2.  Why do some solids dissolve in water but others do not?  Why are some substances gases at room temperature, but others are liquid or solid?  Why does ice water float?  What gives metals the ability to conduct electricity, what makes non-metals brittle?  The answers have to do with … Intermolecular forces

3.  There are 2 types of attraction in molecules: intramolecular bonds & intermolecular forces  We have already looked at intramolecular bonds. The attraction between atoms in a molecule. (ionic, polar, non-polar)  Intermolecular forces (IMF) have to do with the attraction between molecules. Intermolecular forces

5. 1) London Forces 2) Ion-Dipole 3) Dipole - Dipole, 3) H-bonding 5) covalent (network solids) 6) metallic

6.  In a non-polar molecule, on average the electrons are distributed equally over the molecule, but occasionally one side or the other will gain a small excess of electron density. When this occurs, the molecule has a temporary dipole- one side of the molecule has a slight + charge, the other a slight - charge.  When another molecule approaches the first, it can feel this dipole. The electrons around the second molecule can then rearrange so that there is a favorable interaction between the two molecules.

7. Temporary Dipoles Created.

8.  This type of IMF exist between all molecules. Simultaneous attraction between nucleus of one molecule and electrons in an adjacent molecule due to temporary dipoles. London Dispersion Forces Weakest of all IMF’s

10. ◦ Are present in all compounds ◦ Can occur between atoms or molecules ◦ Are due to electron movement ◦ Are transient in nature (dipole-dipole are not permanent - temporary ). ◦ London forces are weaker

11.  Solution: Each of the above molecules is a diatomic halogen, and thus similar in how the molecule behaves. Each sample of the bulk is held together by London dispersion forces. Since iodine is larger than bromine which is larger than chlorine, we expect that the London forces are largest in iodine and smallest in chlorine, so iodine should have the highest melting and boiling points. Laboratory observation: at room temperature, chlorine is a gas, bromine a liquid and iodine a solid. The more electrons, the stronger the London Forces !

12. Dipoles in polar molecule are attracted to ions (cation / anion) in the ionic structure.

14.  Attractive forces that exist between POLAR molecules. Partial charges in the molecule are created due to unequal sharing of electrons in the bond HCl ++ ––

15.  Molecules are attracted to each other in a compound by these partial +ve and –ve dipoles

16. 1. H-bonding is a special type of dipole - dipole attraction that is very strong (about 5x stronger) 2. H is attracted to the lone pair electrons N, O, or F. 3. The high  EN of NH, OH, and HF bonds cause these to be strong forces. Q- Calculate the  EN for HCl and H 2 O A- HCl = 2.9-2.1 = 0.8, H 2 O = 3.5-2.1 = 1.4 4. The fact that H is so small it can get in close to O, F and N and form short bonds. The shorter the bond, the stronger.

20. Liquid water has a partially ordered structure in which hydrogen bonds are constantly being formed and breaking up. The strong hydrogen bonds give water high boiling point and high cohesiveness - surface tension.

21. When small quantities of water are put onto a non-soluble surface and the water pulls into a spherical shapes (drops of water) due to strong hydrogen bonding.

23.  Hydrogen bonding attracts water molecules and other polar molecules by “pulling” them up

24. -Hydrogen bonding allows water to absorb a large amount of heat before its temperature increases drastically

25. In ice each each molecule is hydrogen bonded to 4 other molecules. Because of ordered structure in ice there are less H 2 0 molecules in a given space of volume.

26. Rigid Structure of Ice

27. Glucose (which has hydrogen bonding) is very soluble in water. Cyclohexane (which only has dispersion forces) is not water- soluble.

28.  What type of intermolecular bonding is involved largely depends on two main factors: 1.whether the bonds within a single molecular are polar or not (an unequal distribution of charge between two atoms involved in a chemical bond due to an unequal sharing of electrons), and 2.the overall shape of the molecule (it's molecular architecture - tetrahedral, linear, bent, etc.).

29.  A polar molecule will have one end of the molecule bearing a partial positive charge while another end carries a partial negative charge. Polar molecules must contain polar bonds.

30.  Nonpolar molecules either have no positive and negative ends, because the bonds making up the molecule are nonpolar, or because the entire outer "edge" is negative while the core of the molecule is positive (or vice versa), thus having no oppositely charged ends.

31. Electronegativity & IMFs  EN essentially defines the type of IMF.  Ionic bonds form if the  EN is 1.7 or greater.  Dipole-dipole (polar covalent) is around 0.5-1.7.  Hydrogen bonding is a type of dipole-dipole.  London forces exist in all molecules, but are especially important in non-polar covalent molecules (where  EN is less than 0.5).

32. a) NH 3 : Hydrogen bonding (H + N), London. b) SF 6 : London only (it is symmetrical). c) PCl 3 :  EN=2.9-2.1. Dipole-dipole, London. d) LiCl:  EN=2.9-1.0. Ionic, (London). e) HBr:  EN=2.8-2.1. Dipole-dipole, London. f) CO 2 : London only (it is symmetrical)