Chapter 6 - The Chemical Bond

• For bonding we will look at the ELECTRONS
• a bond will form if a MORE STABLE electron config. is made.

1. Definitions

1. Octet - the 8 outer shell electrons of noble gases (except He)
2. Octet rule - the atoms of the representative elements form bonds so as to have access to eight outer electrons.
3. Valence electrons - The outer s an p electrons in the atom (greatest n)
4. Duet rule - elements that border He form bonds so as to have access to two outer electrons (H, Li, Be).
   

2. Obtaining ideal electron config.

1. A metal may LOSE 1-3 electrons to form a cation.
2. A nonmetal may GAIN 1-3 electrons to form an anion.
3. Atoms may SHARE electrons with each other to meet octet.
   

3. Lewis dot symbols - element symbol plus dots to represent the valence electrons.

1. Formation of positive ions (metals):
   
   
   
   
   
   
   
   
2. Formation of negative ions (nonmetals):
   
   
   
   
   
   
   
   

4. Binary Ionic Compounds

• React sodium with chlorine gas produce NaCl(s)
  Valence Electron from Na is given to Cl to complete OCTET
  
  
  
• Combine Li and O form Li₂O
  2 Valence electrons from 2 Li's given to O to complete OCTET
  
  
  
• Remember Formula for Ionic compound must balance to neutral (zero).

5. Ionic compounds - properties

1. solids at room temperature
2. tend to have high melting points
3. often brittle
4. form a three dimensional array of atoms - LATTICE

6. Bonds

Ionic bonds - electrostatic attractions between the positive charge of the cation and the negative charge of the anion (transfer of electrons)

Covalent bonds - A shared pair of electrons between two atoms
meet the octet or duet rule by combining valence electrons

7. Lewis Structures - show the order and arrangement of atoms in a molecule as well as all the valence electrons.

• Single bond - sharing one pair of electrons between two atoms.
  
  F₂:
  
  
• Double bond - sharing of two pairs of electrons between two atoms.
  
  C₂H₄:
  
  
• Triple bond - sharing of three pairs of electrons between two atoms.
  C₂H₂:

8. Rules for Drawing Lewis Structure:
   1. Decide if ions will be formed
   2. Count total number of valence electrons
   3. Place symbols in order of bonding
      Hydrogen always outside
      Most metallic element in center (usually one found in the least)
      Oxygens do not usually bond together, and are usually outer atoms
   4. Place one pair of electrons between each atom
   5. Place the other remaining electrons until octet or duet rule is met for all atoms (trial and error).
   6. Draw lines to represent single, double, or triple bonds.

Resonance structures - when more than one correct Lewis structure exists.

9. Polar covalent bond - a covalent bond that has a partial separation of charge due to Unequal sharing of electrons. (Polar bond)

Electronegativity - the ability of an atom to attract electrons to itself in a covalent bond (relative electronegativity is what is important). - Fig. 6-6

Dipole - a polar bond with a negative and positive end (two poles)
NONPOLAR -electrons are shared Equally

Difference in electronegativity. determines polarity (EN ³ 1.9 ionic, EN » 0 nonpolar, all others are polar - note #'s are just approximate).

Examples:

• Cl-Cl (DEN = 0)
  nonpolar
• NaCl (DEN = 2.1)
  ionic
• HCl (DEN = 0.9)
  polar

10. Valence Shell Electron Pair Repulsion (VSEPR)

Theory that electron pairs, lone or shared, repel each other to the maximum extent.

Example: BeH2

1. Look at Lewis Structure (note: Be and H follow the duet rule).

 

 

2. For the electrons to be as far away as possible, the two bonds must be on opposite sides of the Be.

 

 

Example CO2:

 

4 Electrons in each C-O bond, but a total of 2 groups.

 

 

Example: two molecules with same # of groups, but different shapes

BF3

SO2

 

11. Polarity of Molecules

• based on VSEPR
• molecule may contain all polar bonds, but be nonpolar
• look at geometry - do the polar bonds cancel?

To Chem142 Home Page