THERMOCHEMISTRY or Thermodynamics

1. THERMOCHEMISTRYorThermodynamics

2. Energy & Chemistry ENERGY is the capacity to do work or transfer heat. HEAT is the form of energy that flows between 2 objects because of their difference in temperature. Other forms of energy — • light • electrical • kinetic and potential

3. Energy & Chemistry All of thermodynamics depends on the law of CONSERVATION OF ENERGY. • The total energy is unchanged in a chemical reaction. • If PE of products is less than reactants, the difference must be released as KE.

4. Internal Energy (E) • PE + KE = Internal energy (E or U) Int. E of a chemical system depends on • number of particles • type of particles • temperature

5. Potential & Kinetic Energy

6. Thermo-dynamics the science of heat transfer (molecular motions). Heat transfers until thermal equilibrium is established.

7. T(system) goes down T(surr) goes up (until it reaches equilibrium Directionality of Heat Transfer • Heat always transfer from hotter object to cooler one. Heat lost = heat gained • EXOthermic: heat transfers from SYSTEM to SURROUNDINGS.

8. T(system) goes up T (surr) goes down Directionality of Heat Transfer • Heat always transfer from hotter object to cooler one. Heat lost = heat gained • ENDOthermic: heat transfers from SURROUNDINGS to the SYSTEM.

9. ENERGY TRANSFER • IDENTIFY SURROUNDINGS AND SYSTEM. • EXOTHERMIC OR ENDOTHERMIC ?

10. WORK?

11. Heat is NOT temperature The increased volume with temperature causes the mercury to rise

12. James Joule 1818-1889 UNITS OF ENERGY 1 calorie = heat required to raise temp. of 1.00 g of H2O by 1.0 oC. 1000 cal = 1 kcal = 1 Calorie (a food “calorie”) But we use the unit called the JOULE 1 cal = 4.184 joules

13. HEAT TRANSFER The quantity of heat transferred depends on: • The quantity of material 2. The size of the temperature change 3. The identity of the material gaining or losing heat

14. HEAT CAPACITY Specific heat = The heat required to raise 1 g of substance 1 ˚K.

16. Heat Calculations Specific heat capacity (J/ g * K) Change in temperature (K) q = C * m * ∆T Mass of substance (g) Heat transfer (J)

19. heat gain/lose = q = (sp. ht.)(mass)(∆T) Specific Heat Capacity If 25.0 g of Al cool from 310 oC to 37 oC, how many joules of heat energy are lost by the Al?

20. Specific Heat Capacity where ∆T = Tfinal - Tinitial q = (0.897 J / g•K)(25.0 g)(37 - 310)K q = - 6120 J Notice that the negative sign on q signals heat “lost by” or transferred OUT of Al.

21. Change of State: Heat of Fusion constant T q = (heat of fusion)(mass) Ice + 333 J/g (heat of fusion) -----> Liquid water

22. Heating/Cooling Curve for Water Evaporate water Heat water Note that T is constant as ice melts Melt ice

23. +333 J/g +2260 J/g Heat & Changes of State What quantity of heat is required to melt 500. g of ice and heat the water to steam at 100 oC? Heat of fusion of ice = 333 J/g Specific heat of water = 4.2 J/g•K Heat of vaporization = 2260 J/g

24. Heat & Changes of State 1. melt ice 0 oC q = (500. g)(333 J/g) = 1.67 x 105 J 2. water from 0 oC to 100 oC q = (500. g)(4.2 J/g•K)(100 - 0)K = 2.1 x 105 J 3. To boil water 100 oC q = (500. g)(2260 J/g) = 1.13 x 106 J 4. Total q = 1.51 x 106 J = 1510 kJ

25. Calorimeter