Misconception 1: "Spontaneous means fast"

Reality: Thermodynamics (spontaneity) and kinetics (rate) are completely separate. C(diamond) → C(graphite) is thermodynamically spontaneous ($\Delta G < 0$) but takes millions of years. A catalyst can make a spontaneous reaction fast but cannot make a non-spontaneous reaction spontaneous.

Misconception 2: "Exothermic reactions are always spontaneous"

Reality: $\Delta H < 0$ alone does not guarantee spontaneity. If $\Delta S < 0$ and $T|\Delta S| > |\Delta H|$, then $\Delta G > 0$ (non-spontaneous). Example: some exothermic condensation reactions become non-spontaneous at high temperature.

Misconception 3: "Entropy always increases in the system"

Reality: The UNIVERSE's entropy increases for spontaneous processes, but the system's entropy can decrease. When water freezes, $\Delta S_{system} < 0$ — but $\Delta S_{surroundings} > 0$ (heat released warms the surroundings), so $\Delta S_{universe} > 0$.

Misconception 4: "$\Delta G$° = 0 at equilibrium"

Reality: $\Delta G = 0$ at equilibrium. $\Delta G^\circ$ is a constant at a given temperature, equal to $-RT\ln K$. $\Delta G^\circ = 0$ only if $K = 1$.

Misconception 5: "Free expansion is the same as reversible expansion"

Reality: Free expansion ($P_{ext} = 0$) is highly irreversible — the system jumps from initial to final state without passing through equilibrium states. Reversible expansion proceeds through infinitely many equilibrium states.

Misconception 6: "Bond enthalpy values are exact"

Reality: Bond enthalpies are average values over many compounds. They give approximate results. The actual $\Delta H$ from calorimetry is more accurate than from bond enthalpy calculations.

Misconception 7: "$C_p - C_v = R$ only for monoatomic gases"

Reality: $C_p - C_v = R$ holds for ALL ideal gases (monoatomic, diatomic, polyatomic). The derivation uses only $PV = nRT$.