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Spontaneous Processes
Spontaneous processes happen on their own without outside intervention.
Examples of spontaneous processes:
Ice melting at 0°C under 1 atm of pressure.
The oxidation of iron metal from moisture into iron (III) oxide.
Decomposition of hydrogen peroxide into water and oxygen.
Introduction to Entropy
The cause for spontaneity is associated with an increase in entropy, as demonstrated by the Gibbs Free Energy equation:
ΔG = ΔH - T·ΔS
Entropy: The probability of how much a system can arrange itself, or in other words, a measure of the disorder of a system. The more probabilities there are, the higher the disorder and entropy.
It is generally favorable for a system to naturally progress towards higher entropy—characterized by increased disorder of components within and the dispersion of energy.
It takes work to order and organize a system.
Real Life Analogies for Entropy:
Society: It takes effort to organize human society. If all forms of effort, such as legislation, were removed, human society would ultimately collapse into chaos.
Confetti: If you shoot confetti out, it is more probable for the particles to settle randomly than for them to all settle in a specific spot.
Change in Entropy (ΔS)
Processes can either increase or decrease the entropy of a system. This can be expressed with ΔS. The formula for calculating a process's change in entropy is:
ΔS = ΣS(products) - ΣS(reactants)
ΔS: Change in entropy.
ΣS: Sum of the standard entropy values of products/reactants. (J/(mol·K))
Predicting Entropy Changes
To predict if a process increased (ΔS > 0) or decreased (ΔS < 0) the entropy of a system, we have to know that certain conditions affect entropy:
Level of Entropy in Phases:
Solids < Liquids < Aqueous Solutions < Gases
The less energetic a phase is, the less freedom of movement it has. Therefore, particles of that phase have a lower probability of forming different arrangements.
Example #1: NH₄Cl (s) → NH₃ (g) + HCl (g)
The ΔS value for the decomposition of NH₄Cl would be positive, as 2 moles of gas are created from 1 mole of a solid.
Example #2: HCl (aq) + H₂O (l) → H₃O⁺ (aq) + Cl⁻ (aq)
The ΔS value for the reaction between HCl and H₂O would be positive, as 2 moles of aqueous ions are created from 1 mole of aqueous ions and 1 mole of a liquid.
Example #3: H₂O (l) → H₂O (s)
The ΔS value for the freezing of H₂O would be negative, as the substance turns from a liquid to a solid.
Temperature: Increasing the average kinetic energy of particles in a system builds disorder.
Having more kinetic energy within particles mean that they have a higher probability of forming different arrangements.
Example: Heating a container of N₂ (g) from 0°C to 20°C.
The ΔS value for the heating of N₂ would be positive, as the temperature of the system increased.
Moles: Increasing the amount of particles of a system increases entropy.
Having more particles within a system can lead to a higher entropy, as there are more particles to arrange and move.
Example: 2HI (g) → H₂ (g) + I₂ (g)
The ΔS value for the decomposition of HI would be positive, as 2 moles of gas are created from 1 mole of gas.
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