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Chapter 11 - Theories of Covalent Bonding

  • A covalent bond develops when the orbitals of two atoms overlap and a pair of electrons occupy the overlap area, according to the basic premise of VB theory.

  • The overlap of the two orbitals indicates that their wave functions are in phase (constructive interference; as referenced to the image attached), and therefore the amplitude between the nuclei rises. This idea underpins the major themes of VB theory:

  • The electron pair is opposing spins.

  • The space produced by the overlapping orbitals has a maximum capacity for two electrons with opposing (paired) spins, as required by the Pauli exclusion principle.

  • A molecule of H2 develops in the simplest scenario when the 1s orbitals of two H atoms overlap and the electrons, with their spins coupled, spend more time in the overlap area (up and down arrows in the attached picture).

  • Bonding orbitals with the greatest overlap. Connection strength is determined by the attraction of nuclei to shared electrons; hence, the higher the orbital overlap, the closer the nuclei are to the electrons, and the stronger the bond.

  • The extent of overlap is determined by the orbital shape and orientation.

  • Atomic orbital hybridization To explain bonding in diatomic compounds such as HF and F2, we imagine a direct overlap of s and/or p orbitals of isolated atoms.

  • But how can the shapes and orientations of C and H atomic orbitals account for the structure of a molecule like methane? A-C atom ([He] 2s 2 2p2) contains two valence electrons in the spherical 2s orbital and one in two of the three mutually perpendicular 2p orbitals.

  • Two CH bonds with a 90° HCH bond angle would occur if the two half-filled p orbitals overlapped the 1s orbitals of two H atoms.

  • Methane, on the other hand, has the formula CH4, not CH2, and its bond angles are 109.5°. Linus Pauling proposed a theory to explain such phenomena.

  • Hybrid orbitals have certain characteristics. Here are some key facts to remember about the hybrid orbitals that develop during bonding:

  • The number of hybrid orbitals created matches the number of atomic orbitals mixed, and the kind of hybrid orbitals formed vary with the types of atomic orbitals combined.

  • A hybrid orbitals shape and orientation optimize its overlap with the orbital of the other atom in the bond.

  • It's helpful to conceive of hybridization as a process in which atomic orbitals mix, hybrid orbitals develop, they overlap other orbitals, and electrons with opposing spins join the overlap zone, establishing stable bonds.

  • In reality, hybridization is a mathematical notion that helps us understand the chemical world.

  • The VSEPR and VB theories are used to explain observed molecule shapes. However, in certain situations, the ideas may be inconsistent with other data.

  • Large nonmetal hydrides are not suitable for hybridization. Consider H2S's Lewis structure and bond angle:

  • According to VSEPR theory, the four-electron groups surrounding H2S point to the corners of a tetrahedron, and the two lone pairs compress the HSH bond angle below the ideal 109.5°.

  • According to VB theory, the S atom's 3s and 3p orbitals combine to create four sp3 hybrids, two of which are filled with lone pairs while the other two overlap 1s orbitals of two H atoms and are filled with bonding pairs.

  • Shapes with enlarged valence shells are less significant for d-Orbital hybridization. The creation of sp3 d and sp3 d2 hybrid orbitals with enlarged valence shells is proposed by the VB theory.

  • However, new quantum-mechanical simulations have revealed that d orbitals have such high energy that they do not successfully hybridize with the considerably more stable s and p orbitals of a given n number.

  • Thus, it now appears that SF6 does not bind using sp3 d2 hybrids; instead, some have hypothesized that SF6 is best to the stable when the bonding orbitals of the central S employ a combination of sp hybrid orbitals and unhybridized 3p orbitals.

  • Others favor molecular orbitals or even ionic structures as explanations. We must remember that molecules do not have to obey our models; rather, we must understand the limitations of our models and change them in light of new data.

  • Because VB theory successfully explains the molecular geometries of molecules with expanded valence shells, we will continue to use the traditional approach of including d-orbital hybridization for molecules with expanded valence shells in this text for simplicity, while acknowledging its limitations.

  • End-to-end overlap of atomic orbitals produces a connection, allowing unrestricted rotation of the molecule's linked components. A multiple bond is made up of a bond and either one (double bond) or two bonds (triple bond).

  • Multiple bonds contain more electron density between their nuclei than single bonds, resulting in higher bond energies. Rotation in a bond is limited by the side-to-side overlap of orbitals.

  • Molecular orbital (MO) theory considers a molecule to be a collection of nuclei with MOs distributed throughout the structure.

  • Atomic orbitals with equivalent energies can be combined or removed to produce bonding or antibonding MOs.

  • Bonding MOs, whether or, have most of the electron density between the nuclei and are lower in energy than the AOs that combine to create them; antibonding MOs have most of the electron density outside the nuclei and are thus greater in energy.

  • MOs are filled in descending order of energy by paired electrons with opposing spins. MO graphs depict energy levels as well as orbital occupancy. Period 2 homonuclear diatomic molecular diagrams

  • Several features of MO theory, including MOThe charge filling, energy-level diagrams, electron configuration, and bond order, are related to previous concepts: Electrons are injected into MOs. Electrons enter MOs in the same way as they do AOs:

  • MOs are filled in ascending energy order (Aufbau principle). An MO can only contain two electrons with opposing spins (Pauli exclusion principle).

  • Before any of the orbitals of equal energy are filled, they are half-filled, with electron spins parallel (Hund's rule).

  • A covalent bond develops when the orbitals of two atoms overlap and a pair of electrons occupy the overlap area, according to the basic premise of VB theory.

  • The overlap of the two orbitals indicates that their wave functions are in phase (constructive interference; as referenced to the image attached), and therefore the amplitude between the nuclei rises. This idea underpins the major themes of VB theory:

  • The electron pair is opposing spins.

  • The space produced by the overlapping orbitals has a maximum capacity for two electrons with opposing (paired) spins, as required by the Pauli exclusion principle.

  • A molecule of H2 develops in the simplest scenario when the 1s orbitals of two H atoms overlap and the electrons, with their spins coupled, spend more time in the overlap area (up and down arrows in the attached picture).

  • Bonding orbitals with the greatest overlap. Connection strength is determined by the attraction of nuclei to shared electrons; hence, the higher the orbital overlap, the closer the nuclei are to the electrons, and the stronger the bond.

  • The extent of overlap is determined by the orbital shape and orientation.

  • Atomic orbital hybridization To explain bonding in diatomic compounds such as HF and F2, we imagine a direct overlap of s and/or p orbitals of isolated atoms.

  • But how can the shapes and orientations of C and H atomic orbitals account for the structure of a molecule like methane? A-C atom ([He] 2s 2 2p2) contains two valence electrons in the spherical 2s orbital and one in two of the three mutually perpendicular 2p orbitals.

  • Two CH bonds with a 90° HCH bond angle would occur if the two half-filled p orbitals overlapped the 1s orbitals of two H atoms.

  • Methane, on the other hand, has the formula CH4, not CH2, and its bond angles are 109.5°. Linus Pauling proposed a theory to explain such phenomena.

  • Hybrid orbitals have certain characteristics. Here are some key facts to remember about the hybrid orbitals that develop during bonding:

  • The number of hybrid orbitals created matches the number of atomic orbitals mixed, and the kind of hybrid orbitals formed vary with the types of atomic orbitals combined.

  • A hybrid orbitals shape and orientation optimize its overlap with the orbital of the other atom in the bond.

  • It's helpful to conceive of hybridization as a process in which atomic orbitals mix, hybrid orbitals develop, they overlap other orbitals, and electrons with opposing spins join the overlap zone, establishing stable bonds.

  • In reality, hybridization is a mathematical notion that helps us understand the chemical world.

  • The VSEPR and VB theories are used to explain observed molecule shapes. However, in certain situations, the ideas may be inconsistent with other data.

  • Large nonmetal hydrides are not suitable for hybridization. Consider H2S's Lewis structure and bond angle:

  • According to VSEPR theory, the four-electron groups surrounding H2S point to the corners of a tetrahedron, and the two lone pairs compress the HSH bond angle below the ideal 109.5°.

  • According to VB theory, the S atom's 3s and 3p orbitals combine to create four sp3 hybrids, two of which are filled with lone pairs while the other two overlap 1s orbitals of two H atoms and are filled with bonding pairs.

  • Shapes with enlarged valence shells are less significant for d-Orbital hybridization. The creation of sp3 d and sp3 d2 hybrid orbitals with enlarged valence shells is proposed by the VB theory.

  • However, new quantum-mechanical simulations have revealed that d orbitals have such high energy that they do not successfully hybridize with the considerably more stable s and p orbitals of a given n number.

  • Thus, it now appears that SF6 does not bind using sp3 d2 hybrids; instead, some have hypothesized that SF6 is best to the stable when the bonding orbitals of the central S employ a combination of sp hybrid orbitals and unhybridized 3p orbitals.

  • Others favor molecular orbitals or even ionic structures as explanations. We must remember that molecules do not have to obey our models; rather, we must understand the limitations of our models and change them in light of new data.

  • Because VB theory successfully explains the molecular geometries of molecules with expanded valence shells, we will continue to use the traditional approach of including d-orbital hybridization for molecules with expanded valence shells in this text for simplicity, while acknowledging its limitations.

  • End-to-end overlap of atomic orbitals produces a connection, allowing unrestricted rotation of the molecule's linked components. A multiple bond is made up of a bond and either one (double bond) or two bonds (triple bond).

  • Multiple bonds contain more electron density between their nuclei than single bonds, resulting in higher bond energies. Rotation in a bond is limited by the side-to-side overlap of orbitals.

  • Molecular orbital (MO) theory considers a molecule to be a collection of nuclei with MOs distributed throughout the structure.

  • Atomic orbitals with equivalent energies can be combined or removed to produce bonding or antibonding MOs.

  • Bonding MOs, whether or, have most of the electron density between the nuclei and are lower in energy than the AOs that combine to create them; antibonding MOs have most of the electron density outside the nuclei and are thus greater in energy.

  • MOs are filled in descending order of energy by paired electrons with opposing spins. MO graphs depict energy levels as well as orbital occupancy. Period 2 homonuclear diatomic molecular diagrams

  • Several features of MO theory, including MOThe charge filling, energy-level diagrams, electron configuration, and bond order, are related to previous concepts: Electrons are injected into MOs. Electrons enter MOs in the same way as they do AOs:

  • MOs are filled in ascending energy order (Aufbau principle). An MO can only contain two electrons with opposing spins (Pauli exclusion principle).

  • Before any of the orbitals of equal energy are filled, they are half-filled, with electron spins parallel (Hund's rule).