Unit 3 Part I: Chemical Bonds (continued!)
7. Covalent Bonds
 One or more electrons are simultaneously attracted to 2 nuclei
 Atoms achieve the octet by sharing electrons
 The bonding atoms’ orbitals overlap, which maximizes attraction between nuclei and
bonding electrons.
o As two atoms approach each other:
 Repulsive force between electrons and between protons of two atoms
 Attractive force between nuclei (protons) and electrons of other atom
 At point of maximum attraction, the attractive forces balance the
repulsive forces
 Atoms bond to form a molecule
 Atoms can share 2, 4, or 6 electrons
o Single bond: share 2 electrons
o Double bond: share 4 electrons
o Triple bond: share 6 electrons
 2 types of covalent bonds:
o Nonpolar covalent bond: electrons are shared equally between atoms
 Same or very similar electronegativity
o Polar covalent bond: electrons are shared unequally
 One atom will have a stronger attraction for the electrons due to
difference in electronegativity
Examples:
8. Properties of Covalent molecules
 Can be any state of matter at room temperature
 Examples:
 Low melting point and boiling point
 Do not conduct electricity
 Shape not considered a crystal lattice
9. Energy of Covalent Bonds
 Breaking bonds = endothermic
 Forming bonds = exothermic
 Dissociation energy: amount of energy required to break a specific covalent bond
o Correlation between bond length and bond strength
o As # of shared electron pairs increases, bond length decreases
o *multiple bonds are stronger
10. Lewis Structures
 Representation of a molecule that shows how the valence electrons are arranged
among atoms. *electron dot diagram for molecules!
1. Add up all of the valence electrons from the atoms
2. Draw a skeleton structure of the molecule
a. Hydrogen is always terminal
b. Put least electronegative element “central”
3. Use two electrons to form a bond between each pair of bound atoms (use a
line!)
4. Arrange remaining electrons to satisfy octet rule
 *Bonding electrons and lone pairs (non bonding or unshared electrons)
 Exceptions to the octet rule
o Electron deficient atoms
 B and Be can have less than an octet (B forms 3 single bonds, Be forms 2
single bonds)
 Example:
o Atoms with expanded octets
 Elements 3rd period and beyond can have an expanded octet due to unfilled
d-orbitals
 Place extra electrons around central atom
 Example:
o Odd electron molecules
 A lone, unpaired e Called free radicals
 Example:
 Resonance
o When two or more valid lewis structures can be drawn for the same molecule
o Electrons are actually delocalized over several atoms or bonds
o Delocalized electrons lower potential energy of molecule