equilibrium

equilibrium - Chapter 15. Chemical Equilibrium 15.1 The...

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Unformatted text preview: Chapter 15. Chemical Equilibrium 15.1 The Concept of Equilibrium Consider colorless frozen N 2 O 4 . At room temperature, it decomposes to brown NO 2 . N 2 O 4 ( g ) → 2NO 2 ( g ) At some time, the color stops changing and we have a mixture of N 2 O 4 and NO 2 . Dynamic Chemical equilibrium is the point at which the concentrations of all species are constant. Assume that both the forward and reverse reactions are elementary processes. We can write rate expressions for each reaction. Forward reaction: N 2 O 4 (g) → 2 ΝΟ 2 ( g) Rate f = k f [N 2 O 4 ] k f = rate constant (forward reaction) Reverse reaction: 2NO 2 (g) → N 2 O 4 (g) Rate r = k r [NO 2 ] 2 k r = rate constant (reverse reaction) Place some pure N 2 O 4 into a closed container. As N 2 O 4 reacts to form NO 2 , the concentration of N 2 O 4 will decrease and the concentration of NO 2 will increase. Thus, we expect the forward reaction rate to slow and the reverse reaction rate to increase. Eventually we get to equilibrium where the forward and reverse rates are equal. At equilibrium: k f [N 2 O 4 ] = k r [NO 2 ] 2 Rearranging, we get: k f / k r = a constant = K eq At equilibrium the concentrations of N 2 O 4 and NO 2 do not change. This mixture is called an equilibrium mixture . This is an example of a dynamic equilibrium. A dynamic equilibrium exists when the rates of the forward and reverse reactions are equal. No further net change in reactant or product concentration occurs. The double arrow ↔ implies that the process is dynamic. Chemical Equilibrium page 2 15.2 The Equilibrium Constant Consider the Haber process : N 2 ( g ) + 3H 2 ( g ) ↔ 2NH 3 ( g ) It is used for the preparation of ammonia from nitrogen and hydrogen. The process is carried out at high temperature (500˚C) and pressure (200 atm) in the presence of a catalyst. Ammonia is a good source of fixed nitrogen for plants. Much of the NH 3 produced industrially is used as a fertilizer. If we start with a mixture of nitrogen and hydrogen (in any proportions), the reaction will reach equilibrium with constant concentrations of nitrogen, hydrogen, and ammonia. However, if we start with just ammonia and no nitrogen or hydrogen, the reaction will proceed and N 2 and H 2 will be produced until equilibrium is achieved. No matter what the starting compositions of reactants and products are, the equilibrium mixture contains the same relative concentrations of reactants and products. Equilibrium can be reached from either direction! We can write an expression for the relationship between the concentration of the reactants and products at equilibrium. This expression is based on the law of mass action . For a general reaction, a A + b B ↔ dD + eE The equilibrium expression is given by: Where K c is the equilibrium constant, K eq ....
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This note was uploaded on 12/10/2009 for the course CHEM 1111 at Colorado.

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equilibrium - Chapter 15. Chemical Equilibrium 15.1 The...

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