1. Phase Changes & Solids AP Chemistry

2. Phase Changes

3. Energy Changes Associated with Changes of State Heat of Fusion: Energy required to change a solid at its melting point to a liquid. (solid = liquid vapor pressure)

4. Energy Changes Associated with Changes of State Heat of Vaporization: Energy required to change a liquid at its boiling point to a gas (liquid vapor pressure = atmospheric/external pressure).

5. Energy Changes Associated with Changes of State The heat added to the system at the melting and boiling points goes into pulling the molecules farther apart from each other. The temperature of the substance does not rise during the phase change.

6. Vapor Pressure At any temperature, some molecules in a liquid have enough energy to escape. As the temperature rises, the fraction of molecules that have enough energy to escape increases. Higher temperature = higher vapor pressure

7. Vapor Pressure As more molecules escape the liquid, the pressure they exert increases.

8. Vapor Pressure The liquid and vapor reach a state of dynamic equilibrium: liquid molecules evaporate and vapor molecules condense at the same rate.

9. Vapor Pressure The boiling point of a liquid is the temperature at which its vapor pressure equals atmospheric pressure. The normal boiling point is the temperature at which its vapor pressure is 760 torr.

10. Evaporation or Boiling? Evaporation only occurs on the surface Boiling occurs within the body of the liquid when vapor pressure = atm pressure Cool Demo!!

11. Types of Solids

12. Solids We can think of solids as falling into two groups: ◦ Crystalline—particles are in highly ordered arrangement.

13. Solids ◦ Amorphous—no particular order in the arrangement of particles. (glass)

14. Types of Bonding in Crystalline Solids

15. Covalent-Network and Molecular Solids Diamonds are an example of a covalent-network solid in which atoms are covalently bonded to each other (sp 3 hybridization, good insulator). ◦ They tend to be hard and have high melting points. ◦ Silica, silicates: form glass, ceramics, semi-conductors ◦ Diamonds can be formed from graphite under high pressures 15,000 atm @ 2800˚C

16. Covalent-Network and Molecular Solids Graphite is an example of a molecular solid in which atoms are held together with van der Waals forces (sp 2 rings, good conductor & lubricant). ◦ They tend to be softer and have lower melting points.

17. Metallic Solids Metals are not covalently bonded, but the attractions between atoms are too strong to be van der Waals forces. In metals, valence electrons are delocalized throughout the solid. Bonding is strong & nondirectional. Easy to move metals but not to separate them.

18. Physical Properties of Metals Conduct heat and electricity. Malleable (can be pressed or hammered into sheets). Ductile (can be drawn into wire). Atoms can slip past each other. ◦ So metals aren’t as brittle as other solids.

19. Electron-Sea Model Metals can be thought of as cations suspended in “sea” of valence electrons. Attractions hold electrons near cations, but not so tightly as to impede their flow.

20. Electron-Sea Model This explains properties of metals— ◦ Conductivity of heat and electricity ◦ Deformation

21. Alloys Mixtures of elements that have properties characteristic of metals. Many ordinary uses of metals involve alloys.

22. Solution Alloys Components of alloys are dispersed uniformly— ◦ In substitutional alloys, solute particles take place of solvent metal atoms. ◦ Particles close in size. ◦ Brass, sterling silver, pewter, plumber’s solder

23. Solution Alloys Components of alloys are dispersed uniformly. ◦ In interstitial alloys, solute particles find their way into holes between solvent metal atoms. ◦ Solute particles smaller than solvent. ◦ Steel = iron with carbon to prevent rusting!