Liquids, Solids, and Intermolecular Forces

1. Liquids, Solids, and Intermolecular Forces Chapter 11

2. Gecko’s Hairy Feet • Nanostructures on the soles of gecko feet. Thanks to about one billion hierarchically organized nanohairs, the gecko can go for a walk on walls and ceilings, unlike people. • Image: Max Planck Institute for Metals Research

4. Polar Molecules • Dipole - A molecule such as HF which has a positive and a negative end. This dipolar character is often represented by an arrow pointing towards the negative charge. • Dipole moments – the measure of the net molecular polarity • Measure of separation of charge • Measured in units of Debyes (D) = Qr (charge x separation)

6. Polar bondsNon-polar molecules

7. Forces holding one molecule to another in a substance. van der Waals forces Dispersion forces London Forces Polar-polar interactions Hydrogen bonding Intermolecular Forces

9. London forces Induced dipole – induced dipole or Nonpolar - Nonpolar interactions

10. Polarizability • The ease with which a molecule/atoms electron cloud can be distorted, thereby inducing a dipole moment. • Increasing the number of electrons increases the polarizability of an atom or molecule.

11. How does polarity affect molecular properties?

15. Dipole – Dipole forces or Polar - Polar interactions

18. Hydrogen Bonds • A special type of polar interaction between a hydrogen atom bonded to an electronegative element and another electronegative element.

26. Solubility Like dissolves Like

27. Solubility • Polar solvents dissolve polar molecules • Nonpolar solvents dissolve nonpolar molecules • Molecules with polar and nonpolar ends are frequently soluble in both polar and nonpolar solvents. • Polar solvents are good for solubilizing salts.

29. Liquids

30. Viscosity • Resistance to flow • If a liquid has strong intermolecular interactions then particles will not flow past each other easily and viscosity will be high.

33. Surface Tension • tendency to minimize surface area

34. In Orbit (Space Shuttle), water droplets are spherical

35. Cohesive Forces – attraction between molecules in a liquid • Adhesive Forces – attraction between liquid molecules and the surface of the tube • Water (red) has stronger adhesive forces. • Mercury has stronger cohesive forces.

36. Capillary Action • The ability of a liquid to flow against gravity up a narrow tube. Attraction of water to glass walls draws water up tubes

37. Vaporization Some Molecules in an open beaker have enough kinetic energy to vaporize from the surface of the liquid.

40. Vapor Pressure The pressure exerted by a vapor in equilibrium with its liquid phase.

42. A liquid boils when thermal energy is high enough to cause molecules in the interior of the liquid to become gaseous, forming bubbles that rise to the surface.

44. or

45. Clausius Clapeyron Equation or

46. Dry ice sublimes at –78oC and has a Hsub of 25.2 kJ/mol. Calculate the vapor pressure of CO2 at –100oC.

48. Calculate the boiling point of water at the summit of Pikes Peak in Colorado where the atmospheric pressure is 447.

50. Crystalline solid – atoms, ions, or molecules lie in an orderly array • typically have flat well defined surfaces called faces. • Amorphous solid – atoms or molecules lie in random jumble.