Predicting the Direction of a Reaction

Equilibrium constants and reaction quotients can be used to predict whether a reaction will favor the products or the reactants.

Learning Objective

Evaluate whether a chemical reaction has reached equilibrium from the reaction coefficient (Q) and the equilibrium constant (K), and use the latter to predict whether the reaction will favor the reactants or products

Key Points

Equilibrium constants are used to predict whether a reaction will favor the products or the reactants. If K > 1, the products dominate the mixture. If K < 1, the reactants dominate the mixture.
If K > 1, the reaction will occur in the forward direction and will occur mostly to completion (it will use up almost all the reactants). If K < 1, it will occur in the reverse direction and will use up almost all the products, turning them into the reactants.
If a reaction is not at equilibrium, you can use the reaction quotient, Q, to see where the reaction is in the pathway. If Q > K, the reactants are favored. If Q < K, the products are favored. If Q = K, the reaction is at equilibrium.

Term

equilibrium
the state of a reaction in which the rates of the forward and reverse reactions are the same

Full Text

Predicting the Direction of a Reaction

Equilibrium constants can be used to predict whether a reaction will favor the products or the reactants. K is the equilibrium ratio of products to reactants. The value for K is large (greater than one) when the products dominate the mixture and small (less than one) when the reactants dominate the mixture:

K > 1: products are favored
K = 1: neither reactants nor products are favored
K < 1: reactants are favored

If the reaction favors the products, it will occur in the forward (left-to-right) direction. If K is very large, the reaction will occur mostly to completion, using up almost all the reactants. If the reaction favors the reactants, it will occur in the reverse (right-to-left) direction. If K is very small, the reaction will use up almost all the products and make them into reactants.

Reaction Quotients and Predicting Direction

The equilibrium constant is only used when a reaction has reached equilibrium. If a reaction is not at equilibrium, you can use the reaction quotient, Q, to see where the reaction is in the pathway:

Q > K: reactants are favored
Q = K: reaction is at equilibrium
Q < K: products are favored

If you know the equilibrium constant for a reaction, and you know all the concentrations, you can predict in what direction the reaction will proceed.

Example 1: If 1.0 x 10^-2 moles of hydrogen gas and 1.0 x 10^-2 moles of iodine gas are placed in a 1-liter flask at 448 °C with 2.0 x 10^-3 moles of HI , will more HI be produced?

Hydrogen Iodide

Space-filling model of hydrogen iodide

Solution: The balanced equation for the reaction is:

$H_2 (g) + I_2 (g) \rightleftharpoons 2HI (g)$

The reaction quotient under these conditions is:

$Q = \frac{[HI]^{2}}{[H_{2}][I_{2}]}$

$Q = \frac {2.0 \cdot 10^{-3}}{(1.0 \cdot 10^{-3})^{2}}=0.040$

This value is smaller than the equilibrium value of 50.53, which tells us there is an excess amount of reactants present. Therefore, the reaction will proceed in the forward direction.

Example 2: If only 1.0 x 10^-3 moles of H₂ and 1.0 x 10^-3 moles I₂ had been used, together with 2.0 x 10^-2 mole of HI, would more HI have been produced spontaneously?

Solution: You can verify that Q = 400. This value is greater than the equilibrium value: there is now too much of the products for equilibrium to exist. Therefore, the reverse reaction is favored.

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