Liquids and Solids

1. Liquids and Solids KMT of Liquids and Solids, Phase Diagram, Vapor Pressure Curve, Heating/Cooling Curve

2. Intermolecular Forces: Liquids, Solids, and Phase Changes • Types of Intermolecular Forces • Properties of liquids and solids • Phase change diagrams • Heating/cooling curve • Vapor pressure curve

3. Brainteaser!!!! • If substances at the same temperature have the same kinetic energy, why are they all not liquids, solids, or gases?

4. Intermolecular Forces • Intermolecular Forces are attractive forces between molecules. Think interstate! • Intramolecular Forces are attractive forces that hold molecules together • Inter vs. Intra • 41 kJ to vaporize 1 mole of H2O • 930 kJ to break all O-H bonds in one H2O molecule Which one is stronger????? Intramolecular forces are stronger than intermolecular forces!!!!

5. Dipole – Dipole Forces • Dipole – molecule with a completely separate positively and negatively charged end • Between polar molecules • What bond is the strongest? • Where is the intermolecular bond?

6. Ion – Dipole Forces • Between polar molecules and ions • Give me an example of an everyday solution between polar molecules and ions!!!!!! • Why are dipoles attracted to ions?

7. London-Dispersion Forces • Intermolecular forces are formed by temporarily induced dipole moments • Usually occurs between identical molecules (Example H2 (g) • How do dipoles become induced? • Electron clouds constantly move and when one molecule collides with another molecule the electrons are temporarily shifted to one side • This creates a momentary negative end and a positive end

8. Hydrogen bonds • Force formed between molecules containing N–H, O–H, or F–H groups, and an electronegative O, N, or F atom. • 10% of the energy in a covalent bond!!!!!!

9. Hydrogen Bonding H2O CH3OH NH3

10. Phases of matter • Gases – molecules are widely separated and the “fluid” is compressible • Liquids – molecules are more tightly packed and liquids are relatively incompressible • Solids – molecules are tightly packed and solids are incompressible and rigid

11. Liquids • IMF’s limit the range of motion of particles in a liquid • Density – Liquids have a higher density at 25 °C than gases • Fluidity – Ability to flow • Viscosity – Measure of the resistance of a liquid to flow • Surface tension – The energy required to increase the surface area of a liquid by a given amount

12. Viscosity • Measure of a liquids resistance to flow • Inversely related to the size of the molecule and the type and strength of intermolecular forces • The higher the temperature the lower the viscosity • If temperature then viscosity • Here’s the tricky part: • If temperature then the liquid starts to flow

13. Surface Tension • The energy required to increase the surface area of a liquid by a given amount • Molecules in the center of a liquid are exposed to IMF from all sides • Molecules on the surface of a liquid are not exposed to IMF from all sides • In order to increase the surface area of a liquid the molecules in the interior of the liquid must move to the surface and the IMF’s must be broken

14. Capillary Action • Water molecules “cling” to the surface of the graduated cylinder by adhesion • Adhesion is the force of attraction between different types of molecules • Cohesion is the force of attraction between the same type of molecules • What force must be strongest for water to cling to the glass tube? • If adhesion forces are stronger than cohesion forces water will be drawn up the sides of the cylinder

15. Solids • Tightly packed molecules that are rigid and cannot be compressed • Density is highest in solids (except in water!!!) • Crystalline solid – solid whose atoms, ions, or molecules are arranged in an orderly, geometric, 3-D structure • Amorphous – atoms are randomly arranged because they typically cool too quickly. No order exists in the solid.

16. Types of Solids Crystalline – a well defined arrangement of atoms; this arrangement is often seen on a macroscopic level. (p.402) • Atomic solids • Ionic solids • Molecular solid • Covalent network • Metallic Units points that can be repeated in three dimensions to form a lattice

17. Phase Changes • Melting – the change from a solid to a liquid • Melting Point – T at which forces holding lattice together are broken • Vaporization- the change from a liquid to a gas • Sublimation – the change from a solid to a gas • Condensation – the change from a gas to a liquid • Deposition – the change from a gas to a solid • Freezing – the change from a liquid to a solid GAS SUBLIMATION VAPORIZATION DEPOSITION CONDENSATION FREEZING SOLID LIQUID MELTING

18. Phase Change Diagrams Look at the liquid solid line and its slope!!!!! • Relationship between T and P • Triple point – P and T at which substance can coexists as a gas, liquid, and solid • Critical point – T at which a substance can no longer remain a liquid regardless of the pressure

19. Phase Diagram for H2O • What is the difference between this diagram and the first? • The liquid solid line leans backwards! Normal Melting and boiling points Vapor pressure curve

20. Vapor Pressure • In a sealed container some water (l) changes phase to become water vapor and exerts a pressure over the surface of the liquid (if the container were open it would be considered partial pressure)

21. Heating Cooling Curve

22. Heating and cooling curve for H2O What bonds are broken? 120 °C steam 100 °C water  steam 50°C liquid water 0 °C ice liquid -10 °C ice Heat added  Why does temperature “stand still”?