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(1) Ice melts spontaneously above 0ºC with the absorption of heat from the surroundings.
ΔH = +6.0 kJ mol–1
Ice ⎯⎯
⎯→ H 2 0 (l) ,
(2) NaCl dissolves spontaneously in water with the absorption of heat.
NaCl (s) + aq ⎯⎯
→ Na + (aq) + Cl – (aq) , ΔH = +3.9 kJ mol–1
Is it possible for a reaction to be nonspontaneous yet exothermic? Explain with example.
Normally exothermic process are spontaneous but there are some processes which are nonspontaneous inspite of being exothermic. For example : H2O(l) at 1 atm. and 0ºC forms ice.
Here both the phases are in equilibrium. This process is non-spontaneous.
H 2 O(l) H 2 O(s) ;
H < 0.
Predict the sign of ΔS in the following processes. Give reasons for your answer.
(a) N 2 O 4 (g) ⎯⎯
→ 2NO 2 (g)
sign of ΔS is +ve as disorder in the product are larger than in the reactants.
(b) Fe 2O 3 (s) + 3H 2 (g) ⎯⎯
→ 2Fe(s) + 3H 2 O(g)
reaction is in equilibrium, ΔS = 0.
(c) N 2 (g) + 3H 2 (g) ⎯⎯
→ 2NH 3 (g)
sign of ΔS is –ve, as disorder in the product are less than in the reactants.
(d) MgCO 3 (s) ⎯⎯
→ MgO(s) + CO 2 (g)
ΔS is +ve, as disorder increases in the products.
(e) CO 2 (g) ⎯⎯
→ CO 2 (s)
ΔS is –ve as disorder decreases in the product.
(f) Cl 2 (g) ⎯⎯
→ 2Cl(g)
ΔS is +ve, as one mole of Cl2(g) gives two moles of Cl(g) atoms and hence, disorder increases.
What can be said about the spontaneity of reactions when
(a) ΔH and ΔS are both positive.
Reactions become spontaneous at high temperature When T ⋅ ΔS > ΔH
(b) ΔH and ΔS are both negative.
Reactions become spontaneous at low temperatures When T ⋅ ΔS < ΔH
(c) ΔH is positive and ΔS is negative.
Reactions are nonspontaneous at all temperatures.
(d) ΔH is negative and ΔS is positive.
Reactions are spontaneous at all the temperatures.
Unique Solutions ®
S.Y.J.C. Science - Chemistry - Part I
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