Intramolecular and Intermolecular Forces

1. Intramolecular and Intermolecular Forces • Intramolecular Forces - The forces between atoms or ions inside a compound such as covalent and ionic bonding (determines the stability of a comound) • H-O-H  2H+ O E = 920 kJ/mol • Intermolecular Forces - The forces between individual particles. (determines the phase a compound is in at room temperature) • H2O(l ) H2O(g) E=40.7 kJ/mol • note: Ionic bonding falls under both categories.

2. Ion-Ion Intermolecular Forces • E q+q-)/d • ionic compounds have strong Intermolecular forces • solids at room temperature and they have high m.p. • ionic compounds with mulitiply charged ions generally have higher m.p. than those with singly charged ions • NaCl mp 801oC • CaF2 mp 1180oC • CaO mp 2580oC

3. Ion-Dipole & Dipole-Dipole intermolecular forces Polar covalent molecules undergo Dipole-Dipole intermolecular interactions which are much weaker than ion-ion intermolecular interactions (E q+q-)/d2). Therefore, polar covalent molecules are generally liquids at room temperature. Ion-Dipole Dipole-Dipole

4. Hydrogen Bonding • Hydrogen bonding occurs among polar covalent molecules containing a hydrogen bonded to F, O or N.

5. Dispersion Forces Dispersion forces - weak attractive forces that all molecules have. dispersion forces are the only intermolecular interaction that small, non-polar molecules undergo. (generally gases at room temp.) The more electrons a molecule/atom has, the stronger the dispersion force

6. Intermolecular Force state • (room temp.) • Strongest Ion-Ion Solid • Ion-Dipole liquid • intermediate hydrogen bonding liquid • Dipole-Dipole liquid • weakest dispersion gas

7. Liquid State - Viscosity • Viscosity - the resistance to flow of a liquid. • Viscosity depends on the strength of the intermolecular forces. • High viscosity - slow flowing or thick: Stronger intermolecular forces • Low viscosity - fast flowing or thin: Weaker intermolecular forces • Temperature increases, viscosity decreases

8. Liquid State - Surface Tension • Surface Tension - a measure of the inward forces that must be overcome to expand the surface area of a liquid. • The molecules on the surface of a liquid only have intermolecular interactions with the molecules toward the interior of the liquid.

9. Liquid State - capillary action • Cohesive Forces - all forces that hold a liquid together. • Adhesive Forces - the force of attraction between a liquid and another surface. • Glass Capillary Tube - a glass tube with a small bore (inner diameter) • Water has a strong attraction to glass (SiO2) therefore water adheres to the glass. There is a strong adhesive force between water and glass. Waters meniscus will be concave in a glass capillary tube. • Mercury does not have a strong attraction to glass and does not adhere to it. Therefore, the cohesive forces in mercury are stronger than the adhesive force between mercury and the glass. Mercury’s meniscus will be convex in a glass capillary tube. (Fig. 13-9) • Capillary action - The movement of a liquid in a capillary tube when the end of the tube is placed in the liquid.

10. Liquid State - Evaporation • Evaporation (vaporization) - the endothermic process by which molecules on the surface of a liquid break away and go into the gas phase. • A molecule must have sufficient kinetic energy to break the intermolecular forces holding it in the liquid phase. (Figure 13-11) • At higher temperatures, more molecules have sufficient kinetic energy to escape and the liquid will evaporate at a faster rate. • Liquid  vapour : evaporation :endothermic • vapour liquid : condensation : exothermic

11. Liquid - vapour pressure • Vapour Pressure - The partial pressure of vapour molecules above the surface of a liquid at equilibrium at a given temperature is the vapour pressure (vp) of the liquid at that temperature. • vapour pressure increases with increasing temperature (Figure 13-13) • volatile liquids have high vapour pressures (weak intermolecular forces) • liquids with low vapour pressures have stronger intermolecular forces

12. Liquid State - Boiling Points • Boiling point - the temperature at which a liquids vapour pressure is equal to the external pressure. (Normal boiling point is the boiling point when the external pressure is 760 torr (1 atm) )

13. Liquid State - Distillation • If a mixture of liquids have sufficiently different boiling points than they can be separated by the process of distillation • (Figure 13-14 pg 498)

14. Liquid State - Heat Transfer • When you add energy (heat) to a system: • 1) You can change the Kinetic Energy (Temp) • or • 2) You can change the Potential Energy (Phase) • You can NOT do both... • Specific heat (Csp (J g-1oC-1)) or molar heat capacity (J mol-1oC) of a liquid - the amount of heat (J) necessary to raise the specified mass one degree celsius/kelvin. (Involves a change in K.E./Temp.) • Enthalpy (heat) of vaporization - Hvap - The amount of energy (kJ) necessary to convert one mole of liquid to vapour. • Liquid + heat  vapour Hvap • vapour liquid + heat -Hvap

15. Solid State - melting point and sublimation • Melting point - the temperature at which a substances solid and liquid phase coexist in equilibrium. (Normal melting point - melting point @ 760 torr/ 1 atm) • Solid  liquid melting endothermic Hfus • liquid  solid freezing exothermic -Hfus • Hfus= 6.02 kJ/mol for water • 6.02 kJ of heat are required to melt one mole of water at 0oC and 1 atm. • 6.02kJ of heat would be released if one mole of water was frozen at 0oC and 1 atm. • Enthalpy of sublimation (deposition) - Hsub - the energy necessary to vaporize one mole of a solid at 1 atm.

16. the three states of matter • page 505 in the book • Phase diagrams - go to saunders 13-17. Picture is in your book also figure 13-18