1. Bonding

2. Ionic Bonding Chemical bonding that results from the electrical attraction between cations and anions is called ionic bonding. Cations give up electrons to anions. Ionic compounds form between metallic ions and non-metallic ions.

3. Octet Rule Noble gases have a very stable structure and are somewhat unreactive. According to the octet rule, atoms bond in order to achieve an electron configuration that is the same as the electron configuration of a noble gas. As an example, Chlorine will gain an electron to have 18, which is how many the atom Argon has. These two atoms are said to be isoelectronic with each other.

4. Recall Lewis Diagrams Valence electrons involved in bonding can be represented by Lewis dot diagrams A chemical symbol represents the nucleus and the core electrons (not involved in bonding). Dots around the symbol represent valence electrons.

5. Recall Drawing Lewis Diagrams 1.Write the element symbol. 2.Draw dots, one for each valence electron 3.Dots should be spread over 4 sides It does not matter what side the dots are placed, but do not start to pair dots until there is one on each side The number of valence electrons is equal to the group number. With one exception. Cl

6. Lewis diagrams for the first 20 elements

7. Representing Ionic Compounds Lewis Diagrams Formation of sodium chloride: Na  +  [Na] + [] Cl  Cl  

8. Bohr Diagram

9. Lewis Diagram

11. It takes two Na to satisfy one sulfur.

12. Lewis Structures for Ionic Compounds Ba O O Ba 2+ 2- Mg Cl Cl Cl Mg 2+ - 2 Ba and O Mg and Cl BaO MgCl 2

13. Covalent Bonds A covalent bond results from the sharing of electrons between two atoms. The most familiar examples of covalent bonding are seen in the interactions of nonmetallic elements with one another. (CO 2, NO) Diatomic molecules are held together by covalent bonding (N 2, O 2, F 2 Cl 2, Br 2, I 2, As 2, H 2 ).

14. Covalent Bonding Formation of hydrogen chloride: H  + Cl    Cl  H  H - Cl  Covalent bond, shared electrons Lone pairs, valence electrons not involved in covalent bond Structural Formula: H-Cl (lone pairs are not drawn)

15. Lewis Structures H   H +  Cl  Cl  H2H2 or H H Cl 2 Cl   Cl   + or Cl Cl   H H Structural Formula: Cl-Cl

16. Multiple Bonds Atoms sometimes need to share more than a pair of electrons. If an atom is sharing two pairs of electrons, this is a double covalent bond. If an atom is sharing three pairs of electrons, this is a triple bond.

17. Double and Triple Bonds Atoms can share 4 electrons to form a double bond or 6 electrons to form a triple bond. The number of shared electron pairs (covalent bonds) that an atom can form is the bonding capacity. O2:O2: N2:N2: = O  N 

18. Multiple Covalent Bonds N N N N N N N N

19. Multiple Covalent Bonds C O O C O O C O O C O O