hybridization and dipole moment

• Sp3d2 hybridization occurs when one s orbital, three p orbitals, and two d orbitals combine to form six hybrid orbitals.
• sp hybridization occurs when mixing of one s and one p- orbital resulting in the formation of two equivalent sp hybrid orbitals.
• Bond parameters:

  • Bond length is the equilibrium distance between nuclei of bonded atoms
  • Bond angle is the angle between orbitals containing bonding electron pairs
  • Bond enthalpy is the energy required to break one mole of bonds
  • Bond order is the number of bonds between two atoms in a molecule
• Resonance structure:

  • Single Lewis structure may not be adequate for representing a molecule
  • Isoelectronic species have the same bond order
  • Resonance structures are used to represent molecules with multiple bonding possibilities
• In an ozone molecule, the two bond lengths are expected to be unequal due to the structure, but experimentally both bond lengths are equal (128pm) and the bonds are intermediate between single and double bonds
• The Lewis structure for ozone does not account for the observed experimental facts
• An alternative structure is proposed where double and single bonds are interchanged, but neither of these structures explains the observed facts
• The actual structure of ozone is a resonance hybrid of the two proposed structures
• Resonance stabilizes the molecule as the energy of the resonance hybrid is less than the energy of any single canonical structure
• Resonance averages the bond characteristics as a whole
• The energy of the O3 resonance hybrid is lower than either of the two canonical forms
• The existence of a completely ionic or covalent bond is an ideal situation, in reality, no bond or compound is completely covalent or ionic
• In homonuclear covalent bonds like H2, O2, Cl2, N2, or F2, the shared pair of electrons is equally attracted by two atoms, resulting in a non-polar covalent bond
• In heteronuclear covalent bonds like HF, the shared electron pair gets displaced more towards the atom with higher electronegativity, resulting in a polar covalent bond
• The dipole moment of a molecule is the product of the magnitude of charge and the distance between the centers of positive and negative charge
• The dipole moment is expressed in Debye units (D) with a conversion factor of 1D=3.33564X10-30 Cm
• The dipole moment is a vector quantity represented by crossed arrow heads on negative and positive centers
• In polyatomic molecules, the dipole moment depends on the individual dipole moments of the bonds and the spatial arrangement of the bonds in the molecule
• The shape of a molecule depends on the number of valence shell electron pairs around the central atom
• Pairs of electrons in the valence shell repel each other to minimize repulsion and maximize distance between them
• Lone pair - lone pair repulsion is greater than lone pair - bond pair repulsion, which is greater than bond pair - bond pair repulsion
• Lone pairs occupy more space and result in greater repulsion compared to bonding pairs, leading to deviations from idealized shapes and alterations in bond angles
• In the valence bond treatment for an H2 molecule, when two H-atoms approach each other, new attractive and repulsive forces begin to operate between the nuclei and electrons of the atoms
• Attractive forces tend to bring two atoms close to each other
• Repulsive forces tend to push two atoms apart
• The bond length in an H2 molecule is 74pm where the net force of attraction balances the force of repulsion
• Types of covalent bonds:

  • Sigma (σ) bond
  • Pi (π) bond
• Sigma (σ) bond:

  • Formed by end to end (head on) overlap of bonding orbitals along the internuclear axis
  • Can be formed by s-s, s-p, or p-p overlapping
• Pi (π) bond:

  • Formed by atomic orbitals overlapping in a way that their axes remain parallel to each other and perpendicular to the internuclear axis
  • Orbital's formed due to sidewise overlapping consist of two saucer type charged clouds above and below the plane of the participating atoms
• Strength of Sigma and Pi bonds:

  • Sigma bond is stronger than Pi bond
  • Sigma bond involves more extensive overlapping than Pi bond
• Hybridization:

  • Concept introduced by Pauling to explain geometrical shapes of polyatomic molecules
  • Atomic orbitals combine to form new set of equivalent hybrid orbitals
  • Hybridization redistributes energies resulting in new orbital's of equivalent energies and shapes
• Salient features of hybridization:

  • Number of hybrid orbitals equals the number of atomic orbitals that get hybridized
  • Hybrid orbitals are always equivalent in energy and shape
  • Hybrid orbitals are more effective in forming stable bonds than pure atomic orbitals
  • Hybrid orbitals are directed in space to minimize repulsion between electron pairs
• Conditions for hybridization:

  • Orbitals in the valence shell of the atom are hybridized
  • Orbitals undergoing hybridization should have almost equal energy
  • Promotion of electrons is not essential prior to hybridization
  • Filled orbitals of valence shell can also participate in hybridization
• Types of hybridization:

  • sp hybridization: mixing of one s and one p orbital resulting in two equivalent sp hybrid orbitals
  • sp2 hybridization: mixing of one s orbital with two p orbitals resulting in three equivalent sp2 hybrid orbitals
  • sp3 hybridization: mixing of one s orbital with three p orbitals resulting in four equivalent sp3 hybrid orbitals
• Hybridization in Water (H2O) and Ammonia (NH3):

  • In NH3, nitrogen undergoes sp3 hybridization with three unpaired electrons and a lone pair
  • In H2O, oxygen undergoes sp3 hybridization with two lone pairs and two hydrogen atoms
  • In Ethane (C2H6), carbon atoms undergo sp3 hybridization forming four sp3 hybrid orbitals
• Hybridization in Ethene (C2H4):

  • In Ethene, carbon atoms undergo sp2 hybridization
  • One sp2 hybrid orbital of each carbon atom overlaps axially to form a sigma bond
  • Two other sp2 hybrid orbitals of each carbon atom form sigma bonds with hydrogen atoms
  • Unhybridized orbital of one carbon atom overlaps with the other to form a weak pi bond
• In ethane molecule, the C-C bond consists of one sp2-sp2 sigma bond and one pi-bond between p-orbitals which are not used in the hybridization
• Type of hybridization in each C-atom in ethane: sp2
• Geometry of hybrid orbitals of each C-atom in ethane: trigonal planar
• Bond angle in ethane: H-C-C: 120°, H-C-H: 117.6°