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Q and K: What's the difference?

All about chemical equilibrium

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On this page:

- What is the reaction quotient?
- How to visualize Q and K

- Example: dissociation of N2O4
- Example: vaporization of solid I2
- Example: vaporization of solid NH4Cl

- Summary
- Concept map

Given a chemical reaction system and its equilibrium constant, together with the concentrations or pressures of the reactants and products actually present, we often need to predict whether, and in which direction, a net change in composition will tend to take place. The concept of the reaction quotient makes this very easy to do, but the similarity of the algebraic forms of K and Q can easily lead to confusion. It is hoped that the more detailed treatment offered below will help.

[Rasmus Grønfeldt]

1 What is the equilibrium quotient?

In the previous section we defined the equilibrium expression for the reaction

a A + b B → c C + d D as

In the general case in which the concentrations can have any arbitrary values (including zero), this expression is called the reaction quotient (the term equilibrium quotient is also commonly used.) and its value is denoted by Q (or Qc or Qp if we wish to emphasize that the terms represent molar concentrations or partial pressures.) If the terms correspond to equilibrium concentrations, then the above expression is called the equilibrium constant and its value is denoted by K (or Kc or Kp.)

K is thus the special value that Q has when the reaction is at equilibrium

The value of Q in relation to K serves as an index how the composition of the reaction system compares to that of the equilibrium state, and thus it indicates the direction in which any net reaction must proceed.

For example, if we combine the two reactants A and B at concentrations of 1 mol L–1 each, the value of Q will be 0÷1=0. The only possible change is the conversion of some of these reacants into products. If instead our mixture consists only of the two products C and D, Q will be indeterminately large (1÷0) and the only possible change will be in the reverse direction.

It is easy to see (by simple application of the le Châtelier principle) that the ratio of Q/K immediately tells us whether, and in which direction, a net reaction will occur as the system moves toward its equilibrium state. A schematic view of this relationship is shown below:

Notice the double-ended arrow in the center. This is to remind you that when the reaction is at equilibrium, the forward and reverse reactions continue as always, but at equal rates, thus cancelling each other to yield a state of no net change. (Remember the Law of Mass Action!)

More formally, we simply look at which of the following describes the relative values of Q and K:

Q/K |
> 1 | Product concentration too high for equilibrium; net reaction proceeds to left.
= 1 | System is at equilibrium; no net change will occur.
< 1 | Product concentration too low for equilibrium; net reaction proceeds to right.

It is very important that you be able to work out these relations for yourself, not by memorizing them, but from the definitions of Q and K.

Problem Example

The equilibrium constant for the oxidation of sulfur dioxide is Kp = 0.14 at 900 K.

2 SO2(g) + O2(g) → 2 SO3(g)

If a reaction vessel is filled with SO3 at a partial pressure of 0.10 atm and with O2 and SO2 each at a partial pressure of 0.20 atm, what can you conclude about whether, and in which direction, any net change in composition will take place?

Solution:

The value of the equilibrium quotient Q for the initial conditions is

Since Q>K, the reaction is not at equilibrium, so a net change will occur in a direction that decreases Q. This can only occur if some of the SO3 is converted back into products. In other words, the reaction will "shift to the left".

2 A visual way of thinking about Q and K

The formal definitions of Q and K are quite simple, but they are of limited usefulness unless you are able to relate them to real chemical situations. The following diagrams illustrate the relation between Q and K from various standpoints. Take some time to study each one carefully, making sure that you are able to relate the description to the illustration.

### Example 1: Dissociation of dinitrogen tetroxide

For the reaction N2O4 → 2 NO2 , Kc = 0.0059 at 298K.

This equilibrium condition is represented by the red curve that passes through all points on the graph that satisfy the requirement that

Q = [NO2]2 / [N2O4] = 0.0059 .

There are of course an infinite number of possible Q's of this system within the concentration boundaries shown on the plot. Only those points that fall on the red line correspond to equilibrium states of this system (those for which Q = K ). The line itself is a plot of [NO2] that we obtain by rearranging the equilibrium expression [NO2] = ([N2O4]K)0.5.

If the system is initially in a non-equilibrium state, its composition will tend to change in a direction that moves it to one that is on the line.

Two such non-equilibrium states are shown. The state indicated by has
Q > K, so we would expect a net reaction that reduces Q by converting some of the NO2 into N2O4; in other words, the equilibrium "shifts to the left". Similarly, in state , Q < K, indicating that the forward reaction will occur.

The blue arrows in the above diagram indicate the successive values that Q assumes as the reaction moves closer to equilibrium. The slope of the line reflects the stoichiometry of the equation. In this case, one mole of reactant yields two moles of products, so the slopes have an absolute value of 2:1.

### Example 2: Phase-change equilibrium

One of the simplest equilibria we can write is that between a solid and its vapor. In this case (as will be explained more fully in the next lesson), the equilibrium constant is just the vapor pressure of the solid. Thus for the process

I2(s) → I2(g)

all possible equilibrium states of the system lie on the horizontal red line and is independent of the quantity of solid present (as long as there is at least enough to supply the relative tiny quantity of vapor.) So adding various amounts of the solid to an empty closed vessel (states and ) causes a gradual buidup of iodine vapor. Because the equilibrium pressure of the vapor is so small, the amount of solid consumed in the process is negligible, so the arrows go straight up and all lead to the same equilibrium vapor pressure.

### Example 3: Heterogeneous chemical reaction

The decomposition of ammonium chloride is a common example of a heterogeneous (two-phase) equilibrium. Solid ammonium chloride has a substantial vapor pressure even at room temperature:

NH4Cl(s) → NH3(g) + HCl(g)

Arrow traces the states the system passes through when solid NH4Cl is placed in a closed container. Arrow represents the addition of ammonia to the equilibrium mixture; the system responds by following the path back to a new equilibrium state which, as the Le Châtelier principle predicts, contains a smaller quantity of ammonia than was added. The unit slopes of the paths and reflect the 1:1 stoichiometry of the gaseous products of the rection.

Summary

Make sure you thoroughly understand the following essential ideas which have been presented above. It is especially imortant that you know the precise meanings of all the highlighted terms in the context of this topic.

- When arbitrary quantities of the different components of a chemical reaction system are combined, the overall system composition will not likely correspond to the equilibrium composition. As a result, a net change in composition ("a shift to the right or left") will tend to take place until the equilibrium state is attained.
- The status of the reaction system in regard to its equilibrium state is characterised by the value of the equilibrium expression whose formulation is defined by the coefficients in the balanced reaction equation; it may be expressed in terms of concentrations, or in the case of gaseous components, as partial pressures.
- The various terms in the equilibrium expression can have any arbitrary value (including zero); the value of the equilibrium expression itself is called the reaction quotient Q.
- If the concentration or pressure terms in the equilibrium expression correspond to the equilibrium state of the system, then Q has the special value K, which we call the equilibrium constant.
- The ratio of Q/K (whether it is 1, >1 or <1) thus serves as an index of how far the system is from its equilibrium composition, and its value indicates the direction in which the net reaction must proceed in order to reach its equilibrium state.
- When Q = K, then the equilibrium state has been reached, and no further net change in composition will take place as long as the system remains undisturbed.

Concept Map