Q.iii. In which of the following, entropy of the system decreases? a. Crystallization of liquid into solid
b. Temperature of crystalline solid is increased from 0 K to 115 K
c. $\mathrm{H_2(g) \rightarrow 2H(g)}$
d. $\mathrm{2\,NaHCO_3(s) \rightarrow Na_2CO_3(s) + CO_2(g) + H_2O(g)}$

Entropy (S) is a state function that measures the degree of molecular disorder or randomness in a system -- the greater the disorder, the larger the entropy. This is illustrated by water's three states: highly-ordered solid ice, less-ordered liquid water, and highly-disordered gaseous steam, so entropy increases (ΔS > 0) on melting and on vaporization; similarly, dissolving solid iodine in water (I2(s) + aq to I2(aq)) or breaking H2 molecules apart into free atoms (H2(g) to 2H(g)) both increase disorder and so have ΔS > 0, while a process that brings particles together into a more ordered arrangement (such as 2H(g) combining into H2(g), or a gas condensing to a liquid) decreases disorder and has ΔS < 0. Quantitatively, the entropy change of a system equals the heat transferred to it in a REVERSIBLE manner (Qrev) divided by the absolute temperature T at which the transfer occurs: ΔS = Qrev / T, giving entropy the SI unit of J K-1 (or J K-1 mol-1 for a molar quantity). Adding heat increases molecular motion and hence disorder, so ΔS is proportional to the heat added; but the SAME amount of heat causes a larger increase in randomness at a LOW temperature than at a high one, so ΔS is inversely related to the temperature at which the heat transfer occurs.