1. A Closer Look at Physical Properties Thermochemistry: States of Matter Thermochemistry:

2. 2 Thermochemistry -Energy of Chemical Reactions -OR- -The study of heat changes that occur during chemical reactions and physical changes of state

3. 3 Energy and Chemistry ENERGY is the capacity to do work or transfer heat. can be: light, electrical, kinetic, potential, chemical HEAT (represented by q) is the form of energy that flows between 2 samples because of a difference in temperature – will always go from a warmer object to a cooler one. WORK is the form of energy that results in a macroscopic displacement of matter such as gas expansion or motion of an object (force x distance) LAW OF CONSERVATION OF ENERGY energy is neither created nor destroyed

4. 4 Energy and Chemistry CALORIE quantity of heat that raises the temperature of 1g of pure water 1°C JOULE SI unit of heat and energy (1kcal = 4186J and 1J = 0.239 cal) HEAT CAPACITY the amount of heat it takes to change an object’s temperature by exactly 1°C

5. A. Kinetic Molecular Theory  KMT Particles of matter are always in motion. The kinetic energy (speed) of these particles increases as temperature increases. Kelvin Temperature scale represents the relationship between temperature and average kinetic energy. This relationship between temp. and energy is directly proportional Example: particles of He at 200K have twice the average KE as particles of He at 100K

6. CHEMICAL ENERGY Chemical Potential Energy: Energy stored in chemicals because of their compositions Chemical bonds are a source of energy BOND BREAKING - requires energy BOND MAKING - releases energy In a chemical reaction : if more energy is released in forming bonds than is used in breaking bonds then EXOTHERMIC... reaction is EXOTHERMIC if more energy is used in breaking bonds than is released in forming bonds then... reaction is ENDOTHERMIC Energy is released as HEAT, LIGHT, SOUND, WORK Energy can be provided by - LIGHT - photochemistry - WORK - electrochemistry - COOLING of surroundings

7. 7 Specific Heat Capacity 1°C (measured in J/(g x °C) A.K.A. Specific Heat: (represented by c) the amount of heat it takes to raise temperature of 1g 1°C (measured in J/(g x °C) A difference in temperature leads to energy transfer. Specific heat capacity= heat lost or gained by substance (J) (mass, g) (T change, K)

8. 8 Specific Heat Capacity SubstanceSpec. Heat (J/gK) H 2 O4.184 Al0.902 glass0.84 Aluminum Water

9. Specific Heat  q = mc  T  q = heat (j)  m = mass (g)  c = specific heat (j/g  C)   T = change in temperature = Tf-Ti (  C)

10. A. Kinetic Molecular Theory - KMT  Particles of matter are always in motion. The kinetic energy (speed) of these particles increases as temperature increases. Kelvin Temperature scale represents the relationship between temperature and average kinetic energy. This relationship between temp. and energy is directly proportional Example: particles of He at 200K have twice the average KE as particles of He at 100K

11. A. Liquids vs. Solids SolidLiquidGas IMF Strength (Intermolecular Force) Very StrongStronger than Gasses Very weak FluidNoYes DensityHigh lower than solids Low CompressibleNo Yes Diffusion RateExtremely slow Slow slower than gasses High

12. B. Liquid Properties  Surface Tension attractive force between particles in a liquid that minimizes surface area

13. B. Liquid Properties  Capillary Action attractive force between the surface of a liquid and the surface of a solid watermercury

14. Evaporation  Definition: the conversion of a liquid to a gas or vapor below its boiling point (a.k.a. vaporization)  What happens… Molecules at the surface break away and go into the gas or vapor state The liquid will evaporate faster as it is heated because KE increases Particles with the highest KE escape first Evaporation in a closed container is different

15.  Particles collide with the walls of the sealed container and produce a vapor pressure above the liquid.  As the container stands, the number of particles turning into vapor increases  Some of the particles will condense and turn to liquid  After some time vapor particles condensing will equal particles vaporizing to create an equilibrium Evaporation in a Closed Container

16. Boiling Point  Definition: the temperature at which the vapor pressure of a liquid is just equal to the external pressure Bubbles = pockets of vapor forming throughout the liquid At lower atmospheric pressures, the boiling point decreases b/c escaping particles are less likely to collide w/external air molecules  The temperature of a boiling liquid never rises above its boiling point With more heat, the liquid will only boil faster

17. C. Types of Solids  Crystalline - repeating geometric pattern covalent network metallic ionic covalent molecular  Amorphous - no geometric pattern decreasing m.p.

18. C. Types of Solids Ionic (NaCl) Metallic

19. C. Types of Solids Covalent Molecular (H 2 O) Covalent Network (SiO 2 - quartz) Amorphous (SiO 2 - glass)

20.  Amorphous solids lack internal structure – atoms are randomly arranged  Allotropes are two or more different molecular forms of the same element in the same physical state (ie: diamond or graphite are allotropes of carbon)  Glasses are amorphous solids that are sometimes called supercooled liquids C. Types of Solids

21. Sublimation  Change of a substance from a solid to a gas or vapor w/o passing through the liquid state Happens with some solids when the vapor pressure is very high Examples: dry ice, iodine

22. Phase diagrams  A single graph that represents the relationships among the solid, liquid, and vapor states or phases of a substance in a sealed container

23.  Given a temperature and pressure, phase diagrams tell us which phase will exist.  Features of a phase diagram: Triple point: temperature and pressure at which all three phases are in equilibrium. Critical point: critical temperature and pressure for the gas. Phase diagrams

24. B. Heating Curves Melting - PE  Solid - KE  Liquid - KE  Boiling - PE  Gas - KE 

25. Copyright 1999, PRENTICE HALLChapter 1125 B. Heating Curves

27.  Temperature Change change in KE (molecular motion) depends on heat capacity During a phase change, adding heat causes no temperature change.

28. Copyright 1999, PRENTICE HALLChapter 1128 Phase Changes Sublimation: solid  gas. Vaporization: liquid  gas. Melting or fusion: solid  liquid. Deposition: gas  solid. Condensation: gas  liquid. Freezing: liquid  solid.

29. Heating Curves  Phase Change change in PE (molecular arrangement) temp remains constant until the phase change is complete  Molar Heat of Fusion (  H fus ) energy required to melt 1 mole of a substance at its m.p. Melting is an endothermic process + ΔH

30. Heating Curves  Molar Heat of Solidification ΔH solid Heat change that occurs when 1 mole of a substances changes from liquid to solid Exothermic process; -ΔH Heat lost is equal to heat gained during melting

31. B. Heating Curves  Heat of Vaporization (  H vap ) energy required to boil 1 mole of a substance at its b.p. Endothermic; +ΔH usually larger than  H fus …why?  EX: sweating, steam burns, the drinking bird

32. Heating Curves  Molar Heat of Condensation Heat released when one mole of a substance changes from gas to liquid Exothermic; -ΔH Heat released is equal to the heat gained during boiling Ex. Steam burns

33. Heat and Changes of State  Heat of combustion (∆H)= the heat of reaction for the complete burning of one mole of a substance  Molar heat of fusion (∆H fus )= the heat absorbed by one mole of a substance in melting from a solid to a liquid at a constant temperature  Molar heat of solidification (∆H solid )= heat lost when one mole of a liquid freezes to a solid at a constant temperature (equal to the negative heat of fusion)  Molar heat of vaporization (∆H vap )= the heat absorbed by one mole of a substance in vaporizing from liquid to a gas  Molar heat of condensation (∆H cond )= heat released by one mole of a vapor as it condenses