Biol 112L - Hydrocarbons and Hybridization

Hydrocarbons

• Hydrocarbons are compounds consisting solely of carbon (C) and hydrogen (H).

Hybridization

• Hybridization involves the overlap of electron orbitals from different atoms to form chemical bonds.
• sp3 hybrid orbitals have bond angles of approximately 109.5°, resulting in tetrahedral geometry, formed from one s orbital and three p orbitals.
• Formation of single-bonded hydrocarbons occurs via sp3 hybrid orbitals between carbon atoms.

sp3 Hybrid Orbitals

• The superscript in sp3 indicates the combination of orbitals, not the number of electrons.
• Asymmetry in sp3 orbitals arises from p orbitals having different algebraic signs in the wave function, leading to strong directional orientation.

Methane and Ethane

• Carbon-carbon bonds in ethane and methane form from the overlap of sp3 hybrid orbitals.

Double Bonds and sp2 Hybridization

• sp2 hybrid orbitals form with a bond angle of about 120° (trigonal geometry), created from one s orbital and two p orbitals.
• Hydrogen bonding in double-bonded hydrocarbons involves strong s bonds from sp2 orbitals and pi bonds from unhybridized p orbitals.

Triple Bonds and sp Hybridization

• Triple bonds form via sp hybrid orbitals with bond angles of approximately 180° (linear geometry).
• Two sp hybrid orbitals create a strong s bond; remaining p orbitals form two pi bonds by sideways overlap, resulting in sharing six electrons.

Examples of Hydrogen Bonds

• In methylamine, the H-N-H bond angle is 107.1°, and C-N-H bond angle is 110.3°, reflecting sp3 hybridization.
• In methanol, the C-O-H bond angle is 108.5° due to two of the oxygen's sp3 hybrid orbitals being occupied by lone pairs.

Phosphorus and Sulfur

• Phosphorus commonly forms five covalent bonds, while sulfur typically forms four.
• O-P-O bond angles generally range from 110 to 112°, indicating sp3 hybridization in phosphorus.

Bond Polarity and Electronegativity

• Polar covalent bonds arise when bonding electrons are attracted more strongly to one atom, causing asymmetrical electron distribution.
• Electronegativity measures an atom's ability to attract shared electrons in a bond, generally increasing left to right across the periodic table and decreasing top to bottom.

Bond Type and Electronegativity Differences

• Nonpolar covalent bonds occur with electronegativity differences of less than 0.5.
• Polar covalent bonds arise from differences between 0.5 and 2.
• Ionic bonds form with differences greater than 2.

Molecular Polarity

• Molecular polarity results from the sum of individual bond polarities and lone-pair contributions, influencing solubility in solvents like water.
• The dipole moment quantifies net molecular polarity, defined by the charge magnitude and distance between charges, measured in debyes.

Dipole Moments

• Structures exhibiting strong dipole moments may contain lone pairs.
• Symmetrical structures can yield zero dipole moments.

Questions for Review

• Identify the more electronegative element from pairs: (a) Li or H (b) B or Br (c) Cl or I (d) C or H.
• Describe the direction of polarization in the C-Cl bond of chloromethane.
• Create three-dimensional models of specified molecules and predict their dipole moments.