Thermodynamics

1. Thermodynamics • Thermodynamics • Way to calculate if a reaction will occur • Kinetics • Way to determine the rate of reactions • Thermodynamic equilibrium rarely attained: • Biological processes – work against thermo • Kinetic inhibitions

2. Thermodynamics very useful • Good approximation of reactions • Tells direction a reaction should go • Basis for estimated rates • Farther from equilibrium, faster rate

3. Thermodynamic definitions • System – part of universe selected for study • Surroundings (Environment) – everything outside the system • Universe – system plus surroundings • Boundary – separates system and surroundings • Real or imagined • Boundary conditions – solutions to Diff Eq.

4. Types of systems • Open system • Exchanges with surroundings • Mass, also heat and work • Closed system • no exchange of matter between with surrounding and system, energy can be exchanged • Isolated system • there is no interaction with surroundings, either energy or matter possible

5. Steady state system • Flux in = flux out • There can be exchange, but no change in total abundance

6. Within Systems • Phase – physically and chemically homogeneous region • Example: saturated solution of NaCl • Species – chemical entity (ion, molecule, solid phase, etc.) • E.g. NaCl (solid) + H20 (liquid) • Also Na+, Cl-, OH-, H+, NaClo, others

7. Components • Minimum number of chemical entities required to define compositions of all species • Many different possibilities • Na+, Cl-, H+, OH- • NaCl – H2O

8. Characteristics of components: • Every species can be written as a product of reactions involving only the components • No component can be written as a product of a reaction involving only the other components

9. Thermodynamic Properties • Extensive • Depends on amount of material • E.g., moles, mass, energy, heat, entropy • Additive • Intensive • Don’t depend on amount of material • Concentrations, density, T, heat capacity • Can’t be added

10. State function • a property of a system which has a specific value for each state (e.g., condition) • E.g., 1 g water @ 25 C • Variables are amount of mass (1 g) and T (25 C) • Path independent • E.g., state would be the same if you condensed steam or melted ice

11. Thermodynamic Laws • Three laws – each derives a “new” state function • 0th law: yields temperature (T) • 1st law: yields enthalpy (H) • 2nd law: yields entropy (S)

12. Zeroth law • If two systems are in thermal equilibrium • No heat is exchanged between the systems • They have the same temperature

13. Measurement of T • Centigrade • 100 divisions between melting and boiling point of water • Kelvin - Based on Charles law • At constant P and m, there is a linear relationship between volume of gas and T • Size of unit is same as centigrade V = a1 + a2J Where V = volume J = temperature a1 & a2 = constants

14. Fig. Levine V (L) T (ºC) Experimental results - extrapolation of results show intercept T @ V = 0 is about -273ºC - Kelvin scale based on triple point of water - defined as being 273.16 K

15. First law • Change in the internal energy of a system is the sum of the heat added (q) and amount of work done (w) on system • Energy conserved

16. Three types of energy • Kinetic and potential – physically defined • Internal – chemically defined

17. Internal energy (U) • Molecular rotation, translation, vibration and electrical energy • Potential energy of interactions of molecules • Relativistic rest-mass energy • In thermo, a system at rest • Kinetic and potential energy = 0 • Thermodynamics considers only changes in internal energy

18. New state function – Enthalpy • PV = work done on/by the system H = U + PV

19. Second Law • A system cannot undergo a cyclic process that extracts heat from a heat reservoir and also performs an equivalent amount of work on the surroundings • i.e., it is impossible to build a machine that converts heat to work with 100% efficiency

20. New state function • Entropy = S • Entropy is variable in definition of Gibbs free energy (G) • G used to determine equilibrium of reactions

21. Equilibrium Thermodynamics • Equilibrium occurs with a minimum of energy in system • Systems not in equilibrium move toward equilibrium through loss of energy Potential + Kinetic energy Minimum or rest energy

22. If system is at constant T and P, measure of energy of system is given by G • G = f(H,S, T) • G and H units = kJ/mol (kcal/mol) • S units = kJ/mol.K (kcal/mol.K) • T is Kelvin scale (K) G = H - TS Equilibrium A, B, C, and D present

23. Consider processes in system at constant T & P • Means system changes • May be chemical reaction • Here D is change in state: DG =DH - TDS D = State2 – State1

24. When system moves toward equilibrium: • may release heat, e.g. DH < 0 • entropy may increase, e.g. DS > 0 • Both may happen • Thus: • DG < 0 for spontaneous reaction • G2 < G1; DG = G2 – G1 < 0 • DG = 0 for process at equilibrium

25. G is an extensive state variable • It depends on the amount of material • The amount of G in a system is divided among components • Need to know how G changes for each component • First look at what variables control G • What is G a function of? • Want to know how G changes if all (or any) other variable change • Change = calculus

26. Math Review (on board)

27. If system is in thermal and mechanical equilibrium: • G = f(P, T, n1, n2, n3…) • Then total differential: (on board) • Infinitesimal change in G caused by infinitesimal change in P, T, n1, n2, n3… • These are values we need to know to know DG

28. Last term defined by Gibbs as chemical potential (m) (on board) • m is the amount that G changes (per mole) with addition of new component • Intensive property (G extensive) • Doesn’t depend on mass of system • For one component system m = G/n • For system at equilibrium, m of all components are identical

29. Equilibrium, activities, chemical potentials (on board)