CHM 101: Hybridization and Shapes of Molecules

Covalent Bonding

• Each atom provides a 1s electron, which is used in covalent bonding through the overlap of the two 1s orbitals, forming a pair within the two overlapping orbitals.
• The two electrons form a strong bond due to the strong attraction they experience to the positive nuclei of the atoms.

Methane (CH4)

• The simplest organic molecule.
• Carbon has an electron configuration of 1s2, 2s2, 2px1, 2py1, 2pz.
• Orbital hybridization takes place, where the four valence orbitals of the carbon combine to form four equivalent hybrid sp3 orbitals.
• Each of the four valence electrons on the carbon occupies a single sp3 orbital.
• The four C-H bonds in methane are arranged with tetrahedral geometry about the central carbon, with each bond having the same length and strength.
• The sp3 hybrid orbitals are oblong in shape, with two lobes of opposite sign but different sizes.
• The bigger lobes are directed towards the four corners of a tetrahedron, making the angle between any two orbitals in methane 109.5°.

Hybridization and Shapes of Molecules

• Learning objectives:
  • Understand sigma and pi bonds, and examples of simple molecules having them.
  • Understand sp2, and sp3 orbital hybridization with specific examples.
  • Understand the principles of the valence shell electron pair repulsion (VSEPR) theory.
  • Understand the concept of bond angle.

Sigma (σ) Bonds

• Sigma bonds are the strongest type of covalent bonds due to the direct overlap of valence orbitals.
• The electrons in these bonds are sometimes referred to as sigma electrons.
• A single bond is usually a sigma bond.

One Hydrogen (H2)

• The simplest molecule to use to explain sigma bonds.
• The bond is formed from the overlap between a half-filled 1s orbital in a hydrogen atom and a sp3 hybrid orbital in the central carbon.

The Solid Dash/Wedge System

• A conventional way of drawing 2-dimensional structures of molecules.
• Solid wedges represent bonds that are meant to be pictured emerging from the plane of the page.
• Dashed wedges represent bonds that are meant to be pictured pointing into, or behind, the plane of the page.
• Normal lines imply bonds that lie in the plane of the page.

Ammonia (NH3)

• Nitrogen is sp3 hybridized.
• The bonding arrangement is also tetrahedral, where the three N-H bonds of ammonia form the base of a trigonal pyramid.
• The lone pair forms the top of the pyramid, and its slightly greater repulsive effect pushes the three N-H bonds away from the top of the pyramid, making the H-N-H bond angles not tetrahedral at 107.3°.

Water (H2O)

• The overlap of sp3 hybrid orbitals on oxygen with 1s orbitals on the two hydrogen atoms makes the bonding in water occur.
• The two non-bonding lone pairs on oxygen would be located in sp3 orbitals.
• A molecule of water is ‘bent’ at an angle of approximately 104.5° (explained by VSEPR).

pi (π) Bonds

• sp2 hybridization occurs in double-bonded compounds like ethene.

Ethene (C2H2)

• Valence bond theory, and the hybrid orbital concept, describes bonding in double-bonded compounds like ethene.
• The 2s, 2px, and 2py orbitals combine to form three sp2 hybrid orbitals, leaving the 2pz orbital unhybridized.
• Characteristics of the ethene molecule relating to its bonding:
  • It is a planar (flat) molecule.
  • Bond angles are approximately 120°, and the C-C bond length is 1.34 Å, significantly shorter than the 1.54 Å single carbon-carbon bond in ethane.
  • There is no rotation about the carbon-carbon double bond.