CHEM303 final exam review PDF

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Dr Geddes. thermodynamics, chemical potential, ionic solutions, equilibrium...

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Lecture 1: Thermodynamics Overview Random error Pressure-volume work w = - [integral] PdV W = - P[delta]V if pressure is constant Mechanical work Work done on a system is positive The 0th Law of Thermodynamics Two systems in thermal equilibrium with a third are in equilibrium with each other Heat q = pt t = time The change in temperature it produces upon flowing into a definite amount of sustance Power p = EI E = potential difference I = electric current Homogenous system Heterogenous system Intensive variable INDEPENDENT of size Extensive variable

Lecture 2: Chemical Equations, The First Law of Thermodynamics (Conservation of Energy) A state function is path-independent Isochoric Adiabatic Internal Energy [delta]U = q + w The 1st Law of Thermodynamics Conservation of energy [delta]U = q for a process at constant volume [delta]H = q at constant pressure Enthalply H = U + PV Heat Capacity C = U/T = q/T at constant volume C = H/T = q/T constant pressure

Lecture 3: The Second Law of Thermodynamics (Entropy) Hess’s Law standard enthalpies of individuals reactions can be combined The Second Law of Thermodynamics total entropy of any system cannot decrease other than by increasing the entropy of another system dS = q / T Order and Disorder Disorder = Disorder Capacity / Information Capacity Order = 1 - Order Capacity / Information Capacity

Lecture 4: Kirchoff’s Law, The Third Law of Thermodynamics Relationship between heat capacities C(p) - C(v) = nR Kirchoff’s Law change of standard reaction enthalpy, based on enthalpy, heat capacity and temperature The Third Law of Thermodynamics the entropy of any pure substance in a perfect crystalline state is 0 at T=0

Lecture 5: Thermodynamic Function, Criteria for Spontaneous Changes and Chemical Equilibrium Combined First and Second Laws dU = TdS - PdV dH = TdS +VdP Helmholtz free energy A = U - TS Useful work, constant T and V Gibbs free energy G = H - TS dA = -PdV - SdT dG = VdP - SdT If Gibbs free energy is less than 0, spontaneous Useful work, constant T and V Entropy of Mixing when two substances mix, available volume changes, increasing uncertainty and entropy Gibbs paradox slightest difference between substances yields big entropy change

Lecture 6: Chemical Potential Chemical potential think about electrical potential but the chemical version! Partial molar quantity V = [delta]V / [delta]n Extent of reaction [delta]squiggle = [derivative]n/v

vi = stoichiometric numbers (positive for products, negative for reactants) n = amount of reactant Ri = reactant Activity a = exp[(mu - md,std) - (RT)] mu = mu,std + RTln(a) mu,std = standard chemical potential [delta]G = [sigma]v(mu,std) Reaction Quotient the relative proportion of products and reactants present in the reaction mixture at some instant time (use current concentrations rather than equilibrium concentrations) Q = [pi](a)^v = (P,c^c)(P,d^d) / (P,a^a)(P,b^b) Le Chatelier’s Principle if system at equilibrium experiences a change, the equilibrium will shift to counteract that change Q < K reaction moves right (actual smaller than equilibrium, must move to products) Activity Coefficient a factor used to account for deviations from ideal K = [pi](a)^v = [pi](c)^v * [pi](gamma)^v [gamma] = a(P,std) / (P,i) P,i = partial pressure [gamma] = a(c,std) / (c,i) [gamma] = a(m,std) / (m,i) [gamma] = a/x,i Activity coefficients have different values on different scales Gases approach ideal as pressure gets very low Fugacity the tendency of a substance to prefer one phase over another A substance has a different fugacity for each phase The lowest will be the most favorable The unit is the same as pressure K as a Function of T ***LnK = -H/RT + S/R*** Differentiate to get the van’t Hoff equation

Lecture 7: Ionic Solutions, Electrochemistry, pH Debye-Huckel constant Am ln[gamma] = -A(m,i)(z^2)(I^0.5) for low ionic strengths (< 0.01 mol/kg) ln[gamma] = -A(m,i)(z^2)(I^0.5) / (1 + BI^0.5) I = ½[sigma]mz^2 = ionic strength z = charge on ion m,i = molality = n/(M,s*n,s) n= mol M,s = molar mass n,s = amount of solvent

Mean ionic activity coefficients a way to “average” activity via multiplication and powers rather than addition and multiplication Nernst Equation equilibrium reduction potential E = E,std - (RT/[vF])lnQ v = number of electrons Q = reaction quotient Standard Half Cell Potentials E,std = E,std,cathode - E,std,anode The Henderson-Hasselbach Equation

Lecture 8: Equilibrium Identical and Independent Binding Cooperative Binding the first ligand to bind makes it easier for the next Anti-cooperative binding Excluded-side binding binding at one site completely excludes binding at another side...