Chemistry 231

1. Chemistry 231 Work, Heat and Internal Energy: The First Law

2. Systems and Surroundings System – the specificpart of the universe of interest to us Surroundings – the part of the universe not contained in the system

3. State of a System (Continued) • 3 types of Systems • open system – exchanges mass and energy • closed system – exchanges energy but no mass • isolated system – no exchange of either mass or energy

4. Types of Systems (Continued) insulation cork Open system Closed System Isolated System

5. Some Definitions • State of a system • the system is in a definite state when each of its properties has a definite value. • Change in state • initial state • final state • Path • initial and final states • intermediate states

6. Some Definitions (Continued) • Process • reversible or irreversible transformation • Cyclic transformation • begins and ends at the same state variables.

7. Some Definitions (Continued) • Isothermal • dT = 0 • Isochoric • dV = 0 • Isobaric • dP = 0

8. Work Unit of work = J = 1 kg m/s2 • Work (w) • any quantity that flows across the system’s boundary and is completely convertible into the lifting of a mass in the surroundings. • How much work was done?

9. A Force Vs. Displacement Curve

10. Work Done in a Closed,Fluid System mass (m) h2 Direction of piston mass (m) Piston (T, P2, V2) h1 Piston (T, P1, V1) State 2 State 1 A single-stage expansion process

11. System and Surroundings • The work done in the surroundings • wsurr= PextDV • The work done by the system • wsys = - wsurr = - PextDV • For an infinitesimal volume change • dwsys = - Pext dV

12. Reversible (Multistage) Expansion • If the system is in equilibrium • Fsys = -Fext • P = Pext • For a simple system • d wrev = - P dV

13. Reversible Transformation in an Ideal Gaseous System Ideal gas as the working fluid.

14. Reversible Transformation (Continued) For an isothermal process (ideal gas as working fluid)

15. Irreversible Transformations dwirr = -Pext dV for a constant external pressure

16. Heat • Heat - the quantity that flows across the boundary of the systemduring a change in state • due to temperature difference between system and surroundings • HOT to COLD (never the other way around)!!!

17. 'Amount of Heat' C - the heat capacity of the system. Measured by determining the temperature change of some known object

18. Macroscopic Heat Flows Integrate the infinitesimal heat flow

19. Heat Flows heat system surroundings Exothermic - system to surroundings Endothermic – surroundings to system

20. Latent Heats • Heat flows during phase changes - latent heats • Latent heat of vapourisation • Latent heat of fusion

21. The Internal Energy Subject our system to a cyclic transformation

22. Cyclic Integrals of Exact Differentials The following would be true for an exact differential

23. The Internal Energy • For a general process The infinitesimal change in the internal energy

24. The Properties of U In general, we write U as a function of T and V

25. Isochoric Changes in U Examine the first partial derivative

26. The Constant Volume Heat Capacity Define the constant volume heat capacity, CV

27. Heat Flows Under Constant Volume Conditions • For a macroscopic system For a system undergoing an isochoric temperature change

28. Isothermal Changes in U Examine the second partial derivative

29. The Joule Experiment O O O O C C F F 50 120 100 40 30 80 20 60 10 40 0 20 10 0 20 20 30 40 40 60 50 A T1, Vm,1, P1 B Stirrer Thermal insulation Valve

30. The Joule Coefficient is known as the Joule coefficient, J. The partial derivative

31. Internal Energy and the Joule Coefficient The change in the internal energy under isothermal conditions is related to the Joule Coefficient

32. Adiabatic Processes For an adiabatic process, q = 0!! The first law becomes

33. Adiabatic Processes for an Ideal Gas For an ideal gas undergoing a reversible, adiabatic process

34. State Changes Under Constant Pressure Conditions mass (m) Piston (T, P, V) Defining the enthalpy of the system Re-examine the piston with the weight on top

35. A Constant Pressure Process • Integrating The first law

36. Enthalpy Define the enthalpy of the system, H

37. The Properties of H In general, we write H as a function of T and P

38. Isobaric Changes in H Examine the first partial derivative

39. The Constant Pressure Heat Capacity Define the constant pressure heat capacity, CP

40. Heat Flows Under Constant Pressure Conditions • For a macroscopic system For a system undergoing an isobaric temperature change

41. Relating CP and CV • In general For an ideal gas

42. Isothermal Changes in H Examine the second partial derivative

43. The Joule-Thomson Experiment Thermal insulation O O O O C O C O O O F F C C F F 50 120 50 120 100 40 100 40 30 80 30 80 20 20 60 60 10 10 40 40 0 0 20 20 10 10 0 0 20 20 20 20 30 30 40 40 40 40 60 50 60 50 T2, P2, Vm,2 T1, P1, Vm,1 Porous Plug

44. The Joule-Thomson Coefficient is known as the Joule-Thomson coefficient, JT. The partial derivative

45. Relating H to the Joule-Thompson Coefficient The change in the enthalpy under constant pressure conditions is related to the Joule-Thomson Coefficient

46. Enthalpy Changes for Reactions The shorthand form for a chemical reaction J = chemical formula for substance J J = stoichiometric coefficient for J

47. Reaction Enthalpy Changes The enthalpy change for a chemical reaction Hm [J] = molar enthalpies of substance J nJ = number of moles of J in the reaction

48. The Enthalpy Change Reaction beginning and ending with equilibrium or metastable states Note – Initial and final states have the same temperature and pressure!

49. Reaction Enthalpies (cont’d) We note that 1 mole of a reaction occurs if

50. A Standard State Reaction A reaction that begins and ends with all substances in their standard states The degree sign, either  or  P = 1.00 bar [aqueous species] = 1.00 mol/ kg T = temperature of interest (in data tables - 25C or 298 K).