2. Chapter 8 sections 1-7 Chapter 9 sections 1-6

3. Ionic – Formed by the electrical attraction of a positive cation to a negative anion. Don’t worry about “Energetics of Ionic Bond Formation” for now!

4. Nonpolar Covalent – A bond between atoms with equal (or nearly equal) electronegativities (the electrons are shared equally – neither atom has a slight charge)

5.  Polar Covalent – A bond between atoms with different electronegativities (the one with the higher electronegativity pulls the electrons more and thus that atom is slightly negative and the other atom slightly positive)

6.  The ability of an atom in a molecule to attract electrons to itself.  Generally increases as we move across a series and generally decreases down a group. (Fluorine has the highest!)  The larger the difference in electronegativity values between two atoms, the more polar the bond between them will be.

7.  Determination of Bond Type Step 1: Check for a metal - if there is a metal the bond is usually IONIC Step 2: If there is no metal subtract the electronegativity values (bigger – smaller) If the difference is: 0.4-ish or less then the bond is NONPOLAR 0.5-ish or greater then the bond is POLAR

8. Electronegativity Difference The bond is:Example(difference) 0-0.4NonpolarCl-Cl(0.0) 0.4-1.0PolarH-Cl(0.9) 1.0-2.0Very polarH-F(1.9) >2.0IonicNaCl(2.1)

9. Ex: Si & N 3.0 – 1.8 = 1.2 very polar  The atom with the higher electronegativity is partially negative  The atom with the lower electronegativity is partially positive

10.  Polar Molecules are also called “dipoles”  Dipole Moment is a quantitative measure of the magnitude of a dipole…the higher the dipole moment the more polar the molecule. Si-N

11.  A way to determine which Lewis Structure is the best representation of a molecule  The best one has atoms with formal charges closest to zero.

12.  Valence Bond Theory – a covalent bond is an overlap of atomic orbitals that allows two electrons to share a space in both atoms.

13.  Sigma ( σ ) bond – the end-to-end overlap of orbitals.  All single bonds are sigma bonds

14.  Pi ( π ) bond – A side-to-side overlap of orbitals  Double bond – 1 sigma & 1 pi  Triple bond – 1 sigma & 2 pi

15.  Consider methane (CH 4 )  The carbon makes 4 bonds, but there are only 2 available orbitals! ↑↓ ↑↓ ↑ ↑ __ 1s 2s 2p

16.  1 of the electrons from the 2s sublevel jumps to the 2p sublevel.  Now there are 4 orbitals available for bonding! ↑↓ ↑_ ↑ ↑ ↑ _ 1s 2s 2p

17.  One of these orbitals is an “s” orbital and the other 3 are “p” orbitals….when scientists look at methane, all the bonds look the same. They become sp 3 hybrid orbitals.

18. Bonds + Lone Pairs around the central atom Common hybridization 2sp 3sp 2 4sp 3 5sp 3 d 6sp 3 d 2 *Multiple bonds count as one!!

19.  Electrons are spread out over a number of atoms in a molecule rather than localized between a pair of atoms.  Often happens in molecules with resonance.