1. Chapter 8
Chemical Bonding

3. Octet Rule
Octet Rule: Atoms will gain, lose, or share valence electrons in order to obtain a stable outer shell electron configuration of 8 valence electrons. (H and He only need 2)

4. Lewis Electron Dot Structures
3 6
5 8 X
Lewis Structures – use dots to represent valence electrons

5. Lewis Electron Dot Structures
Another way

6. Draw the Lewis Structure for:
Phosphorus Chlorine
3. Magnesium Ion Oxygen Ion

7. Chemical Bonds
Chemical Bonds: Attractive forces that hold atoms together due to the mutual attraction between the nuclei and their valence electrons.
Three types of Chemical Bonds
Ionic
Covalent (Polar & Nonpolar)
Metallic

8. Metallic Bonds
Metallic bonds: result from the delocalized sharing of valence electrons between metal atoms. Example: Aluminum, Copper, Iron

9. Electronegativity
Electronegativity – the ability of an atom in a molecule to attract electrons to itself
Metals = low values
(Ionization energy too)
Nonmetals = high values
Note Trend:
Across  increase
Down  decrease

10. (less than 1.7 but more than zero)
Types of Bonds
Type 
Ionic
Polar Covalent
Nonpolar Covalent
Electronegativity Difference…
>1.7
(equal to or more than 1.7)
1.7>x>0
(less than 1.7 but more than zero)
= zero
Valence Electrons are…
Transferred
Shared unequally
Shared Equally
How to Recognize…
Metal & Nonmetal
2 Different Nonmetals
2 of the same Nonmetals
Chemical Formula for one example…
NaCl
H2O H2

11. Ionic Bonding Metals lose electrons and form cations . Na +
Nonmetals gain electrons and form anions. Cl –
3. Ionic Bonds form due to the attractions between the metal ions and nonmetal ions.

12. Lattice Energy
Lattice Energy is the energy needed to separate an ionic lattice into gaseous ions.
Lattice Energy increases with:
 Increasing charge on the ions
LE stronger in MgO than NaCl
 Decreasing distance between the ions
(LE is stronger between smaller radii ions)
LE stronger in LiF than in CsBr
3. The greater the LE, the higher the melting point.

13. Explain These Trends in Lattice Energy
NaCl > RbBr
BaO > KF
CaF2 > BaF2

14. HW - Textbook
#8.1 #8.8 #8.12ab #8.14 #8.19 #8.21 #8.22

15. Covalent Bonds vs Ionic Bonds
Covalent Bonds: form when valence electrons are shared between atoms of 2 nonmetals Ionic Bonds: form when valence electrons are transferred from a metal atom to a nonmetal atom

16. Covalent Bonds Single Bond – sharing of 1 electron pair
Double Bond – sharing of 2 electron pairs
3. Triple Bond – sharing of 3 electron pairs

17. Bond Length & Strength
Single bonds: Longer Bond Lengths Lowest Bond Energies (weak bonds) Triple bonds: Shorter Bond Lengths Highest Bond Energies (strong bonds)

18. Bond Polarity
Nonpolar Covalent Bond – electrons are shared equally between two atoms
Br2
2. Polar Covalent Bond – one atom exerts a greater attraction for the bonding atoms
H2O
3. The greater the electronegativity difference between the atoms, the more polar the bond.

19. Which Bond is Most Polar?
C – O or Si – F B – Br or S - Cl

20. Rules for Drawing Lewis Structures for Molecules & Polyatomic Ions
Find the number of valence electrons in each atom and add them up.
(Note: - ions have gained e-, + ions have lost)
Draw symbols of the atoms near each other in the way they will bond.
Connect the atoms by single bonds (1 e- pair)
Complete octets on atoms bonded to the central atom. (except hydrogen gets 2)
Place leftover electrons on the central atom.
If there aren’t enough electrons to give the central atom an octet, try double or triple bonds.

21. Do These PCl3 CH2Cl2 HCN BrO3 -
Note: The central atom is often written first, but not always. It is usually the less electronegative element.
PCl CH2Cl2
HCN BrO3 -