Ch 5 covers chemical thermodynamics — internal energy, enthalpy, Hess's law, bond enthalpies, entropy, and Gibbs free energy — to predict whether a reaction will be spontaneous.
System + surroundings = universe. First law: ΔU = q + w. At constant pressure: ΔH = q_p (enthalpy). Exothermic: ΔH < 0 (heat released). Endothermic: ΔH > 0 (heat absorbed). Hess's law: ΔH for a reaction is the same regardless of the path. Standard enthalpy of formation: Δ_fH° of elements in standard state = 0. Bond enthalpy: energy needed to break a bond.
Entropy S: measure of randomness/disorder. 2nd law: ΔS_universe > 0 for spontaneous processes. ΔS = q_rev/T. Gibbs free energy: ΔG = ΔH − TΔS. ΔG < 0: spontaneous. ΔG = 0: equilibrium. ΔG > 0: non-spontaneous. A reaction with ΔH < 0 and ΔS > 0 is always spontaneous. Temperature determines spontaneity when ΔH and ΔS have the same sign.
Download: https://ncert.nic.in/textbook/pdf/kech105.pdf | Part I: https://ncert.nic.in/textbook/pdf/kech1ps.zip
Yes. If ΔS is positive and large enough that TΔS > ΔH, then ΔG = ΔH − TΔS < 0 (spontaneous). Example: dissolving ammonium nitrate in water is endothermic but spontaneous because entropy increase (disorder of ions) is large. Temperature also matters — endothermic reactions become spontaneous at higher temperatures.
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