Unit 1: Energy Changes in Chemical Reactions Spontaneous Process & Entropy

Spontaneous reactions → reactions that DOES occur under the given set of conditions Eg. ▪ waterfall runs downhill, but never up ▪ a lump of sugar spontaneously dissolve in a cup of hot coffee, but dissolved sugar does not reappear in its original form all by itself ▪ iron exposed to water and oxygen forms rust, but rust doesn’t spontaneously change back to iron

Non-spontaneous reactions → reactions that DOES NOT occur under the given set of conditions NOTE: Spontaneous does not mean instantaneous. Spontaneous means the reaction tends to occur, but not necessary right away. Eg. The rust on a car develops spontaneously but not instantaneous. (i.e. rust will develop over time, but not immediately)

In general, exothermic reactions tend to proceed spontaneously. However, some endothermic reactions are spontaneous even though the products are less energetically stable than the reactants. Example: 2 NH4NO3(aq) + Ba(OH)2(aq) + energy → 2 NH4OH(aq) + Ba(NO3) 2(aq)

Why do reactions yielding less stable products occur spontaneously? The answer to this question is ENTROPY. Entropy (S) is a measure of disorder or randomness of a system. → increases when disorder increases → ∆S = Sproducts – Sreactants → when entropy increases in a rxn… Sproducts > Sreactants , ∆S > 0 (i.e. ∆S is positive) → when entropy decreases in a rxn… Sproducts < Sreactants , ∆S < 0 (i.e. ∆S is negative)

Entropy is increased when: 1. More molecules are formed. It has been observed that a change that results in a ↑ in S is more likely to occur spontaneously than a change in which S ↓. Entropy is increased when: 1. More molecules are formed. 2. A liquid is formed from solids. 3. A gas is formed from either liquids or solids. 4. A mixture is formed. 5. If the volume of gas increases.

NOTE: → All substances have a positive entropy (i.e. some sort of disorder). → Only a perfect crystal at 0K would have zero entropy (defined as absolute zero entropy). → However, ∆S values can be negative because it is a measure of change, not an absolute value.

Examples: Predicting the Sign of ∆S 1. Solid carbon dioxide sublime into gaseous carbon dioxide. 2. N2O4(g) → 2NO2(g) 3. The synthesis reaction between oxygen and hydrogen forms liquid water.

∆S > 0 (∆S is positive), b/c the particles of a gas are more randomly distributed. → rxn. tends to occur spontaneously ∆S > 0 (∆S is positive), b/c 1 mole of N2O4(g) yields 2 moles of NO2(g) and the state of reactant and product remains the same. → rxn. tends to occur spontaneously 2H2(g) + O2(g) → 2H2O(l) ∆S < 0 (∆S is negative), b/c 3 moles of reactants yields 2 moles of product and the reactants are gases while the product is a liquid. → rxn. tends to be nonspontaneous.

Standard entropy (S°) is the entropy possessed by a substance at SATP (101 kpa, 25°C). Table 17-1 (p. 407) in your textbook has a list of standard entropies for common substances We calculate and predict whether a reaction is likely to occur spontaneously or not using the formula: ∆S° = S°products  S°reactants

Examples: 1. C(s-graphite) → C(s-diamond) S° = 5.694 J/K•mol 2.439 J/K•mol Prediction: very little change ∆S° = S°products  S°reactants = 2.439 J/K•mol  5.694 J/K•mol =  3.225 J/K•mol ∆S < 0, thus reaction is not spontaneous.

Prediction: very little change ∆S° = S°products  S°reactants 2. H2(g) + Cl2(g) → 2HCl(g) S° = 130.6 J/K•mol 223.0 J/K•mol 186.7 J/K•mol Prediction: very little change ∆S° = S°products  S°reactants = 2(186.7 J/K•mol)  1(130.6 J/K•mol + 223.0 J/K•mol) = 19.8 J/K•mol (∆S > 0, thus reaction is spontaneous)

Assignment: Pg. 407; # 8 (a) & (b), Pg. 425; # 29  34