Introduction to Chemical Bonding

1
Introduction to Chemical Bonding
2
Notes Page 1

• Electronegativity the ability of an atom in a
  bond to attract electrons.
• Ion a charged particle. Obtained the charge by
  gaining or losing electrons.
• Positive ion achieved by losing electron(s).
• Negative ion achieved by gaining electron(s).
• Ionic Bond results from a transfer of
  electrons. One atom takes the electrons from the
  other forming ions. Held together by
  electrostatic force. (Opposites attract.)

3

• Covalent Bond results from the sharing of
  electrons between 2 atoms.
• Non-polar covalent bond in which electrons are
  equally shared.
• Polar covalent bonds that have an uneven
  distribution of charge due to unequal sharing of
  electrons, Based on electronegativity.
• Bonds between two unlike atoms are never
  completely ionic and are rarely completely
  covalent.
• The degree to which bonds are ionic or covalent
  depends on the electronegativity differences of
  the bonded atoms. (Table p. 304)

4
Bond characteristic
5
Electronegativity Chart
6
Examples

• H - Cl Bond 2.1 3.0 0.9
• H F 2.1 4.0 1.9
• H H 2.1 2.1 0.0

7
Covalent Bonding

• Molecule resulting particle when 2 or more
  atoms bond covalently.
• Diatomic Molecule molecule consisting of 2
  atoms
• Single bond covalent bond produced by the
  sharing of one pair of electrons.
• Double bond sharing of two pairs of electrons
• Triple bond sharing of 3 pairs of electrons

8
Covalent Bonding contd

• Bond Length the average distance between 2
  bonded atoms.
• Bond Energy energy required to break a chemical
  bond and form neutral atoms.
• Polyatomic ion a charged group of covalently
  bonded atoms.

9
Lewis Structures

• Structural formula indicated the kind, number,
  arrangement, and bonds of the atoms in a
  molecule.
• Atomic symbols represent inner-shell electrons
  nuclei
• Dashes between 2 atomic symbols represent shared
  electron pairs in covalent bonds
• Dots adjacent to only one atomic symbol represent
  unshared or lone electrons.
• Represents a bond (2 electrons)
• Represents an unshared electron

10
Lewis Structures contd

• Central atom is the atom furthest to the left on
  the periodic table (Except hydrogen)
• All atoms need to have 8 electrons in their outer
  level to be stable except H
• H needs 2 electrons in its outer level

11
Drawing Lewis Structures

• Step 1 Find the total number of valence
  electrons that the elements have available.
• Example CH4

12
Drawing Lewis Structures

• Step 2 Find the total of electrons needed to
  have a complete outer shell. (Remember All
  elements need 8 EXCEPT H, which only needs 2)

13
Drawing Lewis Structures

• Step 3 Subtract the HAVES from the NEEDED.
  This is the of electrons that must be SHARED.
• Needed 16
• Have -8
• electrons that must be shared
  8

14
Lewis Structures

• Step 4 Find the number of bonds
• Divide the of SHARED electrons by 2 to get the
  number of bonds.
• 8 4 bonds (represented by 1 dash per
  bond)
• 2
• Step 5 Draw the structure.
• Put the atom furthest left on the periodic table
  (Except H) in the center. Fill in the bonds
  (dashes) and lone e- (dots)

15
Drawing of CH4