Intermolecular Forces (4.3.1)

1. Intermolecular Forces (4.3.1)

2. Dipole-Dipole Forces • Dipole-dipole forces exist between neutral polar molecules. • Polar molecules need to be close together. • There is a mix of attractive and repulsive dipole-dipole forces as the molecules tumble. • If two molecules have about the same mass and size, then dipole-dipole forces increase with increasing polarity.

3. Dipole-Dipole Forces

4. Dipole-Dipole Forces

5. London Dispersion (van der Waals) Forces • Weakest of all intermolecular forces. • It is possible for two adjacent neutral molecules to affect each other. • The nucleus of one molecule (or atom) attracts the electrons of the adjacent molecule (or atom). • For an instant, the electron clouds become distorted. • In that instant a dipole is formed (called an instantaneous/temporary dipole).

6. One instantaneous dipole can induce another instantaneous dipole in an adjacent molecule (or atom). • The forces between instantaneous dipoles are called London dispersion forces.

7. Polarizability is the ease with which an electron cloud can be deformed. • The larger the molecule (the greater the number of electrons) the more polarizable. • London dispersion forces increase as molecular weight increases. • London dispersion forces exist between all molecules. • London dispersion forces depend on the shape of the molecule.

8. The greater the surface area available for contact, the greater the dispersion forces. • London dispersion forces between spherical molecules are lower than between sausage-like molecules.

9. London Dispersion Forces

10. London Dispersion Forces

11. Hydrogen Bonding • Special case of dipole-dipole forces. • By experiments: boiling points of compounds with H-F, H-O, and H-N bonds are abnormally high. • Intermolecular forces are abnormally strong.

12. H-bonding requires H bonded to an electronegative element (most important for compounds of F, O, and N). • Electrons in the H-X (X = electronegative element) lie much closer to X than H. • H has only one electron, so in the H-X bond, the + H presents an almost bare proton to the - X. • Therefore, H-bonds are strong.

14. Hydrogen Bonding

15. Hydrogen bonds are responsible for: • Ice Floating • Solids are usually more closely packed than liquids; • Therefore, solids are more dense than liquids. • Ice is ordered with an open structure to optimize H-bonding. • Therefore, ice is less dense than water. • In water the H-O bond length is 1.0 Å. • The O…H hydrogen bond length is 1.8 Å. • Ice has water molecules arranged in an open, regular hexagon. • Each + H points towards a lone pair on O.

18. Ionic Compound – contains positive ions (usually metals) and negative ions. Held together by ionic bonds. • Covalent Compound – contains 2 or more nonmetals (no ions) • Nonpolar – contains only London dispersion forces (LDF) • Polar – contains LDF and dipole-dipole forces • Polar with H bonded to N, O, or F (with unshared pair) – contains LDF, dipole-dipole forces, and hydrogen bonds. • Larger molecule, stronger LDF (all other factors equal) • Mixture – ions mixed with polar molecules – contains ion-dipole forces (very strong) – like Na+ and Cl- ions in water. Summary of Intermolecular Forces

19. Examples – determine the types of forces present in each: • H2O • CCl4 • SO2 • LiF • Ca(NO3)2 aqueous solution • HF • PCl3

20. Examples – determine the types of forces present in each: • H2O LDF, dipole-dipole, H-bonds • CCl4 LDF • SO2 LDF and dipole-dipole • LiF ionic bonds • Ca(NO3)2 aqueous solution ion-dipole forces • HF LDF, dipole-dipole, H-bonds • PCl3 LDF and dipole-dipole

21. Bonds (covalent, ionic, metallic) • Hydrogen bonds • Dipole-dipole forces • London dispersion forces Relative Strengths of Forces