2.  intermolecular forces are forces of ATTRACTION between covalent molecules  Compared to covalent and ionic bonds, they are very weak – but when there are many, they add up to a significant force  there are three different intermolecular forces: vanderWaal’s (or dispersion) forces, dipole-dipole forces and hydrogen bonding  These forces have an effect on the boiling point and the solubility of substances.

3. vanderWaal’s forces  vanderWaal’s forces of attraction occur when the protons in one atom or molecule attract the electrons in a neighbouring atom or molecule.  Since all particles have protons and electrons, all substances have vanderWaal’s forces  Larger molecules have more protons and electrons, and so have stronger vanderWaal’s forces.

4. vanderWaal’s forces of nonpolar hydrocarbons affect boiling points: When comparing the boiling points of hydrocarbons (non-polar molecules), we see that the boiling point increases as the number of carbons increases. Why is this?

5. vanderWaal’s forces Instantaneous dipole: Induced dipole: Eventually electrons are situated so that tiny dipoles form A dipole forms in one atom or molecule, inducing a dipole in the other

6. Instantaneous & Temporary Dipoles images.tutorvista.com

7.  What trends do you notice? http://wps.prenhall.com/wps/media/objects

8. The boiling points (and melting points) of the halogens INCREASES as their MOLAR MASS increases (ie DOWN the group)

10. vanderWaal’s forces increase as molar mass increases  The reason that the boiling points increase as you go down the group is that as the molar mass increases, the number of electrons increases, and so also does the radius of the atom. The more electrons you have, and the more distance over which they can move, the bigger the possible temporary dipoles and therefore the stronger the vanderWaal’s forces of attraction  This means MORE HEAT ENERGY is required to move the particles apart (ie change from liquid to gas, as in boiling). So boiling points RISE as molar mass RISES. 

11.  Occurs only in polar molecules that have hydrogen and at least one of the following atoms: N, O or F.  These highly electronegative atoms have lone pairs of electrons which are attracted to the hydrogen atoms in neighbouring molecules.  These hydrogen atoms are essentially a proton

12.  Polar substances have a slightly electronegative end and a slightly electropositive end.  Dipole-dipole forces occur when oppositely charged poles momentarily attract one another

13. Polar and non-polar molecules wps.prenhall.com

14. 2. Dipole - Dipole attractions  Neighbouring molecules are attracted to one another through dipole-dipole forces of attraction  the greater the ∆E N of the dipole moment, the stronger the dipole-dipole forces of attraction between molecules, and the higher the boiling point, because it takes more energy to move molecules apart. ++ –– ++ –– ++ –– ++ ––

15. The greater the ∆E N of the dipole moment, the stronger the dipole-dipole forces of attraction between molecules  H-Cl  E N = 0.8 Strongest dipole Highest Boiling Point H-Br  E N = 0.7 H-I  E N = 0.4 Weakest dipole Lowest Boiling Point

16. 3. Hydrogen-bonding  Compare the  E N for H-Cl and O-H bond in H 2 O H-Cl = 2.9-2.1 = 0.8, O-H = 3.5-2.1 = 1.4  The high  E N of N-H, O-H, and H-F bonds cause hydrogen bonding forces to be the strongest IMF (about 5x stronger than normal dipole-dipole forces)

17. NOTE: the hydrogen “bond” is between molecules of water – it is not a TRUE bond!! www.landfood.ubc.ca/courses

18. So What?  The strength of INTERMOLECULAR FORCES of attraction between neighbouring COVALENT molecules affects:  Boiling and melting points  Solubility in polar or nonpolar solvents

19. www.goiit.com

20. Trends in Boiling Points Boiling points increase down a group (as period increases) for two reasons: 1)  EN tends to increase and 2) size increases. A larger size means greater vanderWaal’s forces. Boiling points are higher than expected for H 2 O, HF, and NH 3 because these are capable of hydrogen-bonding with neighbouring molecules (high  EN), creating very strong intermolecular forces. This makes it more difficult to move molecules apart in the evaporation process, resulting in a high boiling point.

21. Remember: HYDROGEN BONDING Is stronger than DIPOLE-DIPOLE ATTRACTIVE FORCES which are stronger than VANDERWAAL’S FORCES

22.  Water is not an organic molecule but is essential for life on this planet  All cells are surrounded inside and out with water – anything that interacts with a cell must first be dissolved in water  Physical properties: ◦ colourless and transparent ◦ liquid at room temperature ◦ density = 1.0 g/mL ◦ m.p. = 0℃b.p = 100℃  water has LD, D-D forces, and H-bonding

23.  Water has cohesive properties – the high number of intermolecular forces causes water molecules to ‘stick’ together Examples: ◦ surface tension – beading of water ◦ water striders – too light to break surface tension ◦ transpiration in plants – transport in xylem tubes

24.  Water has adhesive properties – it’s polar nature causes it to stick to other substances Examples: ◦ capillary action – water ‘climbs’ up small diameter tubes, or ‘bleeds’ through the microscopic pores and channels in paper or other porous substances ◦ this is due to the hydrogen bonding interactions between the water and the surface of the tube (either SiO 2 or the cellulose tubes of paper) ◦ This helps to explain the meniscus inside a tube

28.  Water has outstanding solvent properties  Used to be called the ‘universal solvent’, but this is not a good name, since not everything dissolves in water  The polar nature of water allows any other polar substance or any charged particle to dissolve easily  The δ - will attract the δ + end of solutes, and this attraction will remain once the solute is dissolved.  The same is true for ionic substances – the cation will be attracted to the δ - end of water, and the anion will be attracted to the δ + of water.

30.  Water’s density decreases as it changes from liquid to solid.  This is because the distance between molecules in a crystal lattice (as ice) on average further than when in a liquid. Special Properties of water