1. March 26, 2014 Introduction to Thermochemistry But first, a little fun http://www.youtube.com/watch?v=nByTl6r- mlQ http://www.youtube.com/watch?v=nByTl6r- mlQ

2. Units of Energy The SI unit of energy is the joule (J). An older, non-SI unit is still in widespread use: The calorie (cal). 1 cal = 4.184 J 1 J = 1  kg m 2 s2s2

3. Work Energy used to move an object over some distance. w = F  d, where w is work, F is the force, and d is the distance over which the force is exerted. Work should not be considered a form of energy, but rather a mechanism for the transfer of internal energy

4. Heat Energy can also be transferred as heat. Heat flows from warmer objects to cooler objects. Heat is defined as random internal energy transfer between different bodies at different temperatures. Heat should not be considered a form of energy but rather a mechanism by which internal energy is transferred.

5. Transferal of Energy a)The potential energy of this ball of clay is increased when it is moved from the ground to the top of the wall.

6. Transferal of Energy a)The potential energy of this ball of clay is increased when it is moved from the ground to the top of the wall. b)As the ball falls, its potential energy is converted to kinetic energy.

7. Transferal of Energy a)The potential energy of this ball of clay is increased when it is moved from the ground to the top of the wall. b)As the ball falls, its potential energy is converted to kinetic energy. c)When it hits the ground, its kinetic energy falls to zero (since it is no longer moving); some of the energy does work on the ball, the rest is dissipated as heat.

8. First Law of Thermodynamics Energy is neither created nor destroyed. In other words, the total energy of the universe is a constant; if the system loses energy, it must be gained by the surroundings, and vice versa. Use Fig. 5.5

9. Internal Energy The internal energy of a system is the sum of all kinetic and potential energies of all components of the system; we call it E. Use Fig. 5.5

10. Internal Energy By definition, the change in internal energy,  E, is the final energy of the system minus the initial energy of the system:  E = E final − E initial Use Fig. 5.5

11. Changes in Internal Energy If  E > 0, E final > E initial – Therefore, the system absorbed energy from the surroundings. – This energy change is called endergonic.

12. Changes in Internal Energy If  E < 0, E final < E initial – Therefore, the system released energy to the surroundings. – This energy change is called exergonic or Exothermic

13. Changes in Internal Energy When energy is exchanged between the system and the surroundings, it is exchanged as either heat (q) or work (w). That is,  E = q + w.

14.  E, q, w, and Their Signs

15. Exchange of Heat between System and Surroundings When heat is absorbed by the system from the surroundings, the process is endothermic.

16. Exchange of Heat between System and Surroundings When heat is absorbed by the system from the surroundings, the process is endothermic. When heat is released by the system to the surroundings, the process is exothermic.

17. A system that absorbs heat from the surroundings is an ________ process. 1.Endothermic 2.Exothermic

18. Correct Answer: 1.Endothermic 2.Exothermic

19. The sign of enthalpy change,  H, associated with distillation of salt water is ________. 1.Positive 2.Negative 3.Zero

20. Correct Answer: As is evident in this figure of a distillation apparatus, a heat source is used; therefore the  H must be positive. 1.Positive 2.Negative 3.Zero

21. The sign of enthalpy change,  H, associated with a window washer dropping a squeegee from the top of a skyscraper is ________. 1.Positive 2.Negative 3.Zero

22. Correct Answer: The squeegee is falling, and potential energy is converted to kinetic energy. At the bottom, the squeegee will have lower potential energy, hence the enthalpy change is negative. 1.Positive 2.Negative 3.Zero

23. Hess’s law: if a reaction is carried out in a number of steps,  H for the overall reaction is the sum of  H for each individual step. For example: CH 4 (g) + 2O 2 (g)  CO 2 (g) + 2H 2 O(g)  H = -802 kJ 2H 2 O(g)  2H 2 O(l)  H= - 88 kJ CH 4 (g) + 2O 2 (g)  CO 2 (g) + 2H 2 O(l)  H = -890 kJ Hess’s Law

24. Another Example of Hess’s Law Given: C(s) + ½ O 2 (g)  CO(g)  H = -110.5 kJ CO 2 (g)  CO(g) + ½ O 2 (g)  H = 283.0 kJ Calculate  H for:C(s) + O 2 (g)  CO 2 (g)

25. If 1 mol of compound is formed from its constituent elements, then the enthalpy change for the reaction is called the enthalpy of formation,  H o f. Standard conditions (standard state): Most stable form of the substance at 1 atm and 25 o C (298 K). Standard enthalpy,  H o, is the enthalpy measured when everything is in its standard state. Standard enthalpy of formation: 1 mol of compound is formed from substances in their standard states. Enthalpies of Formation

26. If there is more than one state for a substance under standard conditions, the more stable one is used. Standard enthalpy of formation of the most stable form of an element is zero. Enthalpies of Formation

27. Substance  o f (kJ/mol) C(s, graphite)0 O(g)247.5 O 2 (g)0 N 2 (g)0

28. Using Enthalpies of Formation to Calculate Enthalpies of Reaction For a reaction Note: n & m are stoichiometric coefficients. Calculate heat of reaction for the combustion of propane gas giving carbon dioxide and water. Enthalpies of Formation C 3 H 8 (g) + 5O 2 (g)  3CO 2 (g) + 4H 2 O(  )

29. ∆H o f (kJ/mol):

30. Thermochemistry

31. Enthalpy The quantity of energy transferred as heat during a chemical reaction.  H = H products – H reactants Endothermic (requires energy) is positive Exothermic (released energy) is negative

32. Entropy How random are the particles in a reaction.  S The more random, the higher the value of  S (Positive) The less random, the lover the value of  S (Negative)

33. Gibbs Free Energy The difference between change in enthalpy and entropy  G  G =  H - T  S -  G is spontaneous in the forward direction +  G is not spontaneous in the forward direction.

34. Vocabulary to Note Thermochemistry Calorimetry Temperature Joule Heat Specific Heat Enthalpy (Change, reaction, Formation) Entropy Hess’s Law Free Energy Free Energy Change