I. THERMODYNAMICS (Read Introduction; pp. 803-804 and Section 19.1 and 804-808) A. History 1. Marcellin Berthelot (1827-1907) French chemist; dyes and explosives 1864 – proposed all spontaneous changes are exothermic (ice melting) 1897 - thermochemistry 2. Sadi Carnot (1796-1832) French scientist; founded thermodynamics 1824 – Reflexions sur la puissance motrice de feu (On the Motive Power of Fire) analysis of factors that would make idealized, frictionless steam engine B. Thermochemistry (E = q + w) and the First Law of Thermodynamics (energy is conserved) help us understand heat transfer and work between systems and surroundings, but it does not address the extent to which they occur. For that we need new concepts, namely: thermodynamics, spontaneity and entropy C. Thermodynamics - (Gk thérme- heat, dy’namis power) heat transfer The study of the laws that govern the energy and entropy changes of physical and chemical events. Thermodynamics studies the direction and extent of chemical reactions but not the speed of reaction D. Spontaneity 1. Spontaneous Change - change that occurs by itself without outside assistance; P more stable than R 2. Non-spontaneous Change - change that occurs with continued outside assistance. R more stable than P 3. Natural Direction Examples - rusting metal, burning paper, ice melting T > 0°C, water freezing T < 0°C E. Reversibility 1. Reversible Processes – system is changed and can be exactly restored to its original condition (hypothetical) 2. Irreversible Process - a system is changed and cannot be exactly restored to its original condition 3. NOTE: Any spontaneous process is irreversible; all “real” processes are irreversible F. Isothermal – at constant temperature II. ENTROPY (Read Section 19.2; pp. 808-811 and A Closer Look; p.810) A. Rudolf Clausis (1822-1888) German physicist, Zürich, Würtzburg, Bonn 1850 - enunciated the 2nd law, heat cannot pass from a colder body to a hotter body concluded that entropy in the Universe must increase 1857 Kinetic-Molecular Theory – theory of moving particles B. Entropy - (S) The thermodynamic quantity that describes the degree of randomness of a system. The greater the disorder or randomness in a system, the higher the entropy. C. Entropy is a state function depending only on the change and not the path. And so as a general equation… S = Sfinal - Sinitial and for a chemical system: S = Sproducts - Sreactants

D. The Second Law of Thermodynamics – For any spontaneous process, the entropy of the universe increases. Changes in the universe tend to move towards state of greater disorder. The second law in terms of equations: 1. Reversible Processes – Suniv = Ssystem + Ssurroundings = 0 2. Irreversible Process - Suniv = Ssystem + Ssurroundings > 0 E. Entropy Effects Associated with Melting and Freezing
Change Melting Temperature >mp =mp <mp >mp =mp <mp Sign of Ssys + + + Sign of Ssurr + + + Suniv= Ssys+Ssurr >0 =0 <0 >0 =0 <0 Spontaneity S EQ NS NS EQ S


(READ Section 19.3; pp. 811-819)
F. What does Entropy have to do with molecules? 1. Molecules undergo many specific types of motion including translational, vibrational and rotational motions. Since is would be nearly impossible to consider all the motions of every molecule in a substance, statistical thermodynamics was created to connect the microscopic descriptions of all of the individual molecules within a group to the macroscopic description of the group. The term that is used to describe individual molecules is microstates – a single possible arrangement of molecules at any one instant 2. Lugwig Boltzman (1844-1906) Austrian physicist, Vienna, Graz, Munich, Leipzig; kinetic theory of gases 1872 Boltzman Equation S = k ln W where k is the Boltzman constant, 1.38x10-23 J/K W # of microstates ( Gr Wahrscheinlichkeit – probability) 3. Entropy increases with the number of microstates of the system 4. Entropy increases with increasing: a. Temperature b. Volume c. Number of independently moving particles 5. Entropy increases for processes in which: a. gases form from solids/liquids b. liquids form from solids c. the number of gas molecules increases

S = moles of gasproducts – moles of gasreactants

6. The Third Law of Thermodynamics - the entropy of a pure crystalline at absolute zero is zero

(READ Section 19.4; pp. 820-822)
7. Standard Molar Entropies (S°) - As before, ° means at standard temperature (298°C) and pressure (1 atm) and read the standard molar entropy. S° values can be looked up on a table but in general: a. gases > liquids > solids b. as molecular mass increases, so does the S° value c. as the number of atoms increases, so does the S° value 8. The entropy change in a chemical reaction can be calculated using: S° = S°products - S°reactants 9. Calculating the entropy change in the surrounding depends on the heat absorbed by the system or: S°surroundings = -qsystem /T if isothermal then H°reaction = qsys

III. GIBBS FREE ENERGY (Read Section 19.5 pp. 822-827 and A Closer Look; p. 825) A. Josiah Willard Gibbs (1839-1903) American physicist/mathematician, Yale Gibbs Free Energy - thermodynamic quantity that relates enthalpy, entropy and temperature: G = H - TS “Free” refers to the maximum energy produced in a physical or chemical change not lost to heat that can be theoretically harnessed and therefore available to do useful work or mathematically G = H - TS (kJ)

B. The Gibbs free energy is an indicator of the spontaneity of a reaction or physical change. C. Gibbs free energy is a state function and can be calculated by comparing initial and final states using G = Gfinal - Ginitial and for a chemical system G = Gproducts – Greactants

D. Effects of the sign of G At constant temperature/pressure, change is spontaneous if accompanied by a decrease in the free energy of the system. For a change to be spontaneous, Gfinal must be less than Ginitial and G must be negative. A summary of the relationship of G and spontaneity at constant temperature and pressure 1. if G is negative, the reaction is spontaneous 2. if G is zero, the reaction is at equilibrium 3. if G is positive, the reaction is non-spontaneous; work must be supplied from surroundings to occur

(Read Section 19.6; pp. 827-829)
E. Effects of H, S, and T on Spontaneity 1. When enthalpy is exothermic and entropy is increasing, then the change is spontaneous H is (-) and S is (+), G will be (-) and is spontaneous at all temperatures. Example: collapse of building 2. When enthalpy is endothermic and entropy is decreasing, then the change is non-spontaneous H is (+) and S is (-), G will be (+) and is non-spontaneous at all temperatures. Example: formation of building from a pile of bricks 3. When enthalpy is exothermic and entropy is decreasing, then the change is spontaneous H is (-) and S is (-), G will be (-) and is spontaneous only at low temperatures(below 0°C). Example: freezing of water 4. When enthalpy is endothermic and entropy is increasing, then the change is spontaneous H is (+) and S is (+), G will be (-) and is spontaneous only at high temperatures(above 0°C). Example: melting of ice

(READ Section 19.7; pp. 830-835)
F. Free Energy and Equilibrium 1. When G is equal to 0, the system is considered in a state of equilibrium G = Gfinal - Ginitial or G = Gproducts - Greactants or G=0

2. Using the example of freezing water, H2O(l)  H2O(s) , we know the following: a. when the temperature is below 0°C, the change is spontaneous and b. when the temperature is above 0°C, the change is non-spontaneous. c. when the temperature is exactly 0°C, then G=0, neither spontaneous or non-spontaneous. The ice-water mixture is at equilibrium and as long as no heat is added or taken away, the equilibrium will exist indefinitely. 3. During a phase change, the free energy equation can be used to calculate the phase change temperature. G° = H° - TbS° = 0 G. Equilibrium Constant G = G° + RTln Q or Tb = H°/S°

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