7.0 Chemical Equilibria
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Candidates should be able to:
(a) explain, in terms of rates of the forward and reverse reactions, what is meant by a reversible reaction and dynamic equilibrium
(b) state Le Chatelier’s Principle and apply it to deduce qualitatively (from appropriate information) the effects of changes in concentration, pressure or temperature, on a system at equilibrium
(c) deduce whether changes in concentration, pressure or temperature or the presence of a catalyst affect the value of the equilibrium constant for a reaction
(d) deduce expressions for equilibrium constants in terms of concentrations, Kc, and partial pressures, Kp [treatment of the relationship between Kp and Kc is not required]
(e) calculate the values of equilibrium constants in terms of concentrations or partial pressures from appropriate data
(f) calculate the quantities present at equilibrium, given appropriate data (such calculations will not require the solving of quadratic equations)
(g) show understanding that the position of equilibrium is dependent on the standard Gibbs free energy change of reaction, ∆G⦵ [Quantitative treatment is not required]
(h) describe and explain the conditions used in the Haber process, as an example of the importance of an understanding of chemical equilibrium in the chemical industry
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Reversible reactions can proceed in the forward and backward directions. It is never complete and will reach an equilibrium where both products and reactants are present.
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Dynamic equilibrium for a reversible reaction is achieved when the rates of both forward and reverse equations are equal and concentrations of both reactants and products are also equal.
7.1 Le Chatelier's Principle
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The principle states that when a change is made to a system in equilibrium, the system responds by opposing the change and a new equilibrium is formed.
7.2 Equilibrium Constants
For a reversible equation: aA + bB ⇆ cC + dD
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The equilibrium constant is a measure of the extent to which the products are formed from the reactants before equilibrium is reached.
Kc= [C]^c[D]^d/ [A]^a[B]^b
Kp=(C)^c(D)^d/ (A)^a(B)^b, for reactions involving gases.
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Units of Kc is (mol dm^-3)^c+d-a-b
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Units of Kp is (atm)^c+d-a-b
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The equilibrium constant is constant at a given temperature.
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Kc and Kp are only affected by temperature.
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Kc and Kp are not affected by changes in concentration, pressure and presence of a catalyst.
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A catalyst increases the rate of both forward and backward reaction equally.
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Equilibrium is reached in a shorter time, but there is no change in the amount of product formed.
Position of Equilibrium
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The position of the equilibrium is dependent on ΔG.
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ΔG = -RT ln K, where R = molar gas constant, T = absolute temperature in K and K = equilibrium constant
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When ΔG = 0, the reaction had reached equilibrium.
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The smaller the value of ΔG, the closer the reaction to equilibrium.
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7.3 Haber Process
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Ammonia gas is produced in the Haber process through the following reaction.
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N2 + 3H2 ⇆ 2NH3
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Pressure of 250 atm
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Higher yield of ammonia gas.
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Too expensive to maintain at 250 atm.
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Moderate Temperature of 450 °C
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Too low of a reaction can result in a lower rate of reaction.
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High temperatures can result in a lower yield of ammonia gas (Le Chatelier's Principle).
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Iron catalyst (finely divided) mixed with small amounts of K2O and Al2O3
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Increases the rate of reaction.
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N2 can be obtained from fractional distillation of air
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H2 can be obtained from catalytic cracking of alkanes or natural gas.
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