1.5 Describing Chemical Bonds: Valence Bond Theory
3 min read•may 7, 2024
Covalent bonding forms when atoms share electrons, creating a stable molecule. explains this by describing how overlap. The strength and length of these bonds depend on factors like atom size and .
, the strongest type of , form through direct orbital overlap. They're symmetrical around the and have concentrated electron density between nuclei. Understanding these concepts helps explain molecular structure and reactivity.
Valence Bond Theory and Covalent Bonding
Valence bond theory fundamentals
Explains covalent bonding as the overlap of atomic (regions in space where electrons are likely to be found)
Orbitals involved in bonding called
Orbital overlap leads to formation of covalent bond by sharing electrons between two atoms
Sharing of electrons results in lower energy state compared to isolated atoms (H2 molecule vs. two separate H atoms)
Extent of orbital overlap determines
Greater overlap leads to stronger bond
Overlap influenced by orbital size, shape, and energy (s orbitals have greater overlap than )
of atomic orbitals can occur to form new hybrid orbitals with different shapes and energies
Characteristics of sigma bonds
Formed by direct, end-to-end overlap of atomic orbitals
Involves overlap of s orbitals or end-to-end overlap of p orbitals (s-s, s-p, or p-p overlap)
Cylindrically symmetrical about the bond axis
Electron density concentrated along internuclear axis
Rotation around bond axis does not change electron density distribution
Strongest type of covalent bond due to direct overlap of orbitals
Electron density concentrated between nuclei, providing maximum stabilization
Examples:
H-H bond in H2 molecule (s-s overlap)
C-H bond in methane (s-sp3 overlap)
C-C bond in ethane (sp3-sp3 overlap)
Bond strength and length
Bond strength: measure of energy required to break a bond
Expressed in terms of (BDE)
BDE: energy required to break a bond in a diatomic molecule in the gas phase
Higher BDE indicates stronger bond (C-C bond in ethane has higher BDE than C-H bond in methane)
: distance between nuclei of two bonded atoms
Measured in (pm) or (Å)
Shorter bond lengths generally correspond to stronger bonds (C=C bond in ethene is shorter than C-C bond in ethane)
Factors affecting bond strength and length:
Atom size: Bonds between smaller atoms tend to be shorter and stronger (H-H bond is shorter and stronger than Cl-Cl bond)
Orbital overlap: Greater overlap leads to shorter and stronger bonds (s-s overlap results in shorter and stronger bonds compared to p-p overlap)
: Higher bond order (double or triple bonds) results in shorter and stronger bonds (C≡C bond in ethyne is shorter and stronger than C=C bond in ethene)
: Delocalization of electrons through resonance structures can affect bond strength and length
Calculating bond strength using : E=4πϵ0rQ1Q2
E: electrostatic potential energy
Q1 and Q2: charges on the atoms
ϵ0: permittivity of free space
r: distance between the charges ()
Stronger bonds have more negative electrostatic potential energy (C-C bond in ethane has more negative E than C-H bond in methane)
Advanced Bonding Theories
provides a more comprehensive explanation of bonding by considering the formation of molecular orbitals from atomic orbitals
helps predict molecular geometry based on the arrangement of electron pairs around a central atom
Key Terms to Review (26)
Angstroms: An angstrom (Å) is a unit of length used to measure the size of atoms, molecules, and other small-scale structures. It is a fundamental unit in the description of chemical bonds and molecular geometry.
Atomic Orbitals: Atomic orbitals are the wave-like functions that describe the behavior and spatial distribution of an electron in an atom. They are the fundamental building blocks of atomic structure and play a crucial role in understanding chemical bonding and reactivity.
Bond Axis: The bond axis is an imaginary line that passes through the centers of two bonded atoms, representing the path along which the chemical bond is formed. It is a fundamental concept in the valence bond theory, which describes the formation and nature of chemical bonds.
Bond Dissociation Energy: Bond dissociation energy is the amount of energy required to break a specific chemical bond between two atoms, separating them into individual, free atoms. This term is crucial in understanding the stability and reactivity of molecules, as well as the energetics of chemical reactions.
Bond dissociation energy, D: Bond dissociation energy is the amount of energy required to break a bond between two atoms in a molecule into two separate, radical species. It is measured in kilojoules per mole (kJ/mol) and varies depending on the type of bond and the molecules involved.
Bond length: Bond length is the average distance between the nuclei of two bonded atoms in a molecule. It determines the stability and strength of the bond, varying with bond order and atom size.
Bond Length: Bond length refers to the distance between the nuclei of two bonded atoms in a molecule. It is a crucial parameter in understanding the structure and stability of chemical bonds, as it directly influences the strength and properties of the bond.
Bond Order: Bond order is a fundamental concept in chemical bonding theory that describes the strength and stability of a chemical bond between atoms. It is a measure of the number of shared electron pairs between two atoms and is directly related to the bond's overall strength and length.
Bond strength: Bond strength is the measure of energy required to break a bond between two atoms in a molecule. It directly influences the stability and reactivity of molecules in organic chemistry.
Coulomb's Law: Coulomb's law is a fundamental principle in electrostatics that describes the force of interaction between two stationary, electrically charged particles. It establishes a relationship between the magnitude of the electric force, the charges of the particles, and the distance between them.
Covalent Bond: A covalent bond is a chemical bond formed by the sharing of one or more pairs of electrons between two atoms. This type of bond is the fundamental basis for the stability and structure of many molecules, including those found in the topics of 1.5 Describing Chemical Bonds: Valence Bond Theory, 1.6 sp3 Hybrid Orbitals and the Structure of Methane, and 1.7 sp3 Hybrid Orbitals and the Structure of Ethane.
Electron Pair Repulsion Theory: The electron pair repulsion theory, also known as the VSEPR (Valence Shell Electron Pair Repulsion) theory, is a model used to predict the geometry of molecules based on the arrangement of electron pairs around a central atom. It explains how the repulsive forces between electron pairs in the valence shell of an atom determine the shape of a molecule.
Hybridization: Hybridization is a fundamental concept in chemistry that describes the process of mixing atomic orbitals to form new hybrid orbitals, which are used to explain the geometry and bonding patterns of molecules. This term is closely related to the development of chemical bonding theory, valence bond theory, and molecular orbital theory, as well as the structure and properties of various organic compounds.
Molecular Orbital Theory: Molecular Orbital Theory is a model that describes the behavior of electrons in a molecule by considering the formation of molecular orbitals from the combination of atomic orbitals. This theory provides a more comprehensive understanding of chemical bonding compared to the earlier Valence Bond Theory.
Orbital Overlap: Orbital overlap refers to the interaction and sharing of electron density between two or more atomic orbitals, which is a fundamental concept in understanding the formation of chemical bonds. This term is particularly relevant in the context of valence bond theory, sp3 hybrid orbitals, and sp hybrid orbitals.
Orbitals: Orbitals are regions in an atom where an electron is likely to be found. They are fundamental to understanding chemical bonding and the structure of molecules, as they describe the distribution of electrons around the nucleus of an atom.
P Orbitals: p Orbitals are a type of atomic orbital in which the electron is distributed in a dumbbell-shaped region around the nucleus. They are critical in understanding the formation of chemical bonds, the geometry of molecules, and the behavior of conjugated systems.
Picometers: Picometers are a unit of length in the metric system, equal to one trillionth (10^-12) of a meter. This extremely small unit is often used to measure the size of atoms, molecules, and other nanoscale structures in the context of chemical bonding and molecular structure.
Resonance: Resonance is a fundamental concept in organic chemistry that describes the ability of certain molecules to exist in multiple equivalent structures or resonance forms. This phenomenon arises from the delocalization of electrons within the molecule, leading to the stabilization of the overall structure and the distribution of electron density across multiple atoms.
Sigma (σ): A sigma bond is a type of covalent bond where two atomic orbitals overlap directly between the nuclei of two atoms. It represents the strongest type of covalent chemical bond formed through the head-on (axial) overlap of atomic orbitals.
Sigma (σ) bonds: A sigma bond is the strongest type of covalent chemical bond where two atomic orbitals directly overlap between the nuclei of two atoms. This bond involves the sharing of a pair of electrons by two atoms in a molecule.
Sigma Bonds: Sigma bonds are the strongest type of covalent bonds formed between atoms, characterized by a high electron density along the internuclear axis. They are essential in the development of chemical bonding theory, the description of chemical bonds using valence bond theory, and the understanding of resonance structures and their rules.
Valence bond (VB) theory: Valence Bond (VB) Theory is a model in organic chemistry that describes how atoms in a molecule are connected by bonds formed through the overlapping of atomic orbitals, each containing one electron from the bonding atoms. It emphasizes the importance of electron pairs in covalent bonds and explains the molecular shape based on these interactions.
Valence bond theory: Valence bond theory explains how atoms in a molecule are bonded together by overlapping their atomic orbitals to share electron pairs. It combines the concepts of orbital hybridization and resonance to account for the molecular shape and stability.
Valence Bond Theory: Valence bond theory is a model used to describe the formation of chemical bonds between atoms. It focuses on the pairing of unpaired valence electrons between atoms to create stable covalent bonds, allowing atoms to achieve a more stable electronic configuration.
Valence Orbitals: Valence orbitals are the outermost electron shells of an atom that participate in chemical bonding. They are the orbitals that determine the reactivity and bonding behavior of an element, as they are the electrons most readily available for sharing or transferring during chemical reactions.