1. Chemical Thermodynamics Lecture 1. Chemical Thermodynamics Prepared by PhD Halina Falfushynska

2. Chemical Thermodynamics

3. Chemical Thermodynamics Thermodynamic Systems - Definitions Isolated System: No matter or energy cross system boundaries. No work can be done on the system. Open System: Free exchange across system boundaries. Closed System: Energy can be exchanged but matter cannot. Adiabatic System: Special case where no heat can be exchanged but work can be done on the system (e.g. PV work).

4. Chemical Thermodynamics Thermodynamic State Properties Extensive: These variables or properties depend on the amount of material present (e.g. mass or volume). Intensive: These variables or properties DO NOT depend on the amount of material (e.g. density, pressure, and temperature).

5. Chemical Thermodynamics Units of Energy SI Unit for energy is the joule, J: sometimes the calorie is used instead of the joule: 1 cal = 4.184 J (exactly) A nutritional Calorie: 1 Cal = 1000 cal = 1 kcal establish the equation

6. Chemical Thermodynamics

7. Chemical Thermodynamics Measuring Heat Intuitively, we know temp. is related to heat, and different substances require different amounts of heat to change their temp.  q = C × ΔT heat capacity (C): the amount of heat needed to change the temp of a substance by 1 K  What’s wrong with this?

8. Chemical Thermodynamics Specific Heat Capacity specific heat capacity (C s ): the amount of heat required to change the temp. of 1 g of a substance by 1 K

9. Chemical Thermodynamics Note the tremendous difference in Specific Heat. Water’s value is VERY HIGH.

10. Chemical Thermodynamics Specific Heat Specific Heat C : The amount of energy that raises the temperature of 1 Kg of A substance by 1 K ( or 1C) is called specific heat. In other words, the thermal energy of the system changes by ΔE th = Mc ΔT ( Kelvin and Celsius temperature scales have the same step size. ) The first law of thermodynamics ΔE th = W + Q, In working with solids And liquids, we almost always change the temperature by heating, so W = 0, Then the heat needed to bring about a temperature change ΔT is Q = Mc ΔT. Molar specific heat: The amount of energy that raises the temperature of 1 mol of a substance by 1K. It depends on the molar mass.

11. Chemical Thermodynamics First Law of Thermodynamics Energy is Conserved E = constant (E Kinetic + E Potential + E Internal + E Chemical + E Mechanical + E Electrical ) = constant The first law of thermodynamics is an expression of the principle of conservation of energy. The law states that energy can be transformed, i.e. changed from one form to another, but cannot be created nor destroyed.conservation of energy

12. Chemical Thermodynamics First Law of Thermodynamics Energy cannot be created nor destroyed. Energy can, however, be converted from one form to another or transferred from a system to the surroundings or vice versa.

13. Chemical Thermodynamics

14. Chemical Thermodynamics Isochoric Process Changes in  Heat Added or Removed  Temperature  Pressure (courtesy F. Remer) V= const, A=0 Q v = U 2 -U 1 =  U Changes in  Heat Added or Removed  Temperature  Volume p= const Q p =  U + p  V =  H Isobaric Process Isothermal Process T= const Changes in: Heat Added or Removed  Pressure; Volume

15. Chemical Thermodynamics Reactions at Constant V

16. Chemical Thermodynamics Constant P Calorimetry In calorimetry, we run a controlled reaction and monitor the temp. change. The important things to remember are: q system = -q surroundings q rxn = -q soln q rxn = ΔH rxn

17. Chemical Thermodynamics Work in Ideal-Gas Processes The work done on the system When we press the gas, the gas volume becomes smaller, so the total work done by the environment on the gas

18. Chemical Thermodynamics Work in some special processes Isochoric Process W = 0 Isobaric Process W = -PΔV Isothermal Process Work depends on path isochore isobare isotherme

19. Chemical Thermodynamics Finding work from the P-V diagram W = the negative of the area under the PV curve between Vi and Vf W < 0 W > 0

20. Chemical Thermodynamics Exothermic and Endothermic Processes Endothermic: absorbs heat from the surroundings. An endothermic reaction feels cold,. H > 0 Exothermic: transfers heat to the surroundings. An exothermic reaction feels hot, H < 0 (combustion). The First Law of Thermodynamics

21. Chemical Thermodynamics

22. Chemical Thermodynamics Spontaneous Processes Spontaneous processes are those that can proceed without any outside intervention. The gas in vessel B will spontaneously effuse into vessel A, but once the gas is in both vessels, it will not spontaneously

23. Chemical Thermodynamics Spontaneous Processes Processes that are spontaneous at one temperature may be nonspontaneous at other temperatures. Above 0  C it is spontaneous for ice to melt. Below 0  C the reverse process is spontaneous.

24. Chemical Thermodynamics Reversible Processes In a reversible process the system changes in such a way that the system and surroundings can be put back in their original states by exactly reversing the process. Changes are infinitesimally small in a reversible process.

25. Chemical Thermodynamics Irreversible Processes Irreversible processes cannot be undone by exactly reversing the change to the system. All Spontaneous processes are irreversible. All Real processes are irreversible.

26. Chemical Thermodynamics Gibbs Free Energy 1.If  G is negative, the forward reaction is spontaneous. 2.If  G is 0, the system is at equilibrium. 3.If  G is positive, the reaction is spontaneous in the reverse direction.

27. Chemical Thermodynamics Standard Free Energy Changes Standard free energies of formation,  G f  are analogous to standard enthalpies of formation,  H f .  G  can be looked up in tables, or calculated from S° and  H .

28. Chemical Thermodynamics Free Energy Changes Very key equation: This equation shows how  G  changes with temperature. (We assume S° &  H° are independent of T.)

29. Chemical Thermodynamics Free Energy and Equilibrium Under non-standard conditions, we need to use  G instead of  G°. Q is the reaction quotiant from chapter 15. Note: at equilibrium:  G = 0. away from equil, sign of  G tells which way rxn goes spontaneously.

30. Chemical Thermodynamics Gibbs Free Energy 1.If  G is negative, the forward reaction is spontaneous. 2.If  G is 0, the system is at equilibrium. 3.If  G is positive, the reaction is spontaneous in the reverse direction.

31. Chemical Thermodynamics Change in Gibbs Free Energy

32. Chemical Thermodynamics 32 The work of the expanded ideal gass h s F WORK = Force and Distance