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.

Test your knowledge on sigma and pi bonds, sp2 and sp3 orbital hybridization, and the VSEPR theory for predicting molecular shapes. Dive deep into the concepts taught by Dr. Olubunmi A. Adewusi at Lagos State University.