1) Slater proposed rules to calculate the effective nuclear charge (Zeff) felt by outer electrons by accounting for shielding effects.
2) Zeff is calculated as Z - S, where Z is the nuclear charge and S is the Slater screening constant.
3) Slater's rules provide a method to calculate S by considering the shielding contributions of core and other valence electrons.
This document discusses hybridization theory in chemistry. It explains that hybridization involves combining atomic orbitals to form new hybrid orbitals that explain molecular geometry. The key types of hybridization are sp3, sp2, and sp for molecules with 4, 3, and 2 electron groups around a central atom. sp3 hybridization results in methane's tetrahedral structure. sp2 hybridization gives ethylene's 120° bond angles. sp hybridization linearizes carbon dioxide. Sigma bonds form by hybrid orbital overlap at maximum density, while pi bonds involve unhybridized p-orbital overlap at lower density.
The document discusses molecular orbital theory (MOT) and its application to transition metal complexes. It provides details on:
1) How MOT was developed in the 1930s and uses the linear combination of atomic orbitals (LCAO) method to combine metal and ligand atomic orbitals into molecular orbitals.
2) The principles of ligand field theory, which describes bonding in coordination complexes using sigma and pi bonding between the metal and ligands.
3) How MOT is used to construct molecular orbital diagrams for octahedral transition metal complexes, showing the splitting of metal and ligand atomic orbitals into bonding, non-bonding, and antibonding molecular orbitals.
4) Examples of applying MOT to
This document discusses crystal field stabilization energy (CFSE), which is the energy gap between split d-orbital energy levels caused by ligands interacting with a central metal atom. It provides information on how CFSE is calculated for octahedral and tetrahedral complexes, and factors that affect CFSE such as the nature of ligands and metal cation, complex geometry, and the metal's quantum number.
1) The document summarizes aromaticity and anti-aromaticity, discussing the key criteria for aromatic compounds including being cyclic, conjugated, planar, and having 4n+2 π electrons.
2) It provides examples of aromatic compounds like benzene and non-aromatic compounds that do not meet the criteria. Anti-aromatic compounds are also discussed.
3) NMR spectroscopy is discussed as a way to distinguish aromatic protons from other types of protons based on chemical shift values.
This document provides an introduction to molecular orbital theory (MOT), which describes bonding between atoms using molecular orbitals formed from the combination of atomic orbitals. MOT allows prediction of electron distribution, molecular properties like shape and magnetism, and bond order strength. Principles of MOT include molecular orbitals having lower energy than separated atoms, causing electrons to prefer molecular bonds. Sigma and pi bonds are discussed, with sigma bonds symmetrical along the axis and pi bonds having side-by-side overlap with electron density above and below the axis. Bond order is calculated using the number of electrons in bonding and antibonding orbitals and indicates bond strength, with higher bond order meaning stronger, more stable bonds.
Crystal Field Theory explains the colors of transition metal complexes based on ligand-metal interactions. The electrostatic interaction between ligands and metal d-orbitals splits the d-orbital energies. For an octahedral complex, the d-orbitals point directly at ligands have higher energy than those that bisect ligands. This splitting pattern determines if the complex is high or low spin, which then dictates its color and magnetic properties. The spectrochemical series orders ligands by their ability to cause crystal field splitting, correlating ligand type with complex color.
Slater's rules provide a method to estimate the shielding of electrons and the effective nuclear charge experienced by electrons in an atom. The rules involve writing the electron configuration, ignoring higher energy level electrons, and applying shielding constants of 0.35 for electrons in the same subshell and 0.85 for electrons in the previous subshell. As an example, the rules are used to calculate that the effective nuclear charge experienced by the valence electrons of nitrogen is 3.9 instead of the actual nuclear charge of 7.
Molecular orbital theory(mot) of SF6/CO2/I3-/B2H6sirakash
1) Molecular orbital theory views a molecule as delocalized molecular orbitals formed from linear combinations of atomic orbitals. Bonding molecular orbitals are lower in energy due to constructive interference, while antibonding orbitals are higher in energy due to destructive interference.
2) The document provides examples of applying molecular orbital theory to SF6, CO2, B2H6, and I3- molecules. It describes the atomic orbitals and molecular orbitals formed, including bonding, antibonding, and non-bonding orbitals, and explains how the molecular orbitals rationalize the electronic structures and bonding patterns in these molecules.
This document discusses hybridization theory in chemistry. It explains that hybridization involves combining atomic orbitals to form new hybrid orbitals that explain molecular geometry. The key types of hybridization are sp3, sp2, and sp for molecules with 4, 3, and 2 electron groups around a central atom. sp3 hybridization results in methane's tetrahedral structure. sp2 hybridization gives ethylene's 120° bond angles. sp hybridization linearizes carbon dioxide. Sigma bonds form by hybrid orbital overlap at maximum density, while pi bonds involve unhybridized p-orbital overlap at lower density.
The document discusses molecular orbital theory (MOT) and its application to transition metal complexes. It provides details on:
1) How MOT was developed in the 1930s and uses the linear combination of atomic orbitals (LCAO) method to combine metal and ligand atomic orbitals into molecular orbitals.
2) The principles of ligand field theory, which describes bonding in coordination complexes using sigma and pi bonding between the metal and ligands.
3) How MOT is used to construct molecular orbital diagrams for octahedral transition metal complexes, showing the splitting of metal and ligand atomic orbitals into bonding, non-bonding, and antibonding molecular orbitals.
4) Examples of applying MOT to
This document discusses crystal field stabilization energy (CFSE), which is the energy gap between split d-orbital energy levels caused by ligands interacting with a central metal atom. It provides information on how CFSE is calculated for octahedral and tetrahedral complexes, and factors that affect CFSE such as the nature of ligands and metal cation, complex geometry, and the metal's quantum number.
1) The document summarizes aromaticity and anti-aromaticity, discussing the key criteria for aromatic compounds including being cyclic, conjugated, planar, and having 4n+2 π electrons.
2) It provides examples of aromatic compounds like benzene and non-aromatic compounds that do not meet the criteria. Anti-aromatic compounds are also discussed.
3) NMR spectroscopy is discussed as a way to distinguish aromatic protons from other types of protons based on chemical shift values.
This document provides an introduction to molecular orbital theory (MOT), which describes bonding between atoms using molecular orbitals formed from the combination of atomic orbitals. MOT allows prediction of electron distribution, molecular properties like shape and magnetism, and bond order strength. Principles of MOT include molecular orbitals having lower energy than separated atoms, causing electrons to prefer molecular bonds. Sigma and pi bonds are discussed, with sigma bonds symmetrical along the axis and pi bonds having side-by-side overlap with electron density above and below the axis. Bond order is calculated using the number of electrons in bonding and antibonding orbitals and indicates bond strength, with higher bond order meaning stronger, more stable bonds.
Crystal Field Theory explains the colors of transition metal complexes based on ligand-metal interactions. The electrostatic interaction between ligands and metal d-orbitals splits the d-orbital energies. For an octahedral complex, the d-orbitals point directly at ligands have higher energy than those that bisect ligands. This splitting pattern determines if the complex is high or low spin, which then dictates its color and magnetic properties. The spectrochemical series orders ligands by their ability to cause crystal field splitting, correlating ligand type with complex color.
Slater's rules provide a method to estimate the shielding of electrons and the effective nuclear charge experienced by electrons in an atom. The rules involve writing the electron configuration, ignoring higher energy level electrons, and applying shielding constants of 0.35 for electrons in the same subshell and 0.85 for electrons in the previous subshell. As an example, the rules are used to calculate that the effective nuclear charge experienced by the valence electrons of nitrogen is 3.9 instead of the actual nuclear charge of 7.
Molecular orbital theory(mot) of SF6/CO2/I3-/B2H6sirakash
1) Molecular orbital theory views a molecule as delocalized molecular orbitals formed from linear combinations of atomic orbitals. Bonding molecular orbitals are lower in energy due to constructive interference, while antibonding orbitals are higher in energy due to destructive interference.
2) The document provides examples of applying molecular orbital theory to SF6, CO2, B2H6, and I3- molecules. It describes the atomic orbitals and molecular orbitals formed, including bonding, antibonding, and non-bonding orbitals, and explains how the molecular orbitals rationalize the electronic structures and bonding patterns in these molecules.
This document summarizes the Huckel molecular orbital theory. It describes the theory's key postulates for simplifying calculations for pi-electron systems like ethylene. The postulates state that overlap integrals are zero, coulomb integrals are equal, and exchange integrals are non-zero only for adjacent atoms. For ethylene, the HMO calculations yield two energy levels - a bonding and antibonding level. The coefficients and electron density are also calculated for ethylene's bonding orbital. Finally, the bond order and free valence are determined, showing ethylene has one pi-bond and equal reactivity at both carbon atoms.
The document discusses Crystal Field Theory, which explains the bonding in transition metal complexes. It describes how the electrostatic interaction between ligand electrons and metal d-orbitals results in a splitting of the d-orbital energies. In an octahedral field, the t2g orbitals are stabilized more than the eg orbitals. Crystal Field Theory can explain properties like electronic spectra, magnetic moments, and color of complexes. The magnitude of splitting depends on factors like the metal ion, its charge, the ligands, and can be represented by the crystal field splitting energy Δo.
This document provides an overview of metal carbonyls. It discusses how metal carbonyls are formed from transition metals and carbon monoxide, and examples like nickel tetracarbonyl and iron pentacarbonyl. The molecular orbital diagram of carbon monoxide is shown, explaining why it can participate in pi-backbonding. Infrared spectroscopy is described as a useful technique for analyzing metal carbonyls, as it can distinguish terminal from bridging carbonyl ligands based on the infrared absorption frequency. Factors like metal charge and other ligands that affect the carbonyl stretching frequency are also outlined. Finally, some applications of infrared spectra of metal carbonyls are mentioned.
- The document discusses molecular orbital theory, which describes chemical bonding through the combination of atomic orbitals into molecular orbitals.
- Key features include molecular orbitals being formed from linear combinations of atomic orbitals, with bonding, antibonding, and nonbonding molecular orbitals resulting. Electrons fill these orbitals based on orbital energy.
- The formation of molecular orbitals from atomic orbitals of hydrogen is used as an example, with bonding and antibonding molecular orbitals illustrated.
Metal nitrosyl compounds contain nitric oxide bonded as an NO+ ion, NO- ion, or neutral NO molecule. They can be classified into three classes based on the nitric oxide group present. Metal nitrosyls are coordination compounds where an NO molecule is attached as an NO+ ion to a metal atom or ion. Examples include metal nitrosyl carbonyls such as Co(NO+)(CO)3, metal nitrosyl halides such as Fe(NO+)2I, and metal nitrosyl thio-complexes involving Fe, Co, and Ni metals. These compounds can be prepared through the reaction of nitric oxide with metal compounds like carbonyls, halides, or ferrocyanides. Metal
Crystal field theory proposes that ligands behave as point charges that create an electric field around a central metal ion. This affects the energies of the metal's d-orbitals. In an octahedral complex, ligands along the x, y, and z axes interact more strongly with the dz2 and dx2-y2 orbitals, splitting them into the higher-energy eg set. The dxy, dyz, and dxz orbitals interact less with ligands between the axes, forming the lower-energy t2g set. This splitting of orbital energies, described by the crystal field splitting parameter Δ0, helps explain differences in complexes' magnetic properties.
The document summarizes aromaticity and related topics for chemistry students. It discusses:
- Benzenoid and non-benzenoid aromatic compounds, including their properties and reactions.
- Resonance structures of benzene and how it follows Huckel's rule for aromaticity.
- Classification of compounds based on aromaticity and examples of antiaromatic compounds.
- Aromatic ions and heterocyclic aromatic compounds like pyrrole, furan and pyridine.
The document discusses molecular orbital theory and its application to transition metal complexes. It describes how atomic orbitals of matching symmetry combine to form molecular orbitals, with equal numbers of bonding and antibonding orbitals. Electrons fill the molecular orbitals starting with the lowest energy orbitals. Ligand interactions such as π-accepting and π-donating affect the splitting of orbitals and influence the metal's oxidation state.
B.Sc. Sem-I Unit-IV Aromatic, antiaromatic and non aromatic compounds by Dr P...pramod padole
This document discusses aromaticity and Huckel's rule. It begins by explaining Huckel's rule, which states that cyclic compounds with (4n+2) pi electrons are aromatic, where n is an integer. Examples of aromatic compounds that satisfy Huckel's rule like benzene and naphthalene are given. The characteristics of aromatic compounds such as being cyclic, planar, and having delocalized pi electrons are described. Antiaromatic compounds that have 4n pi electrons are also discussed along with examples like cyclobutadiene. The document concludes by defining non-aromatic compounds that are neither aromatic nor antiaromatic, using cyclooctatetraene as an example.
Phosphine as ligand by Dr Geeta TewariGeeta Tewari
This presentation describes about the nature of phosphine ligands, bonding and reactions of metal phosphine containing complexes. Also explains the similarity and differences of phosphine ligand with NH3 and CO ligands.
1. Organometallic compounds are defined as compounds in which carbon atoms are directly bonded to metal atoms. Organometallic compounds can be classified as ionic, covalent, or electron deficient based on the nature of the metal-carbon bond.
2. Metal carbonyls are compounds where carbon monoxide ligands are bonded directly to metal centers. They can be mononuclear, containing one metal center, or polynuclear, containing multiple metal centers. Common mononuclear metal carbonyls prepared include nickel tetracarbonyl, iron pentacarbonyl, and chromium hexacarbonyl.
3. The document provides details on the nomenclature, classification, preparation, and structures
Crystal field theory and ligand field theory describe how ligands interact with transition metal complexes. Crystal field theory uses an electrostatic model to explain orbital splitting, while ligand field theory uses a molecular orbital approach. Both theories predict that ligands cause the d orbitals on the metal to split into lower energy t2g and higher energy eg sets. The size of this splitting depends on whether ligands are σ-donors only, π-donors, or π-acceptors. π-Acceptors increase splitting while π-donors decrease it. This explains the spectrochemical series from weak to strong field ligands.
Methods of Determining Reaction Mechanisms - Andria D'SouzaBebeto G
This document discusses various methods used to determine reaction mechanisms, including isotopic labeling techniques, stereochemical evidence, crossover experiments, identification of products, and identification of reaction intermediates. Isotopic labeling involves replacing atoms with isotopes like deuterium or carbon-13 to follow the reaction path. Stereochemical evidence from chiral reactants and products can indicate SN1 or SN2 mechanisms. Crossover experiments use non-identical reactants to study intramolecular vs intermolecular rearrangements. Identification of products and intermediates through spectroscopy, isolation, or trapping with reagents provides clues about reaction steps.
This document discusses aromaticity, including its introduction, criteria for aromatic compounds, Hückel's rule, examples of aromatic and anti-aromatic compounds, and non-aromatic compounds. Aromatic compounds are cyclic, planar, and have delocalized pi electrons that follow Hückel's rule of 4n+2 pi electrons. Benzene is used to originally define aromaticity. Resonance contributes greatly to aromatic stability. Anti-aromatic compounds have 4n pi electrons and are destabilized by cyclic pi electron delocalization. Cyclooctatetraene is provided as an example of a non-aromatic compound for not being planar.
Orgel diagrams; D and F/P Orgel Diagrams AafiaAslam
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.
The document summarizes Valence Shell Electron Pair Repulsion (VSEPR) Theory. [1] VSEPR Theory predicts molecular geometry based on electron pair repulsion around a central atom. [2] The theory states that electron pairs around an atom will position themselves as far apart as possible to minimize repulsion. [3] Molecular geometry can be determined by counting the number of electron pairs and their arrangement around the central atom.
This document provides an overview of molecular orbital theory. It explains that molecular orbital theory describes molecules in terms of orbitals and electron configurations similar to atomic orbital theory. The key points are:
- Molecular orbitals are formed from the overlapping of atomic orbitals on different atoms.
- Bonding orbitals are formed from constructive interference and lower the energy, while antibonding orbitals are formed from destructive interference and increase the energy.
- Homonuclear diatomic molecules like H2, O2, and N2 are discussed as examples, with their molecular orbitals, bond orders, and magnetic properties explained.
This document summarizes Crystal Field Theory, which considers the electrostatic interactions between metal ions and ligands. It describes ligands and metal ions as point charges that can have attractive or repulsive forces. This causes the d orbitals of the metal ion to split into two sets depending on if the field created by the ligands is weak or strong. The theory explains color in coordination compounds as being caused by d-d electron transitions under the influence of ligands. However, it has limitations like not accounting for other metal orbitals or the partial covalent nature of metal-ligand bonds.
Slater's rules provide a method to calculate the effective nuclear charge (Zeff) experienced by electrons in atoms and ions. The rules account for shielding of the nuclear charge by inner electrons. Zeff is calculated as the nuclear charge (Z) minus the total shielding (S). S is the sum of shielding values assigned based on orbital type and number of electrons. Comparing Zeff values explains trends like orbital filling order and which electrons are lost in cation formation. However, Slater grouped s and p orbitals together, which is incorrect as s orbitals penetrate the nucleus more than p orbitals.
This document summarizes the Huckel molecular orbital theory. It describes the theory's key postulates for simplifying calculations for pi-electron systems like ethylene. The postulates state that overlap integrals are zero, coulomb integrals are equal, and exchange integrals are non-zero only for adjacent atoms. For ethylene, the HMO calculations yield two energy levels - a bonding and antibonding level. The coefficients and electron density are also calculated for ethylene's bonding orbital. Finally, the bond order and free valence are determined, showing ethylene has one pi-bond and equal reactivity at both carbon atoms.
The document discusses Crystal Field Theory, which explains the bonding in transition metal complexes. It describes how the electrostatic interaction between ligand electrons and metal d-orbitals results in a splitting of the d-orbital energies. In an octahedral field, the t2g orbitals are stabilized more than the eg orbitals. Crystal Field Theory can explain properties like electronic spectra, magnetic moments, and color of complexes. The magnitude of splitting depends on factors like the metal ion, its charge, the ligands, and can be represented by the crystal field splitting energy Δo.
This document provides an overview of metal carbonyls. It discusses how metal carbonyls are formed from transition metals and carbon monoxide, and examples like nickel tetracarbonyl and iron pentacarbonyl. The molecular orbital diagram of carbon monoxide is shown, explaining why it can participate in pi-backbonding. Infrared spectroscopy is described as a useful technique for analyzing metal carbonyls, as it can distinguish terminal from bridging carbonyl ligands based on the infrared absorption frequency. Factors like metal charge and other ligands that affect the carbonyl stretching frequency are also outlined. Finally, some applications of infrared spectra of metal carbonyls are mentioned.
- The document discusses molecular orbital theory, which describes chemical bonding through the combination of atomic orbitals into molecular orbitals.
- Key features include molecular orbitals being formed from linear combinations of atomic orbitals, with bonding, antibonding, and nonbonding molecular orbitals resulting. Electrons fill these orbitals based on orbital energy.
- The formation of molecular orbitals from atomic orbitals of hydrogen is used as an example, with bonding and antibonding molecular orbitals illustrated.
Metal nitrosyl compounds contain nitric oxide bonded as an NO+ ion, NO- ion, or neutral NO molecule. They can be classified into three classes based on the nitric oxide group present. Metal nitrosyls are coordination compounds where an NO molecule is attached as an NO+ ion to a metal atom or ion. Examples include metal nitrosyl carbonyls such as Co(NO+)(CO)3, metal nitrosyl halides such as Fe(NO+)2I, and metal nitrosyl thio-complexes involving Fe, Co, and Ni metals. These compounds can be prepared through the reaction of nitric oxide with metal compounds like carbonyls, halides, or ferrocyanides. Metal
Crystal field theory proposes that ligands behave as point charges that create an electric field around a central metal ion. This affects the energies of the metal's d-orbitals. In an octahedral complex, ligands along the x, y, and z axes interact more strongly with the dz2 and dx2-y2 orbitals, splitting them into the higher-energy eg set. The dxy, dyz, and dxz orbitals interact less with ligands between the axes, forming the lower-energy t2g set. This splitting of orbital energies, described by the crystal field splitting parameter Δ0, helps explain differences in complexes' magnetic properties.
The document summarizes aromaticity and related topics for chemistry students. It discusses:
- Benzenoid and non-benzenoid aromatic compounds, including their properties and reactions.
- Resonance structures of benzene and how it follows Huckel's rule for aromaticity.
- Classification of compounds based on aromaticity and examples of antiaromatic compounds.
- Aromatic ions and heterocyclic aromatic compounds like pyrrole, furan and pyridine.
The document discusses molecular orbital theory and its application to transition metal complexes. It describes how atomic orbitals of matching symmetry combine to form molecular orbitals, with equal numbers of bonding and antibonding orbitals. Electrons fill the molecular orbitals starting with the lowest energy orbitals. Ligand interactions such as π-accepting and π-donating affect the splitting of orbitals and influence the metal's oxidation state.
B.Sc. Sem-I Unit-IV Aromatic, antiaromatic and non aromatic compounds by Dr P...pramod padole
This document discusses aromaticity and Huckel's rule. It begins by explaining Huckel's rule, which states that cyclic compounds with (4n+2) pi electrons are aromatic, where n is an integer. Examples of aromatic compounds that satisfy Huckel's rule like benzene and naphthalene are given. The characteristics of aromatic compounds such as being cyclic, planar, and having delocalized pi electrons are described. Antiaromatic compounds that have 4n pi electrons are also discussed along with examples like cyclobutadiene. The document concludes by defining non-aromatic compounds that are neither aromatic nor antiaromatic, using cyclooctatetraene as an example.
Phosphine as ligand by Dr Geeta TewariGeeta Tewari
This presentation describes about the nature of phosphine ligands, bonding and reactions of metal phosphine containing complexes. Also explains the similarity and differences of phosphine ligand with NH3 and CO ligands.
1. Organometallic compounds are defined as compounds in which carbon atoms are directly bonded to metal atoms. Organometallic compounds can be classified as ionic, covalent, or electron deficient based on the nature of the metal-carbon bond.
2. Metal carbonyls are compounds where carbon monoxide ligands are bonded directly to metal centers. They can be mononuclear, containing one metal center, or polynuclear, containing multiple metal centers. Common mononuclear metal carbonyls prepared include nickel tetracarbonyl, iron pentacarbonyl, and chromium hexacarbonyl.
3. The document provides details on the nomenclature, classification, preparation, and structures
Crystal field theory and ligand field theory describe how ligands interact with transition metal complexes. Crystal field theory uses an electrostatic model to explain orbital splitting, while ligand field theory uses a molecular orbital approach. Both theories predict that ligands cause the d orbitals on the metal to split into lower energy t2g and higher energy eg sets. The size of this splitting depends on whether ligands are σ-donors only, π-donors, or π-acceptors. π-Acceptors increase splitting while π-donors decrease it. This explains the spectrochemical series from weak to strong field ligands.
Methods of Determining Reaction Mechanisms - Andria D'SouzaBebeto G
This document discusses various methods used to determine reaction mechanisms, including isotopic labeling techniques, stereochemical evidence, crossover experiments, identification of products, and identification of reaction intermediates. Isotopic labeling involves replacing atoms with isotopes like deuterium or carbon-13 to follow the reaction path. Stereochemical evidence from chiral reactants and products can indicate SN1 or SN2 mechanisms. Crossover experiments use non-identical reactants to study intramolecular vs intermolecular rearrangements. Identification of products and intermediates through spectroscopy, isolation, or trapping with reagents provides clues about reaction steps.
This document discusses aromaticity, including its introduction, criteria for aromatic compounds, Hückel's rule, examples of aromatic and anti-aromatic compounds, and non-aromatic compounds. Aromatic compounds are cyclic, planar, and have delocalized pi electrons that follow Hückel's rule of 4n+2 pi electrons. Benzene is used to originally define aromaticity. Resonance contributes greatly to aromatic stability. Anti-aromatic compounds have 4n pi electrons and are destabilized by cyclic pi electron delocalization. Cyclooctatetraene is provided as an example of a non-aromatic compound for not being planar.
Orgel diagrams; D and F/P Orgel Diagrams AafiaAslam
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.
The document summarizes Valence Shell Electron Pair Repulsion (VSEPR) Theory. [1] VSEPR Theory predicts molecular geometry based on electron pair repulsion around a central atom. [2] The theory states that electron pairs around an atom will position themselves as far apart as possible to minimize repulsion. [3] Molecular geometry can be determined by counting the number of electron pairs and their arrangement around the central atom.
This document provides an overview of molecular orbital theory. It explains that molecular orbital theory describes molecules in terms of orbitals and electron configurations similar to atomic orbital theory. The key points are:
- Molecular orbitals are formed from the overlapping of atomic orbitals on different atoms.
- Bonding orbitals are formed from constructive interference and lower the energy, while antibonding orbitals are formed from destructive interference and increase the energy.
- Homonuclear diatomic molecules like H2, O2, and N2 are discussed as examples, with their molecular orbitals, bond orders, and magnetic properties explained.
This document summarizes Crystal Field Theory, which considers the electrostatic interactions between metal ions and ligands. It describes ligands and metal ions as point charges that can have attractive or repulsive forces. This causes the d orbitals of the metal ion to split into two sets depending on if the field created by the ligands is weak or strong. The theory explains color in coordination compounds as being caused by d-d electron transitions under the influence of ligands. However, it has limitations like not accounting for other metal orbitals or the partial covalent nature of metal-ligand bonds.
Slater's rules provide a method to calculate the effective nuclear charge (Zeff) experienced by electrons in atoms and ions. The rules account for shielding of the nuclear charge by inner electrons. Zeff is calculated as the nuclear charge (Z) minus the total shielding (S). S is the sum of shielding values assigned based on orbital type and number of electrons. Comparing Zeff values explains trends like orbital filling order and which electrons are lost in cation formation. However, Slater grouped s and p orbitals together, which is incorrect as s orbitals penetrate the nucleus more than p orbitals.
The document discusses electron shielding and how to calculate the effective nuclear charge (Zeff) experienced by an electron. Electrons closer to the nucleus provide more shielding from the nuclear charge than outer electrons. Rules are provided to calculate the shielding constant (S) based on the electronic configuration and determine Zeff by subtracting S from the atomic number (Z). Examples demonstrate applying the rules to calculate Zeff for valence and inner electrons of various elements.
This document discusses atomic structure and properties. It begins by defining the three main subatomic particles - protons, neutrons, and electrons - and their relative masses and charges. It then explains atomic mass units and discusses quantum numbers like principal, azimuthal, magnetic, and spin quantum numbers. The rest of the document covers electron configurations, orbital diagrams, atomic orbitals and subshells, ionization energy, electronegativity, and how these properties relate to an element's position on the periodic table.
The nuclear shell model can explain the observed magic numbers of nuclei by considering the quantized energy levels of nucleons moving in a spherically symmetric nuclear potential. However, it fails to accurately predict properties like binding energies, magnetic dipole moments, and excited state spectra. The collective model improves on this by allowing for a non-spherical nuclear potential, leading to nuclear deformation and rotational excited states. This explains phenomena like large electric quadrupole moments observed in heavier nuclei and the characteristic rotational band structure seen in their low-lying excited states.
The document discusses periodic trends in effective nuclear charge (Zeff) and how it relates to ionization energy. It provides an example calculating the Zeff of the outermost electrons as argon is sequentially ionized from Ar+ to Ar18+. With each additional ionization, an electron is removed and the Zeff of the remaining electrons increases. This leads to higher ionization energies due to the greater attraction felt by the electrons to the increasingly exposed positive nuclear charge.
This document provides an overview of crystal field theory and how it can be used to explain the bonding and spectroscopic properties of transition metal complexes. It discusses how ligands arranged in octahedral, tetrahedral and square planar geometries cause the d-orbitals of the transition metal to split into different energy levels. Factors that influence the size of the crystal field splitting parameter Δo, such as oxidation state, metal identity and ligand type, are also covered.
The document summarizes key concepts about the periodic table including its organization based on electron configuration and blocks (s, p, d, f), rules for filling orbitals like the Pauli exclusion principle and Hund's rule, and examples of writing electronic configurations for elements like sodium, neon, aluminum, argon, and potassium.
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.
NANO106 is UCSD Department of NanoEngineering's core course on crystallography of materials taught by Prof Shyue Ping Ong. For more information, visit the course wiki at http://nano106.wikispaces.com.
This document discusses the magnetic properties of transition metal ions and their coordination complexes. It begins by explaining how transition metal ions have unpaired electrons that contribute to their paramagnetism. It then shows calculations of the spin-only magnetic moment (μS) , total angular momentum magnetic moment (μL+S) , and the effective magnetic moment (μeff) for Cr3+ as an example. For most transition metal ions, the observed μeff is close to the calculated μS. Exceptions are Co2+ and Ni2+ which have higher observed μeff due to spin-orbit coupling. The document also explains how crystal field theory can predict if a complex will be high or low spin based on the ligand's splitting
A Study on Atomic Spectroscopic Term Symbols for Nonequivalent Electrons of (...IJERA Editor
Electronic and magnetic properties of the inorganic molecules and complexes can be understood very well by using the term symbols. Term symbols are a shorthand method used to describe the energy, angular momentum, and spin multiplicity of an atom / ion in any particular electronic state. The Russell-Saunders atomic term symbols provide the information about spectral and magnetic properties of an atom or ion. The atomic terms have been determined for nonequivalent electrons of(n-1)d 1 s 1 p 1 configuration using Russell-Saunders coupling scheme.The total number of microstates computed for the non-equivalent electrons of (n-1)d 1 s 1 p 1 configuration are found to be 120. Among the microstates, two types of atomic terms have been determined namely, quartet (3-types) and doublet (6-types). The ground state term found for this configuration is quartet 4 F with the lower energy state as 4 F2
The document discusses the structure of atoms, including:
1) Solving the Schrodinger equation for hydrogen-like atoms to determine allowed energies and wavefunctions.
2) The quantization of energy levels, orbital angular momentum, and other properties for hydrogen-like atoms.
3) How the concepts for hydrogen-like atoms can be applied to describe multi-electron atoms and molecules using approximations like the orbital model.
The document discusses the concept of effective nuclear charge. It explains that the actual charge experienced by valence electrons is less than the true nuclear charge due to shielding by inner electrons. This decreased charge is called the effective nuclear charge (Zeff). Slater's rules provide a method to calculate the screening constant σ and thus determine Zeff. The concept of Zeff is applied to explain trends in ionization energy, filling of electron shells, and properties of cations, anions, and across the periodic table.
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.
1) Radioactivity is the spontaneous emission of radiation by unstable atomic nuclei. It occurs as the nucleus shifts to a more stable configuration by emitting energy.
2) The principal factor determining nuclear stability is the neutron-to-proton ratio. No nucleus larger than lead-208 is stable as the strong force cannot overcome electrostatic repulsion at larger sizes.
3) The rate of radioactive decay is proportional to the number of nuclei present and follows an exponential decay model expressed as N(t)=N0e-λt, where λ is the decay constant and N0 is the initial number of nuclei.
This document discusses key concepts in quantum mechanics including:
- Planck's quantum theory which established that atoms can only emit or absorb energy in discrete quanta.
- Einstein's explanation of the photoelectric effect using the particle nature of light (photons).
- Bohr's model of the hydrogen atom which explained its spectral lines by postulating discrete electron energy levels.
- Quantum numbers which describe the state of an electron including its orbital, orientation, and spin.
- Electron configuration which shows how electrons fill atomic orbitals according to the aufbau principle.
This document discusses quantum numbers and their role in describing the size, shape, and orientation of atomic orbitals. It explains that there are four quantum numbers - principal, angular, magnetic, and spin. The principal quantum number determines the electron shell or energy level, while the angular and magnetic quantum numbers further specify the subshell and orbital within that subshell. The spin quantum number refers to the spin of the electron. Factors that influence ionization energy such as atomic radius, nuclear charge, and electron shielding are also summarized.
The four quantum numbers - principal (n), azimuthal (l), magnetic (m), and spin (s) - are used to completely characterize electrons in an atom. The principal quantum number indicates the electron's energy level. The azimuthal number denotes the subshell and shape of its orbital. The magnetic number determines the number of orbitals in a subshell. The spin number arises from the electron's intrinsic spin and can have values of ±1/2. Together, the four quantum numbers uniquely identify each electron in an atom.
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1. The document discusses the quantum theory of atomic structure, including quantum numbers, electron configurations, and atomic orbitals. It explains how the principal quantum number n determines an electron's energy level and average distance from the nucleus, while the angular momentum quantum number l determines the shape of its orbital.
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2. Slater’s Rule
INTRODUCTION:
In 1930, a scientist J.C. Slater proposed a set of
empirical rules to understand the concept of Effective
Nuclear Charge and to calculate the Screening
Constant or Shielding Constant.
He proposed a formula for calculation of Effective
Nuclear Charge
Zeff = Z – S
where S is the Slater’s screening constant ,
Z is the Nuclear charge
3. Prior to explaining Slater’s rules , certain terms like
Nuclear Charge, Shielding Effect and Effective Nuclear
Charge have to be Understood.
What is Nuclear Charge?
It is the charge on the nucleus with which it attracts the
electrons of the atom. Basically, the Nuclear Charge is said
to be equal to the Atomic Number( i.e, . the Number of
protons) in an atom. It is denoted by the symbol Z.
What is Shielding Effect?
In case of Multielectron atoms, as the orbitals are filled
up, the electrons in the Inner orbitals shield the electrons in
the Outer orbitals from the Nucleus.
4. So, the electrons in the Outer orbitals do not feel the full
force or charge of the nucleus. Thus, the reduction of
nuclear charge on the Outermost electrons is called
Shielding. Effect or Screening Effect. Shielding Effect is
defined as a measure of the extent to which the intervening
electrons shield the outer electrons from the nuclear charge.
It is denoted by the symbol S
What is Effective Nuclear Charge ?
Effective Nuclear Charge is the actual charge felt by
the outer electrons after taking into account shielding of
the electrons. It is denoted by the symbol Z* or Zeff
5. Slater's Rules:
1) Write the electron configuration for the atom using
the following design;
(1s)(2s,2p)(3s,3p) (3d) (4s,4p) (4d) (4f) (5s,5p) (5d)
( 5f) (6s,6p)……. etc.
2) Any electrons to the right of the electron of interest
contributes nothing towards shielding.
3) All other electrons in the same group as the
electron of interest shield to an extent of 0.35
nuclear charge units irrespective of whether the
electrons are in s, p, d, or f orbitals.
6. 4) In case of 1s electron shielding another 1s electron the
screening constant value is taken to be 0.30.
5) If the electron of interest is an s or p electron: All electrons
with one less value i.e. (n -1) value of the principal
quantum number shield to an extent of 0.85 units of nuclear
charge. All electrons with two or more less values i.e.
( n – 2, n – 3, n – 4 etc.) values of the principal quantum
number shield to an extent of 1.00 units.
6) If the electron of interest is an d or f electron: All electrons to
the left shield to an extent of 1.00 units of nuclear charge.
7) Sum the shielding amounts from steps 2 through 5 and
subtract from the nuclear charge value to obtain the effective
nuclear charge value.
7. Calculate Z* for a valence electron in
fluorine ( Z = 9 ).
Electronic configuration of fluorine is 1s2,2s2,2p5
Grouping it acc. to slater’s rule : (1s2)(2s2,2p5)
Rule 2 does not apply;
Now, one electron out of the 7 valence electrons
becomes the electron of interest. The other remaining
6 valence electrons will contribute 0.35 each towards
shielding.
The electrons in (n – 1) orbitals i.e. 1s orbital will
contribute 0.85 each towards shielding.
8. S = 0.35 x ( No. of electrons in the same shell i.e.
n orbital ) + 0.85 x (No. of electrons in the
(n – 1) shell )
S = 0.35 x 6 + 0.85 x 2 = 3.8
Z* = Z – S = 9 – 3.8 = 5.2 for a valence electron.
Calculate Z* for a 6s electron in Platinum
( Z = 78 )
9. The electronic configuration of Platinum ( Z = 78 ) is
1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d10, 4p6, 5s2, 4d10, 5p6,
6s2, 4f14 , 5d8 ,
.
Grouping it acc. to Slater’s rule :
(1s2)(2s2,2p6)(3s2,3p6) (3d10) (4s2,4p6) (4d10) (4f14)
(5s2,5p6) (5d8) (6s2)
Rule 2 does not apply;
Now, one electron out of the two valence electrons
becomes the electron of interest. The other
remaining valence electrons will contribute 0.35
each towards shielding.
10. The electrons in (n – 1) orbitals i.e. ( 6 – 1 ) = 5th orbital will
contribute 0.85 each towards shielding.
The electrons in (n – 2) , ( n – 3), ( n – 4) ….etc. orbitals
{ i.e. ( 6 – 2 ) , ( 6 – 3), (6 – 4)……etc.} orbitals will
contribute 1.00 each towards shielding.
S = 0.35 x ( No. of electrons in the same shell i.e.
n orbital ) + 0.85 x (No. of electrons in the
(n – 1) shell ) + 1.00 x ( No. of electrons in the (n – 2),
( n – 3)….etc. orbitals )
S = 0.35 x 1 + 0.85 x 16 + 60 x 1.00 = 73.95
Z* = Z – S = 78 – 73.95 = 4.15 for a valence electron.
11. Calculate the Effective Nuclear Charge for
one of the 4f electrons of Cerium ( Z = 58 )
The electronic configuration of Cerium is :
1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 4d10 4f2 5s2 5p6 6s2
Grouping the orbitals acc. to Slater’s :
(1s2) (2s2 2p6) (3s2 3p6) (3d10) (4s2 4p6) (4d10) (4f2)
(5s2 5p6) (6s2)
As we have to calculate the effective nuclear charge
of 4f electrons of Cerium, the electrons lying after the
4f electron will not contribute to shielding
12. Now,
S = 0.35 x ( No. of electrons in the same orbital +
1.00 x( all the electrons in the lower orbitals)
S = 0.35 x 1 + 1.00 x 46 = 46.35
Zeff = Z – S = 58 – 46.35 = 11.65
Calculate the Effective Nuclear Charge in
the periphery of Nitrogen ( Z = 7)
In order to calculate the Effective Nuclear Charge in the
periphery of an atom or ion, the shielding of nuclear
charge by all the electrons present in the electronic
configuration of the atom or ion.
13. The electronic configuration of Nitrogen is 1s2 2s2 2p6
Grouping acc. to Slater’s rule : (1s2) (2s2 2p6)
S = 0.35 x ( No. of electrons in the same orbital ) +
0.85 x (No. of electrons in the (n – 1) orbital
S = 0.35 x 5 + 0.85 x 2 = 3.45
Zeff in the periphery of N - atom = Z – S = 7 – 3.45
= 3.55
14. APPLICATIONS OF SLATER’S RULE
It provides a quantitative justification for the
sequence of orbitals in the energy level diagram.
It helps to explain the filling of ns - orbital (4s, 5s,6s
etc. – orbitals)prior to the filling of (n -1)d orbital (3d,
4d, 5d…etc.).
Let us consider the case of Potassium ( Z = 19), in
which the last electron is added to 4s orbital
15. The configuration of Potassium acc. to Slater is
(1s2 ) ( 2s2 2p6 ) ( 3s2 3p6 ) ( 4s1 )
As the effective nuclear charge on electron in 4s
orbital has to be calculated , the electrons in the
same orbital i.e. n orbital will contribute 0.35 each,
the electrons in (n – 1) orbital i.e. 3s and 3p
orbitals will contribute S = 0.85 each and all the
electrons in (n – 2, n – 3…. Etc ) orbitals i.e. ( 2s ,
2p, 1s) orbitals will contribute S =1.00 each.
16. So,
S = 0 x 0.35 + 8 x 0.85 + 10 x 1.00
= 16.80
Therefore, Effective Nuclear Charge
Z* = Z – S = 19 – 16.80 = 2.20
Now, Let us assume that the last electron
enters the 3d orbital rather than 4s orbital,
Then the configuration acc. to Slater is
(1s2 ) ( 2s2 2p6 ) ( 3s2 3p6 ) ( 3d1 )
17. Here the d electron is under Interest, so the electrons
in the same orbital i.e. 3d orbital will contribute S =
0.35 each, where as the electrons in all the other
orbitals i.e. (3s,3p,2s, 2p, 1s) will all contribute S =
1.00 each.
So,
S = 0 x 0.35 + 18 x 1.00 = 18.00
Therefore Effective Nuclear Charge
Z* = Z – S = 19.00 – 18.00 = 1.00
On comparing the Effective Nuclear Charge of both 4s
and 3d orbitals, we see that the 4s electron is under the
influence of greater Effective Nuclear charge (Zeff = 2.20 )
as compared to 3d electron (Zeff = 1.00) in Potassium
atom.
18. So, the electron in 4s orbital will be more attracted by
the nucleus and will have lower energy than the 3d
electron.
Thus, the last electron will enter in the 4s orbital , rather
than the 3d orbital in case of Potassium atom.
Slater’s rule explain why 4s electrons are lost prior
to 3d electrons during cation formation in case of
Transition elements.
Let us consider the case of Vanadium (Z=23)
The electronic configuration of Vanadium is
1s2 2s2 2p6 3s2 3p6 3d3 4s2
19. After losing 2 electrons, the electronic configuration of
V2+ is
1s2 2s2 2p6 3s2 3p6 3d1 4s2
And not
1s2 2s2 2p6 3s2 3p6 3d3
The above Electronic Configuration can be explained by
Slater’s rules
The Effective Nuclear Charge for 4s electron is
calculated as:
(1s2) (2s2 2p6) (3s2 3p6) (3d3) (4s2)
Now, one of the electrons of the 4s orbital becomes
electron of interest. The second electron however, will
contribute towards shielding effect
20. S = 1 x 0.35 + 11 x 0.85 + 10 x 1.00 = 19.70
So, Zeff = Z – S = 23 – 19.70 = 3.30
Now, lets calculate the Effective Nuclear Charge
for a 3d electron
According to Slater’s rule
(1s2) (2s2 2p6 ) (3s2 3p6 ) (3d3)(4s2)
As 4s orbital lies after the electron under interest it will
contributes nothing towards shielding.
One electron of 3d orbital becomes electron of interest. The
other two 3d electrons will contribute towards shielding
effect.
S = 2 x 0.35 + 18 x 1.00 = 18.70
Zeff = Z – S = 23 – 18.70 = 4.30
21. Comparing the Effective Nuclear Charge of both the
3d and the 4s electron, it is seen that Effective Nuclear
Charge on 3d electron is 4.30 whereas 4s electron has
3.30.
The force of attraction experienced by the 3d
electrons is more as compared to 4s electrons.
The 3d electrons are more tightly held to the nucleus
than 4s electrons.
Thus, the 4s electrons are removed in preference to 3d
electrons.
22. The force of attraction experienced by the 3d
electrons is more as compared to 4s electrons.
The 3d electrons are more tightly held to the nucleus
than
4s electrons.
Thus, the 4s electrons are removed in preference to 3d
electrons.
It helps to explain why size of a cation is
always smaller than its neutral atom.
Let’s take the example of Lithium atom and Lithium
ion
The Electronic configuration of Lithium atom is 1s2 2s1
Write it according to Slater’s (1s2) (2s1)
As 2s orbital has only one electron , it becomes the
electron of interest.
23. Only the 1s electrons will contribute towards shielding.
S = 2 x 0.85 = 1.70
Z eff = Z – S = 3 – 1.70 = 1.30
In case of Lithium (Li+) Ion
The Electronic Configuration of Li+ is 1s2
Grouping acc. to Slater : (1s2 )
In Case of Li+ , one of the 1s electrons becomes electron
of interest and the other 1s electron contributes towards
shielding.
S = 1 x 0.30 = 0.30
Z eff = Z – S = 3 – 0.30 = 2.70
24. Comparison of the Effective Nuclear Charge of Li atom
( Z eff = 1.30) and Li+ ion ( Z eff = 2.70), shows that
Effective nuclear charge of Li+ ion is more than Li atom.
So , the size of Li+ ion is smaller than Li atom.
It explains why a anion is always larger than its
neutral atom
Taking the example of Chlorine atom and Chlorine ion.
In case of Chlorine atom ( Z = 17), the electronic
configuration is 1s2 2s2 2p6 3s2 3p5
Grouping it acc. to Slater’s rule
(1s2 ) ( 2s2 2p6 ) ( 3s2 3p5)
S = 6 x 0.35 + 8 x 0.85 + 2 x 1.00 = 10.90
Zeff = Z – S = 17 – 10.90 = 6.10
25. In case of Chlorine ion Cl-
The electronic configuration is 1s2 2s2 2p6 3s2 3p6
Grouping it acc. to Slater’s rule
(1s2 ) ( 2s2 2p6 ) ( 3s2 3p6)
S = 7 x 0.35 + 8 x 0.85 + 2 x 1.00 = 11.25
Zeff = Z – S = 17 – 11.25 = 5.75
Comparison of the Effective Nuclear Charge of Cl atom
( Z eff = 6.10) and Cl- ion ( Z eff = 5.75), shows that Effective
nuclear charge on Cl atom is more than Cl- ion.
So , the size of Cl- ion is larger than Cl atom .
26. Slater grouped both s and p orbitals together for
calculating effective nuclear charge , which is
incorrect. This is because radial probability
distribution curves show that s orbitals are more
penetrating than p orbitals. So , the s orbitals
should shield to a greater extent as compared to
p orbital.
According to Slater , all the s ,p ,d and f electrons
present in shell or energy level lower than (n – 1)
shell will shield the outer n electrons with equal
contribution of S=1.00 each. This is not justified
as energetically different orbitals should not
contribute equally.
Slater’s rules are less reliable for heavier
elements
* LIMITATIONS OF SLATER’S RULE