1. An Introduction to Thermodynamics.

2. I. Thermodynamics A branch of physical science dealing with heat transfer and temperature.

3. II. Three Phases of Matter

4. A. Solids *Molecules are tightly packed. *High intermolecular forces. *Molecules are vibrating in place. *Have a fixed shape and volume.

5. B. Liquids Molecules are packed. Intermediate intermolecular forces. Molecules are moving around. Have a fixed volume but takes shape of its container.

6. C. Gases Molecules are far apart. No intermolecular forces. Molecules are moving around a lot. No fixed volume (can be compressed) nor shape (takes shape of container).

7. III. Phase Changes A. You are probably aware water can exist in 3 phases  Solid – ice  Liquid – water  Gas – steam/water vapor All of the elements can exist in the 3 phases too, it just depends on their temperature and the external pressure (oxygen is a solid at -220 0 C, gold is a gas at 2820 0 C)

8. Phase Changes B. Solid to liquid  Melting/fusion C. Liquid to gas  Boiling/vaporization/evaporation D. Solid to gas  Sublimation

9. Phase Changes Or the reverse E. Gas to liquid  Condensation F. Liquid to solid  Freezing

10. IV. Phase Changes A. A substance can be converted from a solid to a liquid or a liquid to a gas by  Adding heat  Reducing pressure (liquid to gas) B. A substance can be converted from a gas to a liquid or a liquid to a solid by  Removing heat  Increasing pressure (gas to liquid)

11. heat

12. V. What is temperature? A. Temperature is a measure of the kinetic energy of molecules.  The faster molecules are moving, the higher the temperature. High temperature Low temperature

13. B. Units of Temperature 1. Celsius ( 0 C) – based on the freezing and boiling point of water  a. O 0 C = Water freezes  b. 100 0 C = Water boils C. Temperature is measured with a thermometer

14. VI. What is heat? A. Heat energy (or just heat) is a form of energy which transfers among particles in a substance (or system) by means of kinetic energy of those particles. In other words, under kinetic theory, the heat is transferred by particles bouncing into each other.

15. B. How is heat measured? With a calorimeter Units of heat Joules (J) Calories (cal) (1 calories = 4.18 J = the energy needed to raise the temperature of 1 gram of water by 1°C)

16. VII. What’s the difference between temperature and heat? Note that temperature is different from heat, though the two concepts are linked.heat A. Temperature is a measure of the internal energy of the system, while heat is a measure of how energy is transferred from one system (or body) to another. B. The greater the heat absorbed by a material, the more rapidly the atoms within the material begin to move, and thus the greater the rise in temperature.

17. VIII. Heat explained. A. So things that are hot have high temperatures and lots of heat energy! B. Things that are cold have low temperature and little heat energy!  In other words, something is not really cold, it just has less heat energy

18. IX. Heat transfer A. Heat moves from hot to cold  1. The heat from the water is moving into the ice causing it to melt  2. The heat from the water gets transferred to the ice, so the water cools down and ice warms up

19. B. What does the addition of heat do? 1. Adding heat (which is energy) increases the molecular kinetic energy 2. Adding heat therefore increases temperature (usually ).

20. X. A Phase Change Diagram A B At A the solid is cold, but it absorbs heat energy and at point B is hotter, but it is still a solid!

21. XI. A Phase Change Diagram A BC What has happened to the temperature between these points? So what happens from point B to C to create a flat line? What is happening to the heat energy between these points? Stops rising Continues to be added

22. Where is all the heat going? XII. A Phase Change Diagram Add rest of labels to your diagram The added heat is being used to change the phase of the substance!

23. XIII. During a phase change, A. The heat energy is being used to break down the IMF’s that keep the particles hanging out together instead of making the particles vibrate faster. Red arrows represent the added heat making the particles vibrate faster as the solid heats up. Red arrows represent the added heat making the particles separate (breaking down the imf’s) as the solid begins to melt

24. XV. Heat Calculations How many joules (change calories to joules in your notes)(how much heat) does it take to heat 3 grams of ice from -30 0 C to -10 0 C? Use the formula q = mCpΔT  q = heat (cal or J)  m = mass (g)  Cp = the specific heat of the substance (Specific heat is the amount of heat needed to raise the temperature of one gram of a substance by 1°C).heat temperature  ΔT = the change in temperature (final temperature – initial temperature or T f – T i )

25. How many joules(how much heat) does it take to heat 3 grams of ice from -30 0 C to -10 0 C? Use the formula q = mCpΔT Mass = ΔT = (-10) – (-30) = +20 o C q = (3) (4.184) (20) = So this formula is used when there is NO PHASE CHANGE!! Just something heating up or cooling down. 3g Cp = 4.184 J/g ( o C) 251.04 j

26. XVI. How do you calculate heat when the material is going through a phase change and there is no temperature change? A. For melting/fusion/freezing use q=mHf q = heatm = mass Hf = heat of fusion (a number you will be given)

27. B. For boiling/vaporization/condensation use q=mHv q = heatm = mass Hv = heat of vaporization(a number you will be given)

28. XVII. Sample heat calculation of a substance going through a phase change. A. How much is needed to boil 15 g of water? Use q=mHv Hv = 2260 J/g (for water) q = (15) (2260) = 33 900 J

29. Numbers you will need to do worksheets! Also refer to reference page XVIII. Cp’s are

30. XIX. Heat of fusion and vaporization For water  Hf = 334 J/g (so it takes 334 J to melt 1 gram of water)  Hv = 2260 J/g (it takes 2260 J to melt 1 gram of water)  You may notice on the phase change diagram that the line showing vaporization is longer than line for melting

31. XX. Endothermic vs Exothermic A. Endothermic (heat goes “en”) – absorbs heat to happen  Q (heat) will be + (temp goes up so ΔT is a positive #) B. Exothermic (heat “ex”its) – heat must be released from the substance.  Q (heat) will be - (temp goes down so ΔT is a negative #)

32. You decide – exothermic or endothermic ! Ice melting into water – Steam condensing into water – Sweat evaporating – (this why it cools you – the heat from your skin makes the sweat evaporate, so as the sweat evaporates it takes the heat from your skin) endothermic exothermic endothermic

33. Endothermic Exothermic