2. Lecture 1: Energy and Enthalpy Reading: Zumdahl 9.1 and 9.2 Outline –Energy: Kinetic and Potential –System vs. Surroundings –Heat, Work, and Energy –Enthalpy

3. Energy is the capacity to do work or to produce heat Energy is conserved, it can neither be created nor destroyed, different forms of energy interconvert However, the capacity to utilize energy to do work is limited (entropy)

4. Energy: Kinetic vs. Potential Potential Energy (PE) –Energy due to position or chemical composition –Equals (mgh) in example. Kinetic Energy (KE) –Energy due to motion. –Equals mv 2 /2 in example.

5. Mechanical Energy = KE + PE Energy is the sum of kinetic energy and potential energy. Energy is readily interconverted between these two forms. If the system of interest is isolated (no exchange with surroundings), then total energy is constant.

6. Example: Mass on a Spring Initial PE = 1/2 kx 2 At x = 0: –PE = 0 –KE = 1/2mv 2 =1/2kx 2 Units of Energy Joule = kg.m 2 /s 2 Example: –Init. PE = 10 J –M = 10 kg –Vmax = [2(PE)/M] 1/2 = 1.4m/s 0

7. Energy: Kinetic vs. Potential Potential Energy (PE) –Energy due to position or chemical composition –Equals (mgh) in example. Kinetic Energy (KE) –Energy due to motion. –Equals mv 2 /2 in example.

8. First Law of Thermodynamics First Law: Energy of the Universe is Constant E = q + w q = heat. Transferred between two bodies of differing temperature. Note: q ≠ Temp! w = work. Force acting over a distance (F x d)

9. Applying the First Law Need to differentiate between the system and surroundings. System: That part of the universe you are interested in (i.e., you define it). Surroundings: The rest of the universe.

10. Conservation of Energy Total energy is conserved. Energy gained by the system must be lost by the surroundings. Energy exchange can be in the form of q, w, or both.

11. Heat Exchange: Exothermic Exothermic Reaction. Chemical process in which system evolves resulting in heat transfer to the surroundings Heat flows out of the system q < 0 (heat is lost)

12. Another Example of Exothermic

13. Heat Exchange: Endothermic Endothermic Reaction: Chemical process in which system evolves resulting in heat transfer to the system Heat flows to the system q > 0 (heat is gained)

14. Another Example of Endothermic

15. In exothermic reactions, the potential energy stored in chemical bonds is converted into thermal energy (random kinetic energy), i.e. heat Once we have done that, we have lost the ability to utilize the same potential energy to do work or generate heat again (dissipation)

16. Energy and Sign Convention If system loses energy: E final < E initial E final -E initial =  E < 0. If system gains energy: E final > E initial E final -E initial =  E > 0.

17. Heat and Work Sign Convention If system gives heat q < 0 (q is negative) If system gets heat q > 0 (q is positive) If system does work w < 0 (w is negative) If work done on system w > 0 (w is positive)

18. Example: Piston Figure 9.4, expansion against a constant external pressure No heat exchange: q = 0 System does work: w < 0 (adiabatic)

19. Example (cont.) How much work does the system do? P ext = force/area |w| = force x distance = P ext x A x  h = P ext  V w = - P ext  V (note sign)

20. When it is compressed, work is done to a gas When it is expanded, work is done by the gas (e.g. your car’s engine)

21. Example 9.1 A balloon is inflated from 4 x 10 6 l to 4.5 x 10 6 l by the addition of 1.3 x 10 8 J of heat. If the balloon expands against an external pressure of 1 atm, what is  E for this process? Ans: First, define the system: the balloon.

22. Example 9.1 (cont.)  E = q + w = (1.3 x 10 8 J) + (-P  V) = (1.3 x 10 8 J) + (-1 atm (V final  V init )) = (1.3 x 10 8 J) + (-0.5 x 10 6 l.atm) Conversion: 101.3 J per l x atm (-0.5 x 10 6 l.atm) x (101.3 J/l.atm) = -5.1 x 10 7 J

23. Example 9.1 (cont.)  E = (1.3 x 10 8 J) + (-5.1 x 10 7 J) = 8 x 10 7 J (Ans.) The system gained more energy through heat than it lost doing work. Therefore, the overall energy of the system has increased.

24. Definition of Enthalpy Thermodynamic Definition of Enthalpy (H): H = E + PV E = energy of the system P = pressure of the system V = volume of the system

25. Why we need Enthalpy? Consider a process carried out at constant pressure. If work is of the form  PV), then:  E = q p + w = q p - P  V  E + P  V = q p q p is heat transferred at constant pressure.

26. Definition of Enthalpy (cont.) Recall: H = E + PV  H =  E +  PV) =  E + P  V (P is constant) = q p Or  H = q p The change in enthalpy is equal to the heat transferred at constant pressure.

27. Changes in Enthalpy Consider the following expression for a chemical process:  H = H products - H reactants If  H >0, then q p >0. The reaction is endothermic If  H <0, then q p <0. The reaction is exothermic

28. Enthalpy Changes Pictorially Similar to previous discussion for Energy. Heat comes out of system, enthalpy decreases (ex. Cooling water). Heat goes in, enthalpy increases (ex. Heating water)