2. ________________: mutual electrical attraction between the nuclei & valence e-’s of different atoms that bond together. The type of bonding is determined by the way the valence e-’s are redistributed.

3. ___________:neutral group of atoms held together by covalent bonds. e.g. H 2 O ______________:molecules containing only 2 atoms. E.g. O 2, CO, HF, NO

4. ______________: where chemical compounds tend to form, such that each atom achieves an octet of electrons in its valence shell. This is done by gaining, losing or sharing e - ’s (becoming an ion or entering a covalent bond)

5. e.g. Fluorine gas exists as F 2. F –- F Each F atom has achieved a stable octet by ________________of electrons. (do the same with: O 2, PF 3 )

6. e.g. HCl Chlorine has a stable octet, but H does not. That’s because _______________________ ___________________to elements just not having enough electrons. i.e. H, He, Li, Be & B How many electrons would it need to fill an octet? Is that possible?

7.  When atoms bond they __________________ ______electrons  In dot notation this is represented as two dots between symbols, one from each atom

8. (Do F 2, PF 3, HCl ) The _______________of electrons are also known as lone pairs. The shared pairs can be represented by a dash. (do F 2, & PF 3 ) These representations are known as: ___________________: which show the shared pairs as dots (or dashes) and the unshared pairs as dots.

9. The dots representing the lone pairs can also be dropped. The new representation is known as a ____________________________

10. A single shared pair is known as a _________________ Let’s consider O 2 : (Diagram) The sharing of 2 pairs of electrons between 2 atoms is known as a _______________.

11. Let’s consider N 2 : The sharing of 3 pairs of electrons between 2 atoms is known as a ____________________. Double & Triple Bonds are also known as: Multiple Bonds.

12. We still haven’t explained why carbon can form 4 bonds instead of 2… _________________ Let’s look at Carbon (6): (ec, orbital diag., Lewis, & Structural) It makes sense to assume that Carbon forms 2 covalent bonds.

13. But when Carbon bonds with other atoms, a special thing happens. The 2s & 2p merge together to form an ___________. Now apply Hund’s rule. So now, Carbon has 4 single bonds. Hybridization also applies to Be, B, & Si.

14. Let’s review with some examples: (______________________________ EN atom or C) (e.g’s of NH 3, HCN, C 2 H 6, C 2 H 4, C 2 H 2 ) Show Lewis structure

15. ________:results from 1 atom giving up its valence e-’s (  cation) & transferring them to another atom (  anion)e.g. NaCl __________: results from the sharing of valence e-’s between 2 atoms  Most bonds are between these extremes!

16. ___________________: the bonding valence e- ’s are equally shared by the atoms resulting in equal distribution of electrical charge. e.g. N 2

17. __________________: the bonding valence e-’s are more strongly attracted to the more EN atom resulting in an unequal distribution of the valence e-’s. It is still sharing, not a transfer like in ionic. e.g. CO 2

18. So: Chemical Bonding IonicCovalent Polar Non-Polar The type of bonding can be determined simply by ______________________in the ElectroNegativities ( ∆EN) of the 2 atoms.

20. E.g.’s: A H-F molecule has an EN difference of: 4.0 (for F) – 2.1 (for H) = 1.9 For Na-Cl the EN difference is: 3.0 (for Cl) – 0.9 (for Na) = 2.1 For H-H (H 2 ) the EN difference is: 2.1 (for H) – 2.1 (for H) = 0

21. The difference tells you what type of bonding that is occurring: > 1.7=Ionic < 0.3=Non-Polar Covalent 0.3  1.7=Polar Covalent or: EN difference = 00.31.73.3 I--------I--------------------I---------------------------I non-PolarIonic Polar (see p.162 Fig.6-2 )

22. Going back to the previous examples: H-F ∆ EN = 1.9  Ionic Na-Cl ∆ EN = 2.1  Ionic H-H ∆ EN = 0  Non-Polar Covalent Other examples: Mg-S ∆ EN = 2.5–1.2 = 1.3  Polar Covalent CO 2 ∆ EN = 3.5–2.5 = 1.0  Polar Covalent

23. -forces ______________ molecules. -weaker than ionic & covalent bonds. In Polar Covalent ( ∆ EN=0.3-1.7) The EN difference creates a______________from positive to negative end. (partial charges) e.g.I-Cl=>I---Cl=>ICl (2.5) (3.0)  +  - 0.5 Difference Dipole

24. Dipole-Dipole interaction Intermolecular force _________________ _________________ _________between partial charges Strong intermolecular force.

25. The _________ intermolecular force that has H partially bonded to an electronegative atom. E.g’s: Causes ___________ than normal boiling points  water is a liquid instead of a gas @ room temp. A type of dipole-dipole interaction but stronger e.g. H 2 O, HCl, HF, H 2 S … (do diagrams) I love water!!! (why?)

26. Non-Polar Molecules ( ∆ EN= 0 - 0.3) There is no dipole because the EN diff is too low. But a slight shift of the e-’s to one side creates a temporary dipole, which effects the next molecule, and so on. e.g. Draw Cl 2 gas with London dispersion, aka Van der Waals forces. Other e.g’s? O 2, H 2, F 2 … London Dispersion forces

28. Bond Energy: the_________ ____________ ____________ needed to break a chemical bond.