Graduate
  1. Physical Chemistry  1.1. Thermodynamics  1.1.1. Laws of Thermodynamics  1.1.2. Enthalpy and heat capacity  1.1.3. Entropy and free energy  1.1.4. Phase Equilibrium  1.1.5. Statistical thermodynamics  1.1.6. Thermodynamic cycle  1.1.7. Fugacity and activity  1.2. Quantum Chemistry  1.2.1. Principles of quantum mechanics  1.2.2. Schrödinger Equation  1.2.3. Particle in a box  1.2.4. Operators and eigenvalues  1.2.5. Atomic orbitals  1.2.6. Molecular orbital theory  1.2.7. Perturbation theory  1.2.8. Valence bond theory  1.2.9. Hybridization and chemical bonding  1.3. Chemical kinetics  1.3.1. Rate laws and reaction mechanisms  1.3.2. Collision theory  1.3.3. Transition state theory  1.3.4. Enzyme Kinetics  1.3.5. Chain Reactions and Polymerization  1.3.6. Photochemical reactions  1.4. Statistical mechanics  1.4.1. Partitioning functions  1.4.2. Molecular distribution functions  1.4.3. Boltzmann Distribution  1.4.4. Bose–Einstein and Fermi–Dirac statistics  1.5. Spectroscopy  1.5.1. Rotational Spectroscopy  1.5.2. Vibrational Spectroscopy  1.5.3. Electronic Spectroscopy  1.5.4. Nuclear magnetic resonance spectroscopy  1.5.5. Mass Spectrometry  1.5.6. Raman Spectroscopy  1.5.7. Electron paramagnetic resonance spectroscopy  1.6. Surface and colloidal chemistry  1.6.1. Absorption isotherm  1.6.2. Surface tension and wetting  1.6.3. Colloidal Stability  1.6.4. Catalysis  1.6.5. Emulsions and micelles  1.7. Electrochemistry  1.7.1. Nernst equation  1.7.2. Electrochemical cells  1.7.3. Conductivity and mobility  1.7.4. War  1.7.5. Fuel cells  1.7.6. Electrolysis
  2. Organic chemistry  2.1. Reaction mechanism  2.1.1. Nucleophilic substitution reactions  2.1.2. Electrophilic addition reactions  2.1.3. Elimination reactions  2.1.4. Rearrangement reactions  2.1.5. Radical reactions  2.1.6. Pericyclic reactions  2.2. Spectroscopy and structural determination  2.2.1. UV-Vis Spectroscopy  2.2.2. IR Spectroscopy  2.2.3. NMR Spectroscopy  2.2.4. Mass Spectrometry  2.2.5. X-ray crystallography  2.3. Stereoscopic  2.3.1. Chirality and optical activity  2.3.2. Structural analysis  2.3.3. Geometrical isomerism  2.3.4. Dynamic Stereochemistry  2.4. Organometallic Chemistry  2.4.1. Organolithium and organomagnesium reagents  2.4.2. Palladium-catalyzed cross-coupling reactions  2.4.3. Transition Metal Complexes  2.4.4. Metal–carbon bond  2.5. Polymer chemistry  2.5.1. Polymerization mechanism  2.5.2. Polymer Characteristics  2.5.3. Biodegradable Polymer  2.5.4. Conducting polymer  2.5.5. Supramolecular polymers  2.6. Medicinal chemistry  2.6.1. Drug design and development  2.6.2. Pharmacokinetics and pharmacodynamics  2.6.3. Structure-activity relationships  2.6.4. Molecular docking and drug screening
  3. Inorganic chemistry  3.1. Coordination chemistry  3.1.1. Crystal field theory  3.1.2. Ligand field theory  3.1.3. Spectrochemical Series  3.1.4. Chelation and stability  3.2. Organometallic Chemistry  3.2.1. Metal Carbonyls  3.2.2. Catalysis by organometallic complexes  3.2.3. Metallocenes  3.3. Bio-inorganic chemistry  3.3.1. Metalloproteins and enzymes  3.3.2. The role of metals in biological systems  3.3.3. Metal ion transport and storage  3.4. Solid state chemistry  3.4.1. Crystal Structures  3.4.2. Band theory of solids  3.4.3. Superconductors  3.4.4. Defects in the crystal  3.5. Lanthanides and Actinides  3.5.1. Electronic configuration in lanthanides and actinides  3.5.2. Coordination chemistry of the lanthanides  3.5.3. Magnetic Properties of Lanthanides
  4. Analytical chemistry  4.1. Chromatography  4.1.1. Gas Chromatography  4.1.2. High Performance Liquid Chromatography  4.1.3. Thin Layer Chromatography  4.2. Spectroscopic Techniques  4.2.1. Atomic absorption spectroscopy  4.2.2. X-ray diffraction  4.2.3. Inductively coupled plasma spectroscopy  4.3. Electroanalytical methods  4.3.1. Potentiometry  4.3.2. Voltammetry  4.3.3. Colometry  4.4. Mass Spectrometry  4.4.1. Ionization Technique  4.4.2. Fragmentation Pattern  4.5. Chemometrics  4.5.1. Multidisciplinary analysis  4.5.2. Machine Learning in Chemistry
  5. Theoretical and Computational Chemistry  5.1. Molecular dynamics simulation  5.1.1. Force fields and energy minimization  5.1.2. Monte Carlo simulation in molecular dynamics simulations  5.2. Quantum chemical methods  5.2.1. Hartree–Fock theory  5.2.2. Density functional theory  5.2.3. Semi-empirical methods  5.3. Computational drug design  5.3.1. Molecular Docking  5.3.2. QSAR Modeling  5.3.3. Virtual Screening
  6. Biochemistry  6.1. Enzyme Kinetics  6.1.1. Michaelis-Menten kinetics  6.1.2. Inhibition mechanism  6.2. Metabolism and Bioenergetics  6.2.1. Glycolysis  6.2.2. Citric acid cycle  6.2.3. Electron transport chain  6.3. molecular Biology  6.3.1. DNA replication and repair  6.3.2. Protein synthesis  6.3.3. Gene Regulation  6.4. Structural Biochemistry  6.4.1. Protein Folding  6.4.2. Membrane Biophysics
  7. Environmental Chemistry  7.1. Atmospheric Chemistry  7.1.1. Greenhouse gases  7.1.2. Ozone depletion  7.1.3. Air pollutants and smoke  7.2. Water Chemistry  7.2.1. pH and water hardness  7.2.2. Wastewater treatment  7.2.3. Aquatic Chemistry  7.3. Soil Chemistry  7.3.1. Heavy metal contamination  7.3.2. Soil pH and Buffering  7.3.3. Nutrient cycling  7.4. Toxicology and Chemical Safety  7.4.1. Toxicokinetics  7.4.2. Risk Assessment in Toxicology and Chemical Safety in Environmental Chemistry  7.4.3. Effects of pollutants on the environment

Graduate → Physical Chemistry → Quantum Chemistry ↓


Valence bond theory


Valence bond (VB) theory is a fundamental aspect of quantum chemistry and presents a detailed description of how atoms come together to form molecules. It emphasizes the importance of electron pairing and explains how chemical bonds are formed when atomic orbitals overlap. In VB theory, bonds are formed when atomic orbitals overlap and these overlaps result in sharing of electron pairs between atoms. This sharing leads to the formation of non-covalent or covalent bonds, which are the fundamental interactions between atoms.

Historical background

The roots of valence bond theory can be traced back to the early 20th century. It evolved as scientists refined their understanding of atomic structures and the forces that hold atoms together. Gilbert N. Lewis first introduced the concept of electron sharing in 1916, an idea that became central to understanding covalent bonds. Later, Linus Pauling and Walter Heitler built upon these foundations, establishing VB theory as an important model in chemical bonding.

Fundamental concepts

At the core of VB theory, bonds are viewed as pairings of electrons where orbitals overlap. Let's dive into some key concepts.

Atomic orbitals and hybridisation

Atomic orbitals are regions in an atom where electrons are more likely to be found. These orbitals can combine to form hybrid orbitals, which overlap during molecule formation. The process of hybridization involves mixing different types of orbitals (s, p, d, etc.) to form new hybrid orbitals. For example:

        sp^3 hybridisation: combining one s and three p orbitals to form four equivalent sp^3 orbitals.
        sp^2 hybridisation: combining one s and two p orbitals to form three equivalent sp^2 orbitals.
        sp Hybridization: Combining one s and one p orbital to form two equivalent sp orbitals.

Overlap of orbitals

The strength and nature of a chemical bond depend on the degree of overlap between atomic orbitals. There are two main types of overlap:

  • Sigma (σ) bond: This is the strongest type of covalent bond, formed by end-to-end overlap of orbitals. For example, when two s orbitals or one s and one p orbital overlap, a sigma bond is formed. HH bond in H 2 is a classic example of a sigma bond.
  • Pi (π) bond: This bond is formed by the lateral overlap of p orbitals and is usually weaker than the sigma bond. In double and triple bonds, one is a sigma bond, and the rest are pi bonds. For example, the molecule O 2 has one sigma bond and one pi bond.

Visual example: ethylene (C 2 H 4 )

Ethylene is a simple molecule that shows orbital overlap in VB theory. Consider how the carbon atoms in ethylene hybridize and form bonds:

        Carbon atom hybridization: sp^2

Each carbon in ethylene uses sp^2 hybrid orbitals to form sigma bonds with hydrogen atoms and other carbon atoms. p orbitals of the carbon atoms overlap to form pi bonds.

C C

In this diagram, the blue line represents the sigma overlap, while the red curve represents the pi overlap.

Pauling and Heitler model

Linus Pauling and Walter Heitler significantly advanced the VB theory by incorporating quantum mechanics into the understanding of chemical bonds. Their model involves calculating wave functions to describe the position and energy of the electron in the molecule. This method allows for a quantitative understanding of bond energies and angles.

Comparison with molecular orbital theory

While VB theory focuses on localized electron pairing and orbital overlap, molecular orbital (MO) theory provides a delocalized approach. MO theory describes electrons as occupying molecular orbitals that encompass the entire molecule rather than being restricted to individual atoms. Each theory has its own strengths:

  • VB theory: Able to better explain resonance, hybridisation and individual bonding positions.
  • MO theory: More effective for understanding magnetic properties and electron distribution in a molecule.

Applications of valence bond theory

VB theory provides information about the chemical and physical properties of compounds. Some applications include:

  • Chemical reactivity: VB theory helps in predicting reactions based on the strength and nature of chemical bonds.
  • Materials science: Understanding the bonding in materials such as graphite or diamond, which is directly related to their properties.
  • Spectroscopy: Interpretation of data on molecular vibrations and rotations in terms of bonding details.

Challenges and limitations

VB theory is powerful, but it has its limitations. It can be less effective in systems with highly delocalized electrons, such as in metals or large organic molecules. Additionally, the complexity of manually solving the wave functions restricts its practical use without computational assistance.

Conclusion

Valence bond theory remains an important framework in quantum chemistry, providing a robust way to understand how atoms bond in molecules. Despite its limitations, it provides valuable insights into the behaviour of chemical systems and remains a relevant tool for chemists around the world.


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Valence bond theory

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