Book: UTC: 20 Flashcards
How the tendency of a process to occur by itself is distinct from how long it takes to occur.
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The distinction between a spontaneous and a nonspontaneous change.
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Why the first law of thermodynamics and the sign of ∆Hº cannot predict the direction of a spontaneous change.
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How entropy (S) of a system is defined by the number of microstates over which its energy is dispersed.
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How entropy is alternatively defined by the heat absorbed (or released) at constant T in a reversible process.
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The criterion for spontaneity according to the second law of thermodynamics: that a change increases S_univ.
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How absolute values of standard molar entropies (Sº) can be obtained because the third law of thermodynamics provides a zero point.
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How temperature, physical state, dissolution, atomic size, and molecular complexity influences Sº values.
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How ∆Sº_rxn is based on the difference between the summed Sº values for the reactants and those for products.
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How the surroundings add heat to or remove heat from the system and how ∆S_surr influences overall ∆Sº_rxn.
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The relationship between ∆S_surr and ∆H_sys.
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How reactions proceed spontaneously toward equilibrium (∆S_univ > 0) but proceed no further at equilibrium (∆S_univ = 0).
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How the free energy change (∆G) combines a system’s entropy and enthalpy changes.
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How the expression for the free energy change is derived from the second law.
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The relationship between ∆G and the maximum work a system can perform and why this quantity of work is never performed in a real process.
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