8.1 Valence Bond Theory
3 min read•june 25, 2024
explains how atoms form covalent bonds through orbital . It's all about shared electrons and the strength of their connections. This theory helps us understand why molecules have specific shapes and bond strengths.
Sigma and pi bonds are the building blocks of molecular structures. Sigma bonds form through head-on orbital overlap, while pi bonds result from sideways overlap. These concepts are crucial for grasping how atoms join to create complex molecules.
Valence Bond Theory
Atomic orbital overlap in covalent bonds
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- Covalent bonds form when atomic orbitals overlap enabling electrons to be shared between atoms
- Orbitals must have matching symmetry and orientation to overlap effectively (s orbitals, p orbitals)
- Overlapping orbitals must also have comparable energies to form stable bonds
- Stronger covalent bonds result from greater overlap between orbitals
- Orbital overlap depends on the size and shape of the involved orbitals (s, p, d, f)
- Shared electrons in the overlapping region are attracted to both atomic nuclei lowering the overall energy and stabilizing the bond
- cloud forms in the region of orbital overlap where the shared electrons are typically located
- The extent of orbital overlap influences the and overall structure of the molecule
Sigma vs pi bonds
- Sigma () bonds
- Formed by head-on overlap of atomic orbitals (end-to-end)
- Can involve overlap of s orbitals, p orbitals, or a combination of s and p orbitals (sp, sp², sp³ hybrid orbitals)
- Electron density concentrated along the between the bonded atoms
- Stronger than pi bonds due to more extensive orbital overlap
- Pi () bonds
- Formed by sideways overlap of p orbitals (parallel)
- Require unhybridized p orbitals perpendicular to the internuclear axis
- Electron density concentrated above and below the internuclear axis
- Weaker than sigma bonds due to less orbital overlap
- Multiple bonds between atoms consist of one and one or more pi bonds
- : one sigma bond and one (C=C in )
- : one sigma bond and two pi bonds (C≡C in )
- The number and type of bonds between atoms determine the , which affects bond strength and length
Bond energy and orbital interactions
- is the energy required to break a chemical bond
- Stronger bonds have higher bond energies (H-H 436 kJ/mol, C-C 348 kJ/mol, C=C 614 kJ/mol, C≡C 839 kJ/mol)
- Higher bond energy and stronger bonds result from greater orbital overlap
- Orbital overlap depends on the size, shape, and orientation of the involved orbitals
- As the distance between bonded atoms increases, orbital overlap decreases leading to lower bond energy and a weaker bond
- Optimal bond length occurs when attractive and repulsive forces between atoms are balanced maximizing orbital overlap and bond energy
- Bond energy is also influenced by the difference between bonded atoms
- Polar covalent bonds have slightly lower bond energies than pure covalent bonds due to larger differences (H-Cl vs H-H)
- The bond energy of a molecule can be estimated by summing the average bond energies of all the bonds present
- , where is the molecule's bond energy, is each A-B bond's average bond energy, and is the number of A-B bonds in the molecule
Advanced concepts in bonding
- : the mixing of atomic orbitals to form new hybrid orbitals with different shapes and energies
- : predicts molecular geometry based on electron pair repulsion
- Resonance: occurs when multiple valid Lewis structures can be drawn for a molecule, resulting in a hybrid structure that better represents the actual electron distribution
Key Terms to Review (36)
Bond Energy: Bond energy, also known as bond dissociation energy, is a measure of the strength of a chemical bond. It represents the amount of energy required to break a specific chemical bond between two atoms, separating them into individual, free atoms. This concept is crucial in understanding the stability and reactivity of molecules, as well as the energy changes associated with chemical reactions.
Bond order: Bond order is the number of chemical bonds between a pair of atoms. It indicates the stability and strength of a bond.
Bond Order: Bond order is a concept 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 it plays a crucial role in understanding the properties and reactivity of molecules.
Coordinate covalent bond: A coordinate covalent bond is a type of covalent bond in which one atom provides both electrons for the shared pair. This often occurs between a Lewis base and a Lewis acid.
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 responsible for the stability and structure of many molecules and is a fundamental concept in understanding the topics of 7.2 Covalent Bonding, 7.3 Lewis Symbols and Structures, 7.4 Formal Charges and Resonance, 7.5 Strengths of Ionic and Covalent Bonds, 7.6 Molecular Structure and Polarity, and 8.1 Valence Bond Theory.
D Orbital: The d orbital is a type of atomic orbital in an atom that can hold up to 10 electrons. It is part of the valence shell and plays a crucial role in the formation of chemical bonds and the determination of molecular geometry.
D orbitals: d orbitals are a type of atomic orbital with angular momentum quantum number $l = 2$. They have complex shapes, including cloverleaf patterns and a toroidal shape.
Double bond: A double bond is a type of covalent bond in which two pairs of electrons are shared between two atoms. Double bonds are commonly found in organic molecules and influence molecular geometry and reactivity.
Double Bond: A double bond is a covalent bond in which two pairs of electrons are shared between two atoms, resulting in a stronger and more stable bond compared to a single bond. This type of bond is commonly observed in organic chemistry and plays a crucial role in the structure and reactivity of many chemical compounds.
Electron Density: Electron density is a fundamental concept in quantum mechanics that describes the probability distribution of electrons within an atom or molecule. It is a crucial factor in understanding the behavior and properties of chemical systems.
Electronegativity: Electronegativity is a measure of an atom's ability to attract and hold onto electrons within a chemical bond. It is a dimensionless quantity usually assigned values on the Pauling scale.
Electronegativity: Electronegativity is a measure of an atom's ability to attract shared electrons in a chemical bond. It is a fundamental property that influences the nature and strength of chemical bonds, as well as the physical and chemical properties of substances.
Ethene: Ethene, also known as ethylene, is a simple unsaturated hydrocarbon with the molecular formula C2H4. It is a colorless, flammable gas that is widely used in the chemical industry and is a key intermediate in the production of many organic compounds.
Ethyne: Ethyne, also known as acetylene, is a simple hydrocarbon with the chemical formula C$_2$H$_2$. It is a colorless, flammable gas with a distinctive odor and is widely used in various industrial and commercial applications.
F Orbital: The f orbital is one of the principal electron orbitals in an atom, characterized by a high angular momentum quantum number (l = 3) and a complex spatial distribution. It is the fifth principal energy level of an atom and is found in elements with atomic numbers greater than 57 (lanthanum).
Hybridization: Hybridization is the concept in chemistry where atomic orbitals combine to form new hybrid orbitals that are suitable for the pairing of electrons to form chemical bonds. This idea helps explain molecular geometry and bonding properties, linking the arrangement of atoms in a molecule to their electron configurations and the types of bonds formed.
Internuclear Axis: The internuclear axis refers to the imaginary line that connects the nuclei of two atoms involved in a covalent bond. This axis is a fundamental concept in valence bond theory, as it helps describe the orientation and geometry of bonding between atoms.
Molecular Geometry: Molecular geometry refers to the three-dimensional arrangement of atoms in a molecule. It is a fundamental concept in chemistry that describes the spatial configuration of atoms bonded together and plays a crucial role in understanding the properties and reactivity of molecules.
Node: A node is a point in an atomic or molecular orbital where the probability of finding an electron is zero. Nodes can be classified as either radial or angular depending on their nature.
Overlap: Overlap refers to the interaction between atomic orbitals on different atoms that leads to the formation of a covalent bond. The degree of overlap determines the strength and stability of the bond.
P orbital: A p orbital is a type of atomic orbital that has a dumbbell shape and can hold a maximum of two electrons. It is associated with the angular momentum quantum number l = 1, which means it has three different orientations in space, commonly referred to as p_x, p_y, and p_z. These orbitals play a crucial role in chemical bonding, particularly in the formation of covalent bonds.
Pi bond: A pi bond is a type of covalent bond formed when two atomic orbitals overlap laterally, resulting in an electron density that is concentrated above and below the internuclear axis. This bond occurs alongside a sigma bond in double or triple bonds, and plays a critical role in the structure and reactivity of molecules by influencing their geometry and electronic properties.
Pi bond (π bond): A pi bond (π bond) is a type of covalent bond formed by the lateral overlap of atomic orbitals, usually p-orbitals. It is characterized by electron density concentrated above and below the plane of the nuclei of the bonding atoms.
Polar covalent bond: A polar covalent bond is a type of covalent bond where electrons are shared unequally between atoms, resulting in a dipole moment. This occurs because one atom has a higher electronegativity than the other.
Polar Covalent Bond: A polar covalent bond is a type of chemical bond where the shared pair of electrons between two atoms is unequally distributed, resulting in one atom having a partial positive charge and the other a partial negative charge. This asymmetrical charge distribution gives the molecule an overall polarity.
S orbital: An s orbital is a type of atomic orbital that has a spherical shape and can hold a maximum of two electrons. This orbital is fundamental in understanding how electrons are arranged in atoms and plays a crucial role in bonding, particularly in the formation of covalent bonds and the concept of hybridization.
Sigma Bond: A sigma bond is a type of covalent chemical bond formed by the head-on overlap of atomic orbitals, resulting in a high electron density concentrated along the internuclear axis between two bonded atoms. This type of bond is the strongest and most common type of covalent bond, and it plays a crucial role in the stability and structure of molecules.
Sigma bonds (σ bonds): Sigma bonds (σ bonds) are covalent bonds formed by the head-on overlap of atomic orbitals. They allow for free rotation around the bond axis.
Sp Hybrid Orbital: An sp hybrid orbital is a type of atomic orbital that results from the hybridization of one s orbital and one p orbital. This hybridization occurs in atoms with a valence electron configuration of ns^1 np^1, where n represents the principal quantum number. The sp hybrid orbital has a linear shape and is used to describe the bonding in certain molecules.
Sp² Hybrid Orbital: The sp² hybrid orbital is a type of atomic orbital that arises from the hybridization of one s orbital and two p orbitals in an atom. This hybridization occurs in molecules with trigonal planar geometry, where the central atom forms three equivalent bonds with other atoms, resulting in a planar arrangement of the bonded atoms.
Sp³ Hybrid Orbital: The sp³ hybrid orbital is a type of atomic orbital that arises from the hybridization of one s orbital and three p orbitals in an atom. This hybridization results in the formation of four equivalent, tetrahedrally arranged hybrid orbitals, which are used to describe the bonding in many organic and inorganic compounds.
Triple bond: A triple bond is a type of chemical bond where three pairs of electrons are shared between two atoms. It is typically found in molecules like nitrogen (N₂) and acetylene (C₂H₂).
Triple Bond: A triple bond is a covalent bond in which three pairs of electrons are shared between two atoms, resulting in a very strong and stable chemical connection. This type of bond is found in various chemical structures and plays a crucial role in understanding valence bond theory, multiple bonds, and the properties of nitrogen.
Valence bond theory: Valence bond theory is a model that describes the formation of covalent bonds through the overlap of atomic orbitals. It emphasizes the role of electron pairs and hybridization in bond formation.
Valence Bond Theory: Valence Bond Theory is a model used in chemistry to describe the formation of covalent bonds between atoms. It explains how the sharing of valence electrons between atoms leads to the creation of stable chemical bonds.
VSEPR theory: VSEPR theory, or Valence Shell Electron Pair Repulsion theory, is a model used to predict the geometry of individual molecules based on the repulsion between electron pairs in the valence shell of central atoms. This theory emphasizes that the shape of a molecule is determined by the spatial arrangement of electron pairs, including both bonding and lone pairs, which minimizes the repulsive forces between them.