Initial and Equilibrium Concentrations for H2(g) + I2(g) 2HI(g) at 445 °C

Finding Equilibrium Concentrations When Given K and Initial Concentrations or Pressures STEP 1a: Prepare the ICE table. STEP 1b: Decide in which direction the reaction will proceed. Compare Q to K. (not absolutely necessary, but helpful) STEP 1c: Define the changes of all substances in terms of x. Let x = the concentration gain or loss of a species whose coefficient is 1. (not required, but convenient) If no species has a coefficient of 1, define a 2x or 3x etc. The sign in front of x is negative for species on the side the reaction is proceeding away from, + for substances on the side the reaction is proceeding toward. STEP 2: Substitute into the Law of Mass Action equation Solve for x. For second-order equations, take square roots of both sides or use the quadratic formula. Simplify and approximate the answer for very large or small equilibrium constants, if possible. STEP 3: Calculate final equilibrium values and “check”.

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*Assume you calculated Q and it was smaller than K Chapter 15, Unnumbered Table 1, Page 698

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Relationships between K’s of Related Chemical Equations—I (Reversing Equation) When the reaction is written backward, the equilibrium constant is inverted. For the reaction aA + bB cC + dD, the equilibrium constant expression is as follows: For the reaction cC + dD aA + bB, the equilibrium constant expression is as follows:

Relationships between K’s of Related Chemical Equations—II (Multiplying an Equation) When the coefficients of an equation are multiplied by a factor, the equilibrium constant is raised to that factor. For the reaction aA + bB cC, the equilibrium constant expression is as follows: For the reaction 2aA + 2bB 2cC, the equilibrium constant expression is as follows:

EXAMPLE 15.2 Manipulating the Equilibrium Constant to Reflect Changes in the Chemical Equation Consider the chemical equation and equilibrium constant for the synthesis of ammonia at 25 C : Calculate the equilibrium constant for the following reaction at 25 C : SOLUTION Begin by reversing the given reaction and taking the inverse of the value of K. x ½ Next, multiply the reaction by ½ and raise the equilibrium constant to the ½ power. Calculate the value of K

Relationships between K’s of Related Chemical Equations—III (Summing Equations) When you add equations to get a new equation, the equilibrium constant of the new equation is the product of the equilibrium constants of the old equations. For the reactions (1) aA bB and (2) bB cC, the equilibrium constant expressions are as follows: For the reaction aA cC, the equilibrium constant expression is as follows:

FOR MORE PRACTICE 15.2 Predict the equilibrium constant for the first reaction given the equilibrium constants for the second and third reactions:

Equilibrium Constants for Reactions Involving Gases The concentration of a gas in a mixture is proportional to its partial pressure. Therefore, the equilibrium constant can be expressed as the ratio of the partial pressures of the gases. For aA(g) + bB(g) cC(g) + dD(g), the equilibrium constant expressions are as follows: or

Approximations to Simplify the Math When the equilibrium constant is very small, the position of equilibrium favors the reactants. For relatively large initial concentrations of reactants, the reactant concentration will not change significantly when it reaches equilibrium. Assuming the reaction is proceeding forward [A]equilibrium = ([A]initial  ax)  [A]initial We are approximating the equilibrium concentration of reactant to be the same as the initial concentration.

Checking the Approximation and Refining as Necessary We can check our approximation by comparing the approximate value of x to the initial concentration. If the approximate value of x is less than 5% of the initial concentration, the approximation is valid.