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Lecture

Spontaneous Change (Ch 19).docx

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Department
Chemistry
Course
CHEM 1A03
Professor
David Brock
Semester
Fall

Description
Chapter 19: Spontaneous Change Spontaneous Processes Rusting 4 Fe(s) + 3O2(g)  2Fe2O3(s)  forward process is spontaneous; therefore, the reverse is non-spontaneous  Fe2O3 (s) will not decompose to Fe(s) and O2(g) - The reverse process is non-spontaneous - Making Fe(s) from Fe2O3 (s) reqires carbon - The oxygen is removed as Co2(g) Melting : H2O (s)  H2O (l)  spontaneous above 0C Spontaneity & ΔH - Although many exothermic reactions are spontaneous, there is NOT a direct correlation between ΔH and spontaneity - Thermite - Fe2O3 (s) + 2Al(s) + Al2O3(s) - Cold packs - NH4NO3(s)  NH4+ +NO3- - Reactions that lower their overall energy / give off heat are usually spontaneous Entropy (s): The Boltzmann Equation - S = klnW - S = entropy, k = Boltzmann constant, W = # of microstates - The greater the number of configurations ( or microstates) consistent with the “macrostate”, the greater the entropy of the system Entropy Change (ΔS) - Gas expansion or mixing causes an increase in entropy: - A(g) + B(g)  mixture of A & B - ΔS is a state function and an extensive property - Entropy change between two systems/states becomes less at higher temperatures Entropy Changes for Phase Changes - Solid  Liquid, melting: Sliquid > Ssolid - Liquid  Gas, vaporization: S vapour>Sliquid - ΔS = ΔHtr/Ttr (tr = transition) - Absolute molar entropy increases with increasing temperature - ΔS = qrev/T - Amplifying # of states that molecule can exist in - So we can increase entropy during phase change without changing temp Entropy & Molecular Structure - Molar entropy increase - With increasing moleculer complexity ie. With greater # of vibrational and rotational degrees of freedom available to the molecule (recall heat capacity) - For molecules of the same complexity, the more massive the molecule the greater the entropy eg. S(CH3Br) > S(CH3Cl) - Methane bonds can vibrate in and out towards each other – internal degrees of freedom - Can store a lot more entropy if it has more C-H bonds?? - More degrees of freedom  Higher complexity  Higher heat capacity  higher entropy - The larger an atom is, the more microstates associated with it - Notice S° is in J mol -1K-1 not kJ Air Bags: Entropy Increase - Chemistry of sodium azide: NaN3 - Impact detector detonates – timescale = 50 ms - NaN3  Na(s) + 3/2 N2(g) - Detonation of 130g Sodium Azide produces 73L N2 gas - Air bag contents (NaN3 + KNO3 + SiO2) - Metal oxides are still very reactive  final step involves silicates Which of the following processes leads to the greatest increase in entropy? b) adding 1 mole of MgCl2 to 1L of water  entropy = disorder, look at which will give up the most ions, similar to strong electrolytes, releasing more ions into solution makes a larger/greater disorder, more ions in a random environment  higher disorder  higher entropy Calculating ΔS° - Third law of thermodynamics: the entropy of a pure, perfect crystal at 0K is zero, therefore we have absolute entropes. - But everything has entropy so we have no perfect crystals. - ΔS° = Σv p°(products) -Σv S°rreactants) - Calculations follow strategies used for ΔH° in ch 7 Second Law of Thermodynamics: - ΔSuniverse ΔS system+ ΔSsurroundings 0 - All spontaneous processes result in an increase in the entropy of the universe - For a process at constant P & T recall: - Q(p) = ΔH(sys) and q(surr) = -q(p) = - ΔH(sys) - If we are able to exchange heat reversibly, in infintesnemal steps with the surroundings, then.. - ΔSsurr = -q(rev)/T = - ΔS° - [seenotebook] - ΔS = pos.. might not necessarily happen on its own ? - ΔSuniverse > 0  spontaneous reaction Deriving Free Energy, ΔG - ΔS(universe) = ΔS(system) + ΔS(surroundings) > 0 - ΔS(surr) = - ΔH(system)/T - Electrolysis – adding in another positive ΔS surroundings…? H2O  H2 + O2 - ΔS(universe) = ΔS(system) + ΔS(surroundings) = ΔS(system) - ΔH(system)/T TΔS(universe) = ………… [incomplete] -TΔS(universe) = ΔH(system) -= TΔS(system) - ΔG = -TΔS(universe) - ΔS(univ) > 0 for spontaneous processes; ΔS being neg means its being ordered - Criterion for spontaneous change: neg change in gibbs free energy, ΔG - ΔG = ΔH - TΔS Gibbs free energy, G ; a state function of the system - Enthalpy (system gaining or losing energy, always want to go down in energy; losing energy is favourable) and entropy (systems always want to increase in entropy and become more disordered, disorder is favourable) component - 3. Positive ΔS is good, increasing entrop
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