Organic Chemistry Chapter 1: Structure and Bonding

Structure and Bonding

Organic Chemistry

  • Historical use

    • compounds in living organisms

  • Modern use

    • the study of carbon-containing compounds

1.1 Atomic Structure

  • Protons

    • +1 charge

    • mass = 1.6726 × 10^-27 kg

  • Neutrons

    • 0 charge (neutral)

    • mass = 1.6750 × 10^-27 kg

  • Electrons

    • -1 charge

    • mass = 9.1096 × 10^-31 kg (smallest)

(nucleus has protons and neutrons) (electrons floating around outside)

  • Atomic number (Z)

    • equal to number of protons an atom contains

  • Mass number (A)

    • total number of protons and neutrons

  • Isotopes

    • same atomic number (Z) but a different number of neutrons, so a different atomic mass

    • example: chlorine-35 and chlorine-37

  • Atomic Mass

    • weighted average mass of an element’s natural occurring isotopes

1.2 Orbitals

  • Wave equation

    • mathematical expression describing specific behavior of an electron in an atom

  • Wave function ψ = orbital

    • defined region of space around nucleus where electron may be

    • Types:

      • s

      • p

      • d

      • f

s orbital

1 s orbital

p orbital

3 p orbitals

Electron shells

1.3 Electron Configuration

  • Rules:

    • 1. Orbitals are filled from lowest energy to highest energy

    • 2. Only two electrons can occupy an orbital, and they must be opposite spin

    • 3. If more than one orbital of equal energy is available, one electron is placed in each orbital before electrons are paired.

Carbon → C → 6 (atomic number, so need 6 electrons) → so 6 protons and 6 electrons) → 6 e- → (1s²2s²2p²) (where ² means electrons)

(→ ←) 1s

(→ ←) 2s

(→ ) (→ ) ( ) 2p… the signs are the same so they cannot be paired together, and we need to fill up all empty space first

  • 2s²2p² will be valence electrons

  • 1s² is core

Carbon has 4 valence electrons

1.4 Chemical Bonding Theory

  • Kekule and Couper (1858)

    • Carbon is tetravalent

    • Carbon bonds to other carbons to form extended chains

  • Van’t Hoff and Le Bel (1874)

    • four bonds to carbon are not randomly oriented

    • atoms bonded to carbon arranged in a tetrahedron

Nature of Chemical Bonds

  • Ionic bonds

    • Force of electrostatic attraction between oppositely charged ions

    • More common in inorganic chemistry than in organic chemistry

    • Carbon is less likely to form a cation than metals

    • Carbon is less likely to form an anion than nonmetals

  • NaCl Na+ Cl

Covalent Bonding

  • An electron pair that is shared between two atoms

  • Molecules are neutral groupings of atoms held together by covalent bonds

  • 1916 by Lewis

    • Lewis structures (electron-dot structures)

    • covalent bond :

  • Octet rule

    • Atoms gain, lose, or share electrons to give a stable electron configuration characterized by 8 valence electrons.

    • Exception: Hydrogen only has 2 electrons (via sharing its 1s electron with another atom)

H - Cl

Example: Draw the line-bond structures of H2O and NH3 .

Practice: Draw the electron-dot and line-bond structures for methanol, CH3OH.

Example: Draw the line-bond structures of the following organic molecules.

C_3H_8

CH3CH2CH2CH2CH3