Dynamic Equilibrium

Introduction: Why reactions do not always go to completion ⚖️

students, when you see a chemical reaction written on paper, it can look like the reactants just turn completely into products. In real life, many reactions do not behave that neatly. Some reactions go nearly to completion, but many important reactions in chemistry settle into a balanced state where both reactants and products are present at the same time. This state is called dynamic equilibrium.

In this lesson, you will learn how to:

explain the key ideas and vocabulary of dynamic equilibrium,
describe what happens at the particle level,
connect equilibrium to the ideas of how much reaction happens and how far a reaction goes,
use examples to understand how equilibrium appears in IB Chemistry HL.

A useful real-world picture is a busy train station 🚆. People are constantly arriving and leaving, but if the number arriving each minute equals the number leaving each minute, the total number of people in the station stays constant. Dynamic equilibrium works in a similar way. The system looks unchanged overall, but particles are still reacting in both directions.

What dynamic equilibrium means

Dynamic equilibrium is reached in a closed system when the rate of the forward reaction equals the rate of the reverse reaction. At that point, the concentrations of reactants and products stay constant over time, although the reactions continue.

The key word is dynamic. Nothing has stopped. Instead, two opposite processes happen at the same rate.

For a reversible reaction such as

$\mathrm{N_2O_4(g) \rightleftharpoons 2NO_2(g)}$

the forward reaction produces $\mathrm{NO_2}$ from $\mathrm{N_2O_4}$, and the reverse reaction turns $\mathrm{NO_2}$ back into $\mathrm{N_2O_4}$. At equilibrium, molecules keep changing from one form to the other, but the overall amounts remain constant.

This does not mean the amounts of reactants and products are equal. It only means their concentrations are no longer changing.

A few important terms:

Reversible reaction: a reaction that can proceed in both the forward and reverse directions.
Closed system: a system in which matter cannot enter or leave.
Equilibrium position: the relative amounts of reactants and products at equilibrium.
Equilibrium constant, $K$: a number that describes the ratio of product and reactant concentrations at equilibrium for a given temperature.

How equilibrium develops over time

Imagine starting with only reactants in the container. At first, the forward reaction is fastest because there are many reactant particles available. As products form, the reverse reaction becomes more likely because more product particles are present.

The system changes in three stages:

Initially: only reactants may be present, so the forward reaction is much faster than the reverse reaction.
Approaching equilibrium: product concentration increases, so the reverse reaction speeds up.
At equilibrium: the forward and reverse reaction rates are equal.

This can be shown with a concentration-time graph. The concentration of reactants decreases and then levels off, while the concentration of products increases and then levels off. The flat part of the graph does not mean the reaction has stopped; it means equilibrium has been reached.

For a reaction such as the Haber process,

$\mathrm{N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g)}$

equilibrium is very important because industrial chemists want to maximize the amount of ammonia produced while keeping the process practical and efficient.

The equilibrium constant and the idea of extent of reaction

The equilibrium constant links dynamic equilibrium to the broader IB theme of how far a reaction goes. It helps us judge whether products or reactants are favored at equilibrium.

For a general reaction

\mathrm{aA + bB \rightleftharpoons cC + dD}

the equilibrium constant in terms of concentration is

$K_c = \frac{[C]^c[D]^d}{[A]^a[B]^b}$

where the square brackets mean concentration in $\mathrm{mol\,dm^{-3}}$.

A large $K_c$ means the equilibrium mixture contains relatively more products. A small $K_c$ means the equilibrium mixture contains relatively more reactants. Importantly, $K_c$ depends only on temperature.

The extent of reaction describes how far a reaction proceeds before equilibrium is reached. A reaction with a very large equilibrium constant lies far to the product side, meaning the extent of reaction is large. A reaction with a very small equilibrium constant lies far to the reactant side, meaning the extent of reaction is limited.

However, a large $K_c$ does not mean the reaction is fast. Speed and position are different ideas. A reaction can be very slow and still have products favored at equilibrium.

What can change equilibrium? Le Châtelier’s principle

students, equilibrium is stable, but it is not fixed forever. If conditions change, the system responds to reduce the effect of the change. This is called Le Châtelier’s principle.

The main changes are concentration, pressure, and temperature.

1. Changing concentration

If you add more reactant, the system shifts in the direction that uses up that reactant. If you remove product, the system shifts to replace it.

Example: In the equilibrium

$\mathrm{Fe^{3+}(aq) + SCN^{-}(aq) \rightleftharpoons FeSCN^{2+}(aq)}$

adding more $\mathrm{Fe^{3+}}$ can make the red complex ion $\mathrm{FeSCN^{2+}}$ form more strongly. This is often used in school experiments to show visible equilibrium shifts.

2. Changing pressure

Pressure changes matter only for reactions involving gases. Increasing pressure favors the side with fewer gas molecules because that reduces pressure.

For the Haber process,

$\mathrm{N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g)}$

there are $4$ moles of gas on the left and $2$ on the right. Increasing pressure shifts the equilibrium to the right, increasing ammonia yield.

3. Changing temperature

Temperature is special because it changes $K_c$ itself.

For an exothermic forward reaction, heat acts like a product.
For an endothermic forward reaction, heat acts like a reactant.

If temperature increases, equilibrium shifts in the endothermic direction.

For the Haber process, the forward reaction is exothermic. Lower temperature favors ammonia formation, but too low a temperature makes the reaction too slow. This is a classic example of balancing rate and equilibrium position.

Dynamic equilibrium and reaction rates

This topic connects strongly to rates of reaction, another part of Reactivity 2. At the start of a reversible reaction, the forward rate is high and the reverse rate is low. As products build up, the reverse rate increases. At equilibrium:

\text{rate of forward reaction} = \text{rate of reverse reaction}

This equality of rates is why concentrations remain constant.

The collision theory helps explain this. Particles must collide with enough energy and the correct orientation for reaction to happen. At equilibrium, collisions are still happening all the time, but the overall number of successful forward and reverse reactions is balanced.

This shows why equilibrium is dynamic rather than static. A static system would mean no reactions happen at all, which is not the case.

Common IB Chemistry HL pitfalls

students, students often confuse equilibrium with equal amounts. Remember:

equal rates, not equal concentrations,
a closed system is required,
equilibrium can be reached from either side,
catalysts do not change the equilibrium position.

A catalyst speeds up both the forward and reverse reactions by lowering the activation energy. It helps the system reach equilibrium faster, but it does not change $K_c$ or the final equilibrium composition.

Also, solids and pure liquids are not included in equilibrium expressions because their concentrations are effectively constant.

For example, in

$\mathrm{CaCO_3(s) \rightleftharpoons CaO(s) + CO_2(g)}$

only $\mathrm{CO_2(g)}$ appears in the equilibrium expression:

$K_c = [CO_2]$

under a given temperature.

Why dynamic equilibrium matters in the real world

Dynamic equilibrium is important in industry, biology, and the environment 🌍. The Haber process makes fertilizers that support global agriculture. The balance between dissolved gases and gases above liquids affects oceans and drinks. In the body, many biochemical processes involve reversible reactions and equilibria that help maintain stable conditions.

In acid-base chemistry, equilibrium ideas help explain buffers. In atmospheric chemistry, equilibrium helps describe the formation and breakdown of gases such as ozone. These examples show that equilibrium is not just a classroom idea. It helps explain how systems stay balanced while still changing.

Conclusion

Dynamic equilibrium is one of the most important ideas in Reactivity 2 because it connects how fast reactions happen with how far they go. In a closed system, a reversible reaction reaches a state where the forward and reverse rates are equal, concentrations stay constant, and reactions continue at the particle level. Le Châtelier’s principle helps predict how the system responds to changes in concentration, pressure, and temperature. For IB Chemistry HL, this idea is essential for understanding equilibrium position, the equilibrium constant, and how chemists design processes that produce useful amounts of products.

Study Notes

Dynamic equilibrium happens in a closed system.
At equilibrium, the rate of the forward reaction equals the rate of the reverse reaction.
Concentrations stay constant, but reactions do not stop.
A reversible reaction can reach equilibrium from either side.
The equilibrium position shows the relative amounts of reactants and products.
The equilibrium constant is written for a general reaction as

$ K_c = \frac{[C]^c[D]^d}{[A]^a[B]^b}$

A large $K_c$ means products are favored; a small $K_c$ means reactants are favored.
Changing concentration, pressure, or temperature can shift equilibrium.
Increasing pressure favors the side with fewer gas molecules.
Temperature changes can shift equilibrium and change $K_c$.
A catalyst speeds up reaching equilibrium but does not change $K_c$.
Dynamic equilibrium links the ideas of amount of change, rate of reaction, and extent of reaction in Reactivity 2.