Energy and Phase Changes

1. Energy and Phase Changes

2. Energy and Phase Changes • Three states or phases of matter • Solid • Liquid • Gas • Energy is involved in every phase change. • Some require energy to occur (endothermic) • Some release energy (exothermic) • Based on the breaking or formation of bonds between the particles (atoms, molecules, fun)

3. Intermolecular Forces Differing attractive forces between molecules cause some materials to be solids, some to be liquids, and some to be gases at the same temperature. Section 12-2

4. Gases • Kinetic Molecular Theory • Constant random motion • High energy movement spreads molecules throughout container. • Large space between gas molecules • Elastic collisions • Little or no intermolecular attraction or repulsion. Section 12-3

5. Liquids Kinetic molecular theory still involved, but intermolecular forces are strong enough to greatly limit movement. Forces of attraction keep molecules closely packed in a fixed volume, but not in a fixed position. Molecules can flow past each other. Liquids are much denser than gases because of the stronger intermolecular forces holding the particles together. Section 12-3

6. Solids Solids: particles’ attractive intermolecular forces are greater than their kinetic energy of motion. • Particles in a solid vibrate in a fixed position. • Most solids are more dense than their liquid. • Note: Ice is not more dense than liquid water. • It FLOATS! Section 12-3

7. Solids (cont.) Crystalline solidsare solids with atoms, ions, or molecules arranged in an orderly, geometric shape. Section 12-3

8. Heating Curve for Water http://media.wwnorton.com/college/chemistry/chemtours/interface.asp?chapter=chapter_05&folder=heating_curves

9. Phase Changes • Phase changes that require energy to occur: • Melting • Vaporization • Sublimation • Phase changes that release energy: • Freezing • Condensation • Deposition Section 12-4

10. Phase Changes Involve Heat Transfer Heat is the transfer of energy from an object at a higher temperature to an object at a lower temperature. Section 12-4

11. Phase Changes That Require Energy (cont.) When ice is heated, the ice eventually absorbs enough energy to break the hydrogen bonds that hold the water molecules together. When the bonds break, the particles move apart and ice melts into water. Themelting point of a crystalline solid is the temperature at which the forces holding the crystalline structure together are broken and it becomes a liquid. Section 12-4

12. Phase Changes That Require Energy (cont.) Particles with enough energy escape from the liquid and enter the gas phase. Section 12-4

13. Phase Changes That Require Energy (cont.) Vaporizationis the process by which a liquid changes to a gas or vapor. Evaporationis vaporization only at the surface of a liquid. Section 12-4

14. Phase Changes That Require Energy (cont.) In a closed container, the pressure exerted by a vapor over a liquid is called vapor pressure. Section 12-4

15. Phase Changes That Require Energy (cont.) The boiling point is the temperature at which the vapor pressure of a liquid equals the atmospheric pressure. Section 12-4

16. Phase Changes That Require Energy (cont.) Sublimationis the process by which a solid changes into a gas without becoming a liquid. Dry ice = solid CO2 changing directly to gas. Section 12-4

17. Phase Changes That Release Energy As heat flows from water to the surroundings, the particles lose energy. Thefreezing point is the temperature at which a liquid is converted into a crystalline solid. The bonds that are formed create a more stable situation and excess energy is released. Section 12-4

18. Phase Changes That Release Energy (cont.) As energy flows away (cooling) from water vapor, the velocity decreases. The process by which a gas or vapor becomes a liquid is called condensation. Deposition is the process by which a gas or vapor changes directly to a solid, and is the reverse of sublimation. Section 12-4

19. Heat Energy of Phase Changes • Heat of fusion (Hf) = heat absorbed when a solid melts • Energy needed to break bonds (endothermic) • Heat of solidification or crystallization (Hs) = heat released when a liquid freezes (= - Hf ) • Energy released from the formation of bonds (exothermic) • Heat of vaporization (Hv) = heat absorbed when a liquid changes into a gas • Energy needed to break bonds • Heat of condensation (Hc) = heat released when a gas condenses to become a liquid (= - Hv ) • Energy released from the formation of bonds

20. Heating Curves - Where is the energy going? • Heat of vaporization: • Heat of fusion

21. Heat Energy of Phase Changes q = c x m x ∆T Leads to: q = m x Hf (or Hv, Hs, Hc) For water (memorize): Hf = 80 cal/g or 6.02 kJ/mol (melting/freezing) Hv= 540 cal/g or 40.6 kJ/mol (boiling/condensing) C liquid water = 1.00 cal/g*deg C or 4.184J/ g*deg C C ice = 0.50 cal/g*deg C or 2.1 J/ g*deg C C water vapor = 0.50 cal/g*deg C or 2.1 J/ g*deg C

22. Phase Diagrams A phase diagram is a graph of pressure versus temperature that shows in which phase a substance will exist at various T and P. Section 12-4

23. Phase Diagrams (cont.) The triple point is the temperature and pressure at which all three phases of a substance can coexist. Above the critical point temperature the substance can no longer be found in a liquid state at any P. Section 12-4

24. Phase Diagrams (cont.) The phase diagram for different substances are different from water. Section 12-4