1. Unit 6 Notes – Part II Mr Nelson 2010

2. Bonding & Lone Pairs Electron pairs that are shared are called bonding pairs Electron pairs that are not bonded or shared are called lone or unshared pairs

3. Steric Number The total electron pairs as the steric number. Double or Triple bonds count as 1 steric number. The central atom in this molecule, A, has a steric number of four.

4. Determining the shape of molecules Electrons, whether they be bonding or lone pairs, repel each other. By assuming the electron pairs are placed as far as possible from each other, we can predict the shape of the molecule.

5. VSEPR VSEPR: Valence Shell Electron Pair Repulsion –Electron pairs will arrange so they are as far apart as possible. ALL HAVE NEGATIVE CHARGES –ALL e - pairs determine the shape, bonded or not, but ONLY bonded pairs determine the name of the shape.

6. Molecular Shapes The shape of a molecule plays an important role in its reactivity. By noting the number of bonding and lone pairs we can easily predict the shape of the molecule.

7. Different Bond Types Ionic (extremely polar) –Electrons are transferred Covalent –Polar – uneven sharing of electrons –Nonpolar – even sharing of electrons

8. Nonpolar Covalent Bonds

9. Polar Covalent Bonds

10. Ionic Bonds

11. Nonpolar, polar, and ionic bonds (a) – a nonpolar covalent bond (b) – a polar covalent bond (c) – an ionic bond

12. Electronegativity Electronegativity is the ability of atoms in a molecule to attract electrons to themselves. On the PT, EN increases: –…from left to right across a row. EN decreases –…from the top to bottom of a group (column).

13. Bond Properties and Electronegativity

14. Bond properties and electronegativity Any bond can be classified by subtracting the EN of the 2 elements involved

15. Polar Covalent Bonds Electrons are not always shared equally in compounds. Oxygen pulls harder on the electrons it shares with hydrogen than hydrogen does. Oxygen’s end of the molecule has more electron density than the hydrogen end.