Undergraduate
  1. General chemistry  1.1. Introduction to Chemistry  1.2. Atomic Structure  1.2.1. Subatomic particles  1.2.2. Atomic Model  1.2.3. Quantum numbers  1.2.4. Electron Configuration  1.2.5. Periodic trends  1.2.6. Isotopes and their applications  1.3. Chemical bond  1.3.1. Ionic bond  1.3.2. Covalent bonds  1.3.3. Metal Bond  1.3.4. Hybridization  1.3.5. Molecular geometry and VSEPR theory  1.3.6. Bond polarity and dipole moment  1.4. Stoichiometry  1.4.1. Mole concept  1.4.2. Empirical and molecular formula  1.4.3. Limiting reactant and excess reactant  1.4.4. Percent Yield and Purity  1.4.5. Dilution calculation  1.5. States of matter  1.5.1. Gas Laws  1.5.2. Matter and Intermolecular Forces  1.5.3. Solid and Crystal Structures  1.5.4. Phase diagram  1.5.5. Plasma and supercritical fluid  1.6. Chemical reactions  1.6.1. Types of Chemical Reactions  1.6.2. Balancing Chemical Equations  1.6.3. Thermochemical reactions  1.6.4. Reaction energy profiles  1.7. Solutions and Mixtures  1.7.1. Concentration units  1.7.2. Syndrome properties  1.7.3. Solubility and Solubility Rules  1.7.4. Henry's Law and Raoult's Law  1.7.5. Partition coefficient  1.8. Chemical equilibrium  1.8.1. Law of mass action  1.8.2. Le Chatelier's principle  1.8.3. Reaction quotient  1.8.4. To use this online calculator for Equilibrium Constant, enter Equilibrium Constant (Eq.) and hit the calculate button. Here is how the Equilibrium Constant calculation can be explained with given input values -> 0.05 = 0.05/0.  1.8.5. Acid-base balance  1.9. Acids and bases  1.9.1. Arrhenius, Brønsted–Lowry, and Lewis definitions  1.9.2. pH and pOH  1.9.3. Acid-base titration  1.9.4. Buffer Solution  1.9.5. Acid-base indicator  1.9.6. Polyprotic Acids and Bases  1.10. Electrochemistry  1.10.1. redox reactions  1.10.2. Galvanic and Electrolytic Cells  1.10.3. Nernst equation  1.10.4. Electrolysis  1.10.5. Faraday's laws of electrolysis  1.11. Thermodynamics  1.11.1. Laws of Thermodynamics  1.11.2. Enthalpy, entropy and Gibbs free energy  1.11.3. Spontaneity of reactions  1.11.4. Thermodynamic cycles (Hess's law, Carnot cycle)  1.12. Kinetics  1.12.1. Rate law and reaction order  1.12.2. Activation Energy and Catalyst  1.12.3. Collision theory  1.12.4. Reaction mechanism  1.12.5. Transition state theory
  2. Organic chemistry  2.1. Structure and relationships  2.1.1. Hybridisation in Carbon Compounds  2.1.2. Resonance and aromatization  2.1.3. Molecular orbitals  2.1.4. Inductive and mesomeric effects  2.2. Hydrocarbons  2.2.1. Hydrocarbons  2.2.2. Alkene  2.2.3. Alkynes  2.2.4. Aromatic compounds  2.2.5. Cycloalkanes and Conformation  2.3. Functional Group  2.3.1. Alcohols and phenols  2.3.2. Ethers and epoxides  2.3.3. Aldehyde and ketone  2.3.4. Carboxylic acids and derivatives  2.3.5. Amin  2.3.6. Esters and amides  2.3.7. Thiols and sulfides  2.4. Stereoscopic  2.4.1. Isomerism in Stereochemistry  2.4.2. Chirality and optical activity  2.4.3. Structural analysis  2.4.4. Diastereomers and Enantiomers  2.5.1. Substitution reactions  2.5.2. Elimination reactions  2.5.3. Addition reactions  2.5.4. Radical reactions  2.5.5. Rearrangement reactions  2.5.6. Pericyclic Reactions  2.6. Spectroscopy and structural analysis  2.6.1. Infrared Spectroscopy  2.6.2. Nuclear magnetic resonance spectroscopy  2.6.3. Mass Spectrometry  2.6.4. UV-visible spectroscopy  2.6.5. X-ray crystallography  2.7. Polymer chemistry  2.7.1. Addition and Condensation Polymers  2.7.2. Polymerization mechanism  2.7.3. Biodegradable Polymer  2.7.4. Conducting and smart polymers
  3. Inorganic chemistry  3.1. Coordination chemistry  3.1.1. Ligands and coordination compounds  3.1.2. Crystal field theory  3.1.3. Molecular Orbital Theory in Coordination Compounds  3.1.4. Jahn–Taylor distortion  3.2. Main Group Chemistry  3.2.1. Alkali and alkaline earth metals  3.2.2. Halogens and Noble Gases  3.2.3. Transition metals and their complexes  3.2.4. Lanthanides and Actinides  3.3. Solid state chemistry  3.3.1. Crystal lattices and unit cells  3.3.2. Defects in solids  3.3.3. Electrical and magnetic properties  3.3.4. Superconductivity  3.4. Bio-inorganic chemistry  3.4.1. Metal ions in biology  3.4.2. Enzymatic reactions with metals  3.4.3. Metal poisoning and detoxification  3.4.4. Metalloproteins and Metalloenzymes
  4. Physical chemistry  4.1. Quantum Chemistry  4.1.1. Wave–particle duality  4.1.2. Schrödinger Equation  4.1.3. Quantum Numbers and Orbitals  4.1.4. Atomic and molecular spectroscopy  4.2. Statistical mechanics  4.2.1. Maxwell–Boltzmann distribution  4.2.2. Partitioning functions  4.2.3. Bose–Einstein and Fermi–Dirac statistics  4.3. Chemical thermodynamics  4.3.1. Gibbs Free Energy  4.3.2. Phase transition  4.3.3. Thermodynamic efficiency  4.4. Surface chemistry  4.4.1. Absorption and Catalysis  4.4.2. Colloids and Emulsions
  5. Analytical Chemistry  5.1. Classical methods  5.1.1. Gravimetric Analysis  5.1.2. titrimeter  5.2. Instrumental methods  5.2.1. Chromatography  5.2.2. Spectroscopy  5.2.3. Electroanalytical methods  5.2.4. X-ray diffraction
  6. Industrial Chemistry  6.1. Petrochemistry  6.2. Chemical engineering principles  6.3. Green Chemistry  6.4. Pharmaceuticals and Drug Synthesis
  7. Biochemistry  7.1. Carbohydrates  7.2. Proteins and enzymes  7.3. Lipids and membranes  7.4. Nucleic acids  7.5. Metabolism and Bioenergetics  7.6. Enzyme kinetics and mechanism

Undergraduate → General chemistry → Chemical bond ↓


Hybridization


Hybridization is a concept in chemistry used to predict and explain the geometry and properties of molecules. Developed in the early 20th century, the concept helps explain how atomic orbitals mix to form new, equivalent hybrid orbitals during the formation of chemical bonds. The idea of hybridization is important in explaining the shapes and bonds of molecules in a range of substances.

Understanding atomic orbitals

Before diving into hybridization, it is important to understand what atomic orbitals are. Atomic orbitals are the regions where electrons are likely to be found around the nucleus of an atom. These orbitals are named as s, p, d and f depending on their shapes and energy levels.

- The s orbital is spherical. - The p orbitals are dumbbell-shaped and oriented along the x, y, and z axes. - The d and f orbitals have more complex shapes.

In an atom, electrons fill these orbitals according to the Aufbau principle, Hund's rule, and the Pauli exclusion principle. When atoms come together to form molecules, their atomic orbitals combine, or hybridize, to form new orbitals that affect the shape and energy of the molecule.

What is hybridization?

Hybridization is the process of combining atomic orbitals into new hybrid orbitals that are suitable for pairing electrons to form chemical bonds in molecules. This process allows the atom to form stronger and more stable chemical bonds.

In general, hybridization involves the mixing of orbitals on the same atom, rather than on different atoms. Different types of hybridization are named based on the characteristics of the new orbitals. Common types include:

  • sp
  • sp 2
  • sp 3
  • sp 3 d
  • sp 3 d 2

Hybridisation in methane: An example of sp 3 hybridisation

Let's take the example of methane (CH 4 ) to understand how hybridisation works. Methane has one carbon atom bonded to four hydrogen atoms. In its basic state, carbon has one 2s orbital and three 2p orbitals.

Carbon's ground state electron configuration: 1s 2 2s 2 2p 2

To form four equivalent bonds with hydrogen, carbon undergoes sp 3 hybridization. In this process, the 2s orbital combines with all three 2p orbitals to form four equivalent sp 3 hybrid orbitals.

Each sp 3 hybrid orbital = (1s + 3p)
C H H

Each of these sp 3 hybrid orbitals overlaps with the 1s orbital of the hydrogen atom, forming four σ (sigma) bonds, which are equal in energy and size. This gives methane a tetrahedral geometry with a bond angle of about 109.5 degrees.

Other types of hybridization

sp hybridisation in acetylene

Acetylene (C 2 H 2 ) is a classic example of sp hybridization. In this molecule, each carbon atom is bonded to one hydrogen and another carbon atom. Here, the carbon atom undergoes sp hybridization by combining one 2s and one 2p orbital to produce two sp hybrid orbitals.

Each sp hybrid orbital = (1s + 1p)

sp hybrid orbitals form a σ bond with hydrogen and another σ bond between the two carbon atoms. The remaining unhybridized 2p orbitals form two π (pi) bonds, giving acetylene a linear shape with a bond angle of 180 degrees.

sp 2 hybridisation in ethene

Ethene (C 2 H 4 ), or ethylene, shows sp 2 hybridization. Each carbon atom in ethene is bonded to two hydrogen atoms and another carbon atom. This is achieved by combining one 2s orbital with two 2p orbitals to form three sp 2 hybrid orbitals.

Each sp 2 hybrid orbital = (1s + 2p)

In ethene, sp 2 hybrid orbitals participate in σ bonding with hydrogen atoms and the other carbons. The remaining unhybridized 2p orbitals on each carbon overlap sideways to form a π bond, resulting in a planar structure with a 120-degree bond angle.

H H C C

sp 3 d hybridization in phosphorus pentachloride

Phosphorus pentachloride (PCl 5 ) is a good example of sp 3 d hybridization. In this molecule, phosphorus is surrounded by five chlorine atoms. The 3s orbital, three 3p orbitals, and one 3d orbital of phosphorus combine to form five equal sp 3 d hybrid orbitals.

Each sp 3 d hybrid orbital = (1s + 3p + 1d)

These hybrid orbitals arrange themselves in a trigonal bipyramidal structure, where three of the chlorines are in the equatorial position with an angle of 120°, and two of the chlorines are in the axial position with an angle of 180°.

sp 3 d 2 hybridisation in sulfur hexafluoride

An example of sp 3 d 2 hybridization is seen in sulfur hexafluoride (SF 6 ). In this compound, sulfur forms six bonds using the 3s, 3p, and two 3d orbitals, forming six equivalent sp 3 d 2 hybrid orbitals.

Each sp 3 d 2 hybrid orbital = (1s + 3p + 2d)

These are arranged in an octahedral geometry where all bond angles are 90 degrees. This structure is highly symmetrical, which gives SF 6 its unique properties.

Importance of hybridization

The concept of hybridisation is important for understanding molecular geometry and bonding. It allows:

  • Prediction of bond angles and molecular shapes.
  • Understanding the equivalence of bonds in molecules such as methane.
  • Explain the properties of complex compounds like transition metal complexes.

Although various methods of chemical bonding theories exist, hybridization provides a straightforward way to view and predict molecular structure and interactions.

Limitations of hybridization

Although hybridisation is a useful concept in general chemistry, it has its limitations. It is often not applicable to:

  • Molecules with non-covalent characteristics.
  • Transition metals (due to complex electron interactions).
  • Compounds involving heavier elements, where relativistic effects are important.

Advanced methods such as molecular orbital theory and valence bond theory can provide more accurate descriptions for such systems.

Conclusion

Hybridisation is a fundamental concept in chemistry that helps us understand how atoms bond and arrange themselves in three-dimensional space. Through examples such as methane, acetylene and phosphorus pentachloride, we can appreciate the variety of structures that arise from this process of mixing atomic orbitals. Although it may have its limitations, hybridisation remains an essential part of the chemist's toolkit for understanding the molecular world.


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