CHM136H1 Lecture Notes - Lecture 2: Lone Pair, Bond Length, Molecular Geometry
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30 Sep 2015
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CHM138H1 - Lecture 2 - Introduction to Organic Chemistry (Continued)
Chemical Bonding Theory through the Ages
● 1858: Kekulé and Couper independently observed that C always has four bonds, which
they called “affinity units”
● 1874: van’t Hoff and Le Bel proposed that the four C bonds have specific spatial
directions
● Atoms surround carbon as corners of a tetrahedron
● 1916: Lewis proposed that organic compounds have covalent bonds from sharing
electrons
Indicating Covalent Bonds
● Lewis structures (electron dot) show valence electrons of an atom as dots. For example:
○ H has one dot, representing its 1s electron
○ C has four dots (2s22p2)
● Kekulé structures (line-bond) have a line drawn between two atoms indicating a two-
electron covalent bond
Methane Ammonia Water
Lewis
Kekulé
● A stable molecule results at completed shell: octet (eight dots) for main-group electrons
(two for H) → the OCTET RULE
Valency Rules and Examples
1. Atoms with one/two/three valence electron(s) form one/two/three bond(s)
2. Atoms with four or more valence electrons form as many bonds as they need electrons
to fill the s and p levels of their valence shells to reach a stable octet
3. C has four valence electrons (2s22p2) and forms four bonds (e.g. methane)
4. N has five valence electrons (2s22p3) but forms only three bonds (e.g. ammonia)
5. O has six valence electrons (2s22p4) but forms only two bonds (e.g. water)
Remember the HONC rule:
H = 1 bond
O = 2 bonds
N = 3 bonds
C = 4 bonds
Non-Bonding Electrons
● Valence electrons not used in bonding are called non-bonding electrons (or lone pair
electrons)
● E.g. nitrogen atom in ammonia shares six valence electrons in three covalent bonds,
and the remaining two valence electrons are a non-bonding lone pair
Bond Formation Theories
● Covalent bond forms when a singly occupied orbital on one atom overlaps a singly
occupied orbital on another atom