Hybridization of atomic orbitals is a game-changer in understanding molecular structure. It explains how atoms form stronger, more stable bonds by mixing their orbitals. This process allows for diverse molecular shapes and bond types we see in nature.

By predicting hybridization states, we can figure out a molecule's geometry and bonding patterns. This knowledge is crucial for grasping how molecules behave and interact, making it a cornerstone of molecular physics and chemistry.

Hybridization in covalent bonding

Concept and role of hybridization

• Hybridization mixes atomic orbitals to form new hybrid orbitals with different shapes and energies
• Hybrid orbitals form by combining s and p orbitals within the same principal quantum shell
  • Example: sp hybridization combines one s orbital and one p orbital
• Hybridization allows for stronger, more stable covalent bonds by maximizing orbital overlap between bonding atoms
• The type of hybridization an atom undergoes depends on the number of electron domains (bonding and lone pairs) surrounding it

Importance of hybridization in covalent bonding

• Hybridization explains the observed geometries of molecules that cannot be predicted by the simple overlap of atomic orbitals
• Hybrid orbitals have shapes and orientations that maximize the overlap between bonding atoms
  • Greater overlap leads to stronger, more stable covalent bonds
• Hybridization allows for the formation of multiple bonds (double or triple bonds) by leaving unhybridized p orbitals available for pi (π) bonding
• Understanding hybridization helps predict the structure, stability, and reactivity of molecules

Hybridization state of atoms

Determining hybridization state

• The hybridization state of an atom is determined by the number of electron domains (bonding and lone pairs) around the central atom
• Molecules with two electron domains around the central atom adopt a linear geometry and undergo sp hybridization
  • Example: BeF2 (beryllium fluoride)
• Molecules with three electron domains around the central atom adopt a trigonal planar geometry and undergo sp² hybridization
  • Example: BF3 (boron trifluoride)
• Molecules with four electron domains around the central atom adopt a tetrahedral geometry and undergo sp³ hybridization
  • Example: CH4 (methane)

Exceptions and special cases

• Exceptions to the above rules exist in cases of multiple bonds (double or triple bonds) or molecules with expanded octets
• In molecules with multiple bonds, the hybridization state is determined by the total number of electron domains, including both sigma (σ) and pi (π) bonds
  • Example: C2H4 (ethene) has three electron domains (two C-H sigma bonds and one C=C pi bond) and undergoes sp² hybridization
• Molecules with expanded octets, such as SF6, involve the participation of d orbitals in hybridization
  • Example: SF6 undergoes sp³d² hybridization, resulting in an octahedral geometry

Formation of hybrid orbitals

Types of hybrid orbitals and their geometries

• sp hybridization mixes one s orbital and one p orbital, resulting in two linear sp hybrid orbitals oriented 180° apart
  • Example: CO2 (carbon dioxide)
• sp² hybridization mixes one s orbital and two p orbitals, resulting in three trigonal planar sp² hybrid orbitals oriented 120° apart
  • Example: SO3 (sulfur trioxide)
• sp³ hybridization mixes one s orbital and three p orbitals, resulting in four tetrahedral sp³ hybrid orbitals oriented 109.5° apart
  • Example: NH3 (ammonia)

Unhybridized orbitals and multiple bonds

• The remaining unhybridized p orbitals, if any, can form pi (π) bonds with other atoms
• Pi bonds result from the sideways overlap of unhybridized p orbitals
  • Example: In C2H4 (ethene), the carbon atoms are sp² hybridized, and the remaining unhybridized p orbitals form a pi bond between the carbons
• Multiple bonds consist of one sigma (σ) bond formed by the overlap of hybrid orbitals and one or more pi (π) bonds formed by the overlap of unhybridized p orbitals
  • Example: N2 (nitrogen) has a triple bond consisting of one sigma bond (sp-sp overlap) and two pi bonds (p-p overlaps)

Hybridization theory for molecular structure

Predicting molecular structure using hybridization

• Determine the number of electron domains (bonding and lone pairs) around the central atom in a molecule
• Based on the number of electron domains, predict the hybridization state of the central atom (sp, sp², or sp³)
• Use the hybridization state to determine the geometry of the molecule (linear, trigonal planar, or tetrahedral)
  • Example: NH3 has four electron domains (three bonding and one lone pair), so the N atom is sp³ hybridized, and the molecule has a tetrahedral electron domain geometry
• Assign the appropriate hybrid orbitals to the central atom and the unhybridized atomic orbitals to the surrounding atoms
• Determine the types of bonds formed between the atoms (sigma (σ) or pi (π) bonds) based on the overlap of the assigned orbitals

Bond angles and molecular shape

• Predict the bond angles between the atoms based on the geometry of the molecule
  • Linear geometry (sp hybridization) has a bond angle of 180°
  • Trigonal planar geometry (sp² hybridization) has bond angles of 120°
  • Tetrahedral geometry (sp³ hybridization) has bond angles of 109.5°
• The molecular shape may differ from the electron domain geometry due to the presence of lone pairs
  • Example: NH3 has a tetrahedral electron domain geometry but a trigonal pyramidal molecular shape due to the lone pair on the N atom
• Lone pairs occupy more space than bonding pairs, causing a slight decrease in bond angles compared to the ideal geometry
  • Example: In H2O (water), the bond angle is 104.5° instead of the ideal 109.5° due to the two lone pairs on the O atom