Presentation Transcript

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, no exchange of energy or matter
5. Steady state system Flux in = flux out There can be exchange, but no change in total abundance
6. Parts of 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. 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
9. State function a property of a system which has a specific value for each state (e.g., condition) E.g., 1 g water @ 25º C A couple of state functions for this sytem are amount of mass (1 g) and T (25º C) There are others we will learn about Path independent E.g., state would be the same if you condensed steam or melted ice For the values of the state functions, it doesn’t matter how the state got there
10. 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)
11. Zeroth law If two systems are in thermal equilibrium No heat is exchanged between the systems They have the same “temperature” T is the newly defined state function How is temperature defined?
12. 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
13. Fig. Levine V (L) T (ºC) 1 mole of N2 at constant P 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
14. 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
15. Three types of energy Kinetic and potential – physically defined Internal – chemically defined
16. 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
17. Potential + Kinetic energy + internal energy Minimum or rest energy Here only internal energy, U
18. New state function – Enthalpy (H) PV = pressure * volume = work done on/by the system Units – energy, e.g. j, kJ, cal etc. Extensive 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 Extensive = units of energy/T, e.g. kJ/K 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
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 (kcal) S units = kJ/K (kcal/K) T is Kelvin scale (K) G = H - TS Equilibrium A, B, C, and D present
23. A + B ↔ C + D Equilibrium A, B, C, and D present
24. 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
25. 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
26. Non-equilibrium system: Equilibrium system A + B → C + D DG ≠ 0 A + B ← C + D DG ≠ 0 A + B ↔ C + D DG = 0
27. 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
28. Math Review (on board)
29. 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
30. 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
31. Equilibrium, activities, chemical potentials (on board)