Understanding Points
Reactivity 1.4.1—Entropy, S, is a measure of the dispersal or distribution of matter and/or energy in a system. The more ways the energy can be distributed, the higher the entropy. Under the same conditions, the entropy of a gas is greater than that of a liquid, which in turn is greater than that of a solid.
Reactivity 1.4.2—Change in Gibbs energy, ΔG, relates the energy that can be obtained from a
chemical reaction to the change in enthalpy, ΔH, change in entropy, ΔS, and absolute temperature, T.
Reactivity 1.4.3—At constant pressure, a change is spontaneous if the change in Gibbs energy, ΔG, is negative.
Reactivity 1.4.4—As a reaction approaches equilibrium, ΔG becomes less negative and finally
reaches zero.
Entropy: the distribution of available energy among the particles
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The more ways the energy can be distributed, the higher the entropy.
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A measure of “disorder” in a system
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Units: J K-1 mol-1
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If the reaction becomes more disordered → positive entropy → ∆S= +ve
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If the reaction becomes more ordered → negative entropy → ∆S= -ve
Factors that increase entropy of a system
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Changes of state
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solid < liquid < gas increasing order of entropy
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e.g. CaCO3 (s) → CaO (s) + CO2 (g) ∆S= +ve
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Number of Moles
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Increasing the number of moles, increases the entropy of a system
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e.g. N2O4(g) → 2NO2 (g) ∆S= +ve
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Increasing movement of particles (e.g. by increasing temperature)
Calculating entropy changes within a system
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Unlike enthalpy, the absolute values of entropy can be directly measured
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∆S⊖ = ∆S⊖products - ∆S⊖reactants
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e.g. calculate the entropy change in the following reaction
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3H2 (g) + N2 (g) ⇌ 2NH3 (g
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∆S⊖ = 2(192) - [3(131) + 192] = -201JK-1mol-1
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∆S⊖ <0, Reaction has become more ordered
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Spontaneity is a reaction’s ability to proceed and move towards completion without the application of work (self-sustained reaction without additional input of energy).
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Change in Gibbs free energy ΔG: a measure of spontaneity
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ΔG⊖ = ΔH⊖ - TΔS⊖
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If ΔG < 0 the reaction is spontaneous
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If ΔG = 0 the reaction is at equilibrium
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If ΔG > 0 the reaction is not spontaneous
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the more negative the value of ΔG , the further the equilibrium lies to the right
ΔH | ΔS | T | ΔG | Spontaneity |
- | + | low & high | - | spontaneous |
+ | - | low & high | + | not spontaneous |
- | - | low | - | spontaneous at low temp,
magnitude of TΔS < ΔH |
high | + | not spontaneous | ||
+ | + | low | + | not spontaneous |
high | - | spontaneous at high temp,
magnitude of TΔS > ΔH |
Calculating Gibbs Free Energy Changes
1. From Enthalpy and Entropy Values
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Is the following reaction spontaneous at 500K? If not at what temperature does it become spontaneous?
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CaCO3 (s) → CaO (s) + CO2 (g)
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The standard enthalpies of formation of CaCO3 , CaO, CO2 are -1207, -636, -394 kJmol-1
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The standard entropies of CaCO3 , CaO, CO2 are 93, 40, 214 JK-1mol-1
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Enthalpy change ( ∆H⊖ = ∆H⊖f products - ∆H⊖f reactants)
→ [-636+(-394)] - (-1207) = + 177 kJmol-1
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Entropy change (∆S⊖ = ∆S⊖products - ∆S⊖reactants )
→ (40 + 214) - (93) = + 161JK-1mol-1
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Gibbs Free Energy change
→ (+177000) - (500 x 161) = +96500 J mol-1
→ +96500Jmol-1/1000 = +96.5 kJ mol-1 ∴since ∆𝐺 = +ve, the reaction is non-spontaneous at 500K
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Temperature at which the reaction becomes spontaneous ,TΔS > ΔH
→ T x 161 = +177000
→ T = 177000 / 161
→ T = 1099K
2.From Gibbs Free Energy of formation changes
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∆G⊖ = ∆G⊖f products - ∆G⊖f reactants