sigma bonds to hydrogen in a tetrahedral structure.
2. In ethene and ethyne, carbon's orbitals hybridize differently to form pi bonds in addition to sigma bonds. Ethene carbons hybridize to form 3 sp2 orbitals and 1 p orbital, allowing 2 sigma bonds and 1 pi bond to other carbons. Ethyne carbons hybridize to form 2 sp orbitals and 2 p orbitals, allowing 1 sigma bond and 1 pi bond to each
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This document discusses galvanic cells and cell potential. It begins by defining oxidation-reduction reactions and half-reactions. It then explains how galvanic cells use redox reactions to produce an electric current and discusses the components of galvanic cells including the salt bridge, electrodes, and direction of electron and ion flow. The document introduces standard reduction potentials and how to calculate cell potential from half-cell potentials. It explains how cell potential depends on concentration using Le Châtelier's principle and the Nernst equation. Examples are provided to demonstrate how to calculate cell potentials, determine reaction spontaneity, and predict changes in potential with changing concentrations.
Coordination Chemistry.ppt
This document discusses coordination chemistry and Alfred Werner's theory of coordination compounds from the late 19th century. Some key points:
1. Werner proposed that transition metals form complexes where the metal ion is at the center coordinated by other ligands. He termed the metal-ligand bonds as primary and secondary valences.
2. Through experiments with cobalt chloride and ammonia complexes, Werner showed that the complexes have defined compositions and proposed formulae such as [Co(NH3)6]3+.
3. Werner's theory established the field of coordination chemistry and that transition metals can exhibit multiple oxidation states through bonding with ligands in complexes. His work resolved a longstanding controversy over chemical bonding models.
week2-d3-Introduction to Bonding.pdf
This document provides an introduction to bonding, including:
- How chemical bonds form via the sharing or transfer of valence electrons between atoms
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Chapter7
This document discusses several key atomic and molecular properties including electron configurations, ionization energies, atomic and ionic radii, and electron affinity. It explains how atoms gain or lose electrons to form ions that have noble gas configurations. Cations are typically smaller than the parent atom, while anions are larger. Trends in various properties across the periodic table are also examined, such as how ionization energy generally increases moving left to right and up a group, and how atomic radius decreases with increasing nuclear charge. Diagonal relationships between elements are explained in terms of similar cation charge densities.
Chemical Bonding I: Basic Concepts
For Chem 1:
Significanceof the ELectron in Bonding
The Octet Rule
Lewis Symbol/Structures
Formal Charge
Polyatomic Ions
Types of Bonds (Ionic, Covalent, Coordinate Covalent, Metallic Bonds, Multiple Bonds)
Exceptions to the Octet Rules
Oxidation Number is not included in the class discussion and exam. ;D
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Coordination Chemistry
- 1. BIOCHEMISTRY TOPIC : 1. Classification of complexes 2. EAN (eg) 3. Nomenclature (eg) - MINHAJUL ABEDEN - GOVERNMENT SCIENCE COLLEGE AUTONOMOUSUS, BANGALORE COORDINATION CHEMISTRY
- 2. 1. Ligands definition ● ion/atom ● 1 or more donor atom ● Each donate 1 pair of e- ● I.e lewis base 1. Classification 1. based on how the ligands bond to CMA ● Monodentate ● Bidentate ● Polydentate ● Ambidentate ● chelating 1. Based on charge of complex formed ● Cationic ● Anionic ● neutral 2.3 Based type of ligands ● Homoleptic ● heteroleptic
- 3. 1. LIGANDS An ion or molecule that binds to a central metal atom to form a coordination complex & the binding force is ccordination bond.
- 4. CLASSIFICATION 1.1 Based on their donor atoms [number of donor atoms] 1. MONODENTATE ligand Ligand donate 1 pair of e- To bond to CMA by one bond 2. BIDENTATE 2 pair of e- 2 bonds to CMA 3 POLYDENTATE more than 2 more than 2 bonds to CMA
- 5. 1.2 Number of donor atoms a. Ambidentate More than 2 donor atoms Ligates thru one donor atom a. Chelating Polydentate ligands 1 ligand Ligates thru 2 donor atoms Forms ring strusture
- 6. e
- 7. 1.3 Based on charge of complex formed a. Anionic The coordination complex is -vely charged (when ionic sphere removed) a. Cationic a. Neutral NOTE : the ligands can also be classified based on charge on the ligand specie: i.e Neutral, Cationic, anionic ligand
- 8. 1.4 BASED ON TYPE OF LIGAND (complexes are classified) a. Homoleptic complex only one type of ligand [Co(NH3)6]3+ a. Heteroleptic more than one type of ligands [Co(NH3=)4 Cl2]+
- 9. EAN (eff atomic number) Total number of e- present with the CMA (includes the e - donated by ligands). The CMA accepts e- to complete octate confn EAN= Z - X + Y Z atomic num. of CMA X e- lost by CMA on cmplx frmn y e- gained frm ligand
- 10. EAN concept explained : EAN = eff. atomic num. EAN = total num of e- present with CMA including those donated by ligands (in stable complex) EAN = the number is always equal to atomic no. of nearest noble gas (nearest to metal in periodic table) EAN answer is always equal to Z of noble gas EAN = (total no. of e- in CMA) + (no. of e- donated by ligands) EAN = [(atomic no. Of CMA) - (Oxydn state of CMA)] + [no. e- donated by ligands] EAN = (z) - ( x ) + [ y ] EAN = z-x+y ● Oxidation state is subtracted from Z because many times metals have +ve O.S i.e they lost e- ● eg Fe2+ have 24 e- compared to ground state config on 26
- 11. Solving EAN 1. Note the atomic no. of CMA at ground state = Z 2. Find out O.S of CMA = x 3. Determine coordination number = y 4. Substitute in EAN = Z-X+Y Note its necessary to remember - atomic number & oxidation state of transition elements - names and formula of important complexes
- 12. EAN examples 1. Potassiunferrocyanide K4 [ Fe (CN)6 ] Note : EAN = Z - X +Y Ferro : +2 charge CN has -1 charge in its ionic stage [ Fe (CN)6 ] is the coon sphere K4 is the ionic sphere a. Calculate O.S of Fe (Fe loses two e- to become Fe2+ with 24 electrones) ( that will be the X) -4 = X - 6 X = +2 a. (Coordination num) * 2 = Y Y = 2 * 6 Y= 12 a. Z = atomic number Z = 26 EAN = 26 - 2 + 12 EAN = 36 ( i.e the config of noble gas Krypton ) Ean of potassiumferrocyanide is 36 hence its a stable complex
- 13. 2 . Potassiunferricyanide K3 [ Fe (CN)6 ] Note : EAN = Z - X +Y Ferro : +3 charge CN has -1 charge in its ionic stage [ Fe (CN)6 ] is the coon sphere K3 is the ionic sphere a. Calculate O.S of Fe (Fe loses three- to become Fe3+ with 23 electrones) ( that will be the X) -3 = X - 6 X = +3 a. (Coordination num) * 2 = Y Y = 2 * 6 Y= 12 a. Z = atomic number Z = 26 EAN = 26 - 3 + 12 EAN = 35 Ean of potassiumferricyanide is 35 , IT ACTS AS OXYDISING AGENT
- 14. 3. Cupraammoniumsulphate [ Cu (NH3)4 ] SO4 Note : EAN = Z - X +Y Cupra: +2 charge Sulphate has -2 charge in its ionic stage [Cu (NH3)4 ] is the coon sphere SO4 is the ionic sphere a. Calculate O.S of Cu (Cu loses two e- to become Cu2+ with 29 electrones) ( that will be the X) -2 = X - 4 X = +2 a. (Coordination num) * 2 = Y Y = 2 * 4 Y= 8 a. Z = atomic number Z = 29 EAN = 29 - 2 + 8 EAN = 35 Ean of Cupraammoniumsulphate is 35 IT ACTS AS OXYDISING AGENT