1. 4.1 Lewis Theory of Bonding

2. Types of bonding conditions between elements Low Electronegativity and low Ionization energy (Metals) High electronegativity and High Ionization energy (Non-metals) Low Electronegativity and low Ionization energy (Metals) Metallic bondingIonic bonding (transferring of electrons between atoms) High electronegativity and High Ionization energy (Non-metals) Ionic bondingCovalent bonding (sharing of electrons between atoms)

3. Lewis Structures Gilbert Lewis created before the development of quantum mechanics Evidence - Known chemical formulas -Octet Rule -Valence -Electron shell model

4. Key Ideas Nobel gas configuration conveys stability e-’s most stable when paired Chemical bonds are formed to achieve an octet – Achieved by exchanging e-’s between a metal and a non-metal = ionic bond – Achieved by sharing e-’s between 2 non-metal= covalent bond

5. Lewis Diagrams for Ionic Compounds  Identify the number of valence shell electrons and determine the charge on the ion using the “stable octet rule”.  Write the elemental symbol, place dots to represent the electrons in the valence shell, enclose in square brackets and write the ionic charge as a superscript. [Na] + or [ Cl ] -

6. Lewis Diagrams or Structures A convention developed to “show” the relationship between atoms when they form bonds. Why is it necessary? predict where the electrons are in a molecule needed to predict the shape of a molecule

7. Lewis Diagrams for Ionic Compounds  Identify the number of valence shell electrons and determine the charge on the ion using the “stable octet rule”.  Write the elemental symbol, place dots to represent the electrons in the valence shell, enclose in square brackets and write the ionic charge as a superscript. [Na] + or [ Cl ] -

8. Lewis Diagrams for Covalent Compounds Draw the Lewis Diagram for nitrogen trifluoride (NF 3 ). Step 1. Count the valence electrons N = 5 F = 7 5 + 3( 7) = 26 valence electrons

9. Lewis Diagrams for Covalent Compounds Step 2. Write a skeletal structure. Use the least electronegative atom in the centre Electronegativity: N = 3.0 & F = 4.0 FF F N = a pair of e - (a single bond)

10. Lewis Diagrams for Covalent Compounds Step 3. Complete the octets for each terminal atom (except H) FF F N    

11. Lewis Diagrams for Covalent Compounds Step 4. Assign any additional electrons as lone pairs on the central atom  FF F N   

12. Ms. van Gaal says- Try it! Draw Lewis structures for (a) Cl 2(g), (b) NaCl (s), (c) NH 3(g).

13. Lewis Diagrams for Covalent Compounds Example 3. Chlorate ion, ClO 3 - ((1 x 7) + (3 x 6) + 1) = 26 ClOO O  :   : 

14. Co-ordinate covalent bonds In some covalent compounds, the bonds between atoms occur because one atom has donated both electrons to the covalent bond. This is called a coordinate covalent bond. N : H H H H+H+ + N H H H H + Nitrogen supplies the two lone pair electrons to this N-H bond. The H + ion has no electrons.

15. Co-ordinate covalent bonds To determine the number of coordinate covalent bonds – subtract the bonding capacity (lone valence electrons) from the number of bonds the atom has. N H H H H + Nitrogen Bonds  4 Bonding capacity  3 Coordinate bonds  4-3=1

16. Lewis Structures and Quantum Mechanics … It just won’t go away! As Lewis was putting together his theory, the quantum mechanical model was in its infancy - Thus there are correlations between them The four sides of the atom correspond to the s and p block Hund’s rule is being observed as one electron is placed on each side of an element before adding another

17. Exceptions to the Octet rule H, C, N, O, and halogens generally obey the octet rule in all circumstances Boron is a molecule that has an underfilled octet in bonding situations. Some elements will allow overfilled octets.

18. Work Pg. 200 #1-2 Pg. 204 #1-2 Pg. 205 # 1-6

19. Sulfur hexafluoride SF 6

20. Phosphorus pentachloride PCl 5