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Lecture 1: Energy and Enthalpy

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. Energy is the capacity to do work or to produce heat

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Lecture 1: Energy and Enthalpy

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  1. 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

  2. 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)

  3. 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 mv2/2 in example.

  4. 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.

  5. Example: Mass on a Spring • Initial PE = 1/2 kx2 • At x = 0: • PE = 0 • KE = 1/2mv2=1/2kx2 • Units of Energy Joule = kg.m2/s2 • Example: • Init. PE = 10 J • M = 10 kg • Vmax = [2(PE)/M]1/2 = 1.4m/s 0

  6. 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 mv2/2 in example.

  7. 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) First Law of Thermodynamics

  8. 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.

  9. 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.

  10. 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)

  11. Another Example of Exothermic

  12. 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)

  13. Another Example of Endothermic

  14. 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)

  15. Energy and Sign Convention • If system loses energy: Efinal < Einitial Efinal-Einitial = DE < 0. • If system gains energy: Efinal > Einitial Efinal-Einitial = DE > 0.

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

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

  18. Example (cont.) • How much work does the system do? • Pext = force/area • |w| = force x distance = Pext x A x Dh = PextDV • w = - PextDV (note sign)

  19. 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)

  20. Example 9.1 • A balloon is inflated from 4 x 106 l to 4.5 x 106 l by the addition of 1.3 x 108 J of heat. If the balloon expands against an external pressure of 1 atm, what is DE for this process? • Ans: First, define the system: the balloon.

  21. Example 9.1 (cont.) DE = q + w = (1.3 x 108 J) + (-PDV) = (1.3 x 108 J) + (-1 atm (Vfinal - Vinit)) = (1.3 x 108 J) + (-0.5 x 106 l.atm) • Conversion: 101.3 J per l x atm (-0.5 x 106 l.atm) x (101.3 J/l.atm) = -5.1 x 107 J

  22. Example 9.1 (cont.) DE = (1.3 x 108 J) + (-5.1 x 107 J) = 8 x 107 J (Ans.) The system gained more energy through heat than it lost doing work. Therefore, the overall energy of the system has increased.

  23. 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

  24. Why we need Enthalpy? • Consider a process carried out at constant pressure. • If work is of the form D(PV), then: DE = qp + w = qp - PDV DE + PDV = qp qp is heat transferred at constant pressure.

  25. Definition of Enthalpy (cont.) • Recall: H = E + PV DH = DE + D(PV) = DE + PDV (P is constant) = qp • Or DH = qp • The change in enthalpy is equal to the heat transferred at constant pressure.

  26. Changes in Enthalpy • Consider the following expression for a chemical process: DH = Hproducts - Hreactants If DH >0, then qp >0. The reaction is endothermic If DH <0, then qp <0. The reaction is exothermic

  27. 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)

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