Dynamic Equilibrium

Summary: Dynamic equilibrium is the state reached by a reversible reaction in a closed system when the rate of the forward reaction and the rate of the backward reaction become equal, resulting in constant (but not necessarily equal) concentrations of all reactants and products, with no net macroscopic change despite ongoing molecular-level activity in both directions.

Tags: dynamic-equilibrium equilibrium reversible-reactions closed-system le-chatelier igcse-chemistry chemical-equilibria

Created: 2026-07-18

Dynamic equilibrium is a central concept in chemical equilibria describing the condition reached when a reversible reaction proceeds in a Closed System and the rate of the forward reaction becomes exactly equal to the rate of the backward reaction, so that the concentrations of all chemical species present — both reactants and products — remain constant over time, though they are not necessarily equal to one another. The term “dynamic” is essential: it distinguishes this state from a static equilibrium by emphasising that both the forward and reverse reactions continue to occur at the molecular level without interruption — reactant particles are still being converted into product particles, and product particles are still being converted back into reactant particles, but because these two processes happen at identical rates, there is no net change in the amounts of any substance present. Because the system is closed (no matter can enter or leave), the total mass of the system is conserved and the concentrations observed at equilibrium reflect the intrinsic thermodynamics of the reaction — specifically, the balance between enthalpy and entropy captured by the equilibrium constant, (K_c), which is fixed at a given temperature and independent of the starting concentrations. The position of equilibrium — whether it lies to the left (more reactants) or to the right (more products) — depends on the relative magnitudes of the forward and reverse rate constants, which in turn are governed by the activation energies of the two opposing pathways; a reaction with a lower activation energy in the forward direction will tend to favour products and produce a high equilibrium yield, as seen in the Haber process where careful manipulation of temperature, pressure, and catalyst conditions shifts the equilibrium position to maximise the yield of ammonia. Crucially, dynamic equilibrium can only be established in a closed system — if the system is open and products or reactants are allowed to escape, the forward and reverse rates can never become equal and a steady-state dynamic equilibrium is never achieved; this is why reactions such as the thermal decomposition of calcium carbonate (CaCO₃(s) ⇌ CaO(s) + CO₂(g)) reach equilibrium only when carried out in a sealed container, as the carbon dioxide gas would otherwise diffuse away and drive the reaction irreversibly to completion. At the macroscopic scale, observable properties of the system — colour, pressure, concentration, pH — remain constant at equilibrium, but at the microscopic scale, individual molecules are continuously undergoing reaction in both directions, a duality that is often illustrated by isotope-labelling experiments in which radioactive tracer atoms introduced after equilibrium is reached are found to distribute themselves between both reactant and product species, confirming that chemical exchange is ongoing even though no bulk change is detectable.

Definition and Core Principles

Dynamic equilibrium is defined by two simultaneous conditions:

Rate equality: (r_{\text{forward}} = r_{\text{reverse}}) — the speed at which reactants are consumed in the forward reaction exactly matches the speed at which products are consumed in the reverse reaction.
Constant concentrations: ([R_1], [R_2], \ldots, [P_1], [P_2], \ldots) are all unchanging over time — but not necessarily equal to each other.

Both conditions must be satisfied. A system in which concentrations happen to be momentarily steady but the rates are not equal (for example, during the initial stages of a reaction approaching equilibrium) is not at dynamic equilibrium. Equally, a system in which forward and reverse rates happen to be equal but concentrations are still changing due to changes in volume or temperature is not at dynamic equilibrium.

The Meaning of “Dynamic”

The word dynamic is the single most frequently tested conceptual point on this topic in the IGCSE examination. It means:

Reactions are still happening in both directions.
Individual particles are still colliding, bonds are still breaking and forming.
At any instant, some fraction of reactant molecules are being converted to products, and an equal fraction of product molecules are being converted back to reactants.
It is not a frozen or stopped state — it is a state of balanced, continuous flux.

A useful analogy is a supermarket with an equal number of customers entering and leaving per minute: the total number of people inside stays the same, but individuals are constantly coming and going. The observable “population” is constant, but there is continuous motion at the individual level.

Static vs. Dynamic Equilibrium

The distinction between static equilibrium (no change at any level) and dynamic equilibrium (no net macroscopic change, but ongoing microscopic change) is tested in both chemistry and physics contexts:

Feature	Static Equilibrium	Dynamic Equilibrium
Macroscopic properties	Constant	Constant
Microscopic activity	None	Continuous in both directions
Reaction rates	Zero (both directions)	Equal and non-zero
Examples	A book resting on a table (physics); a completed irreversible reaction	Reversible reaction in a closed system (chemistry)
Disturbed by…	A force (physics)	Changes in temperature, pressure, or concentration (Le Chatelier)

A chemical equilibrium is never static: even at equilibrium, molecules possess kinetic energy and undergo collisions that can result in reaction, provided the collision energy exceeds the activation energy in the relevant direction.

Requirements for Dynamic Equilibrium

Closed System

A Closed System is one in which no matter can enter or leave, although energy (heat, light) may be exchanged with the surroundings. A sealed flask, a stoppered test tube, or a sealed reaction vessel are all examples of closed systems.

If a system is open — for instance, a reaction carried out in an open beaker — gaseous products can escape into the atmosphere, making it impossible for the reverse reaction to proceed at a rate that catches up with the forward reaction. The reaction will go to completion (all limiting reactant consumed) rather than reaching equilibrium.

Example: CaCO₃(s) ⇌ CaO(s) + CO₂(g)

In a sealed container: CO₂ cannot escape; the reverse reaction (CaO + CO₂ → CaCO₃) can occur; equilibrium is established when the rate of decomposition equals the rate of recombination.

In an open container: CO₂ diffuses away; the reverse reaction cannot occur; all CaCO₃ eventually decomposes; no equilibrium is reached.

Reversible Reaction

Dynamic equilibrium can only be established for reactions that are reversible — that is, reactions in which the products can react together under the same conditions to regenerate the reactants. Reactions that go essentially to completion (irreversible reactions), such as combustion or neutralisation between a strong acid and a strong base, do not reach dynamic equilibrium because the reverse reaction has a negligibly small rate.

A reversible reaction is denoted by the ⇌ symbol:

[ a\text{A} + b\text{B} \rightleftharpoons c\text{C} + d\text{D} ]

Sufficient Time

Equilibrium is not instantaneous. A reversible reaction begins with only forward reaction occurring (because no products are present). As products accumulate, the reverse reaction begins to compete. Over time, as reactant concentrations fall and product concentrations rise, the forward rate decreases and the reverse rate increases until they meet. The time required depends on the kinetics (activation energies, temperature, catalyst presence) and may range from milliseconds to days.

Constant Temperature

Because the equilibrium constant (K_c) (and (K_p) for gaseous systems) depends on temperature, a system at dynamic equilibrium must be at a uniform and constant temperature. If the temperature changes, the equilibrium position shifts (see Le Chatelier’s Principle below) and the system re-establishes equilibrium with a new set of constant concentrations.

Graphical Representation

The approach to equilibrium and the attainment of the equilibrium state can be represented graphically.

Rate vs. Time

A rate-time graph shows the forward and reverse reaction rates as functions of time:

At (t = 0): forward rate is at its maximum (reactant concentrations are highest); reverse rate is zero (no products present).
As time progresses: forward rate decreases (reactants consumed); reverse rate increases (products accumulate).
At equilibrium: the two rate curves merge into a single horizontal line; forward and reverse rates are equal and constant.

Rate
 ↑
 |  Forward rate (decreasing)
 |  ╲
 |   ╲___________ Equilibrium (rates equal)
 |   ╱
 |  ╱  Reverse rate (increasing)
 | ╱
 └────────────────────────────→ Time

Concentration vs. Time

A concentration-time graph shows how the amounts of reactants and products change:

Reactant concentration: starts high, decreases over time, plateaus at equilibrium.
Product concentration: starts at zero, increases over time, plateaus at equilibrium.
At equilibrium: all concentration curves become horizontal; no further net change.

Concentration
 ↑
 | Reactant (decreasing)
 | ╲
 |  ╲_____________
 |  ╱
 | ╱  Product (increasing)
 |╱________________
 └────────────────────────────→ Time

Key Exam Point

A common IGCSE graph-interpretation question asks students to explain why the curves become horizontal. The correct answer is not “the reaction has stopped” but rather “the forward and reverse reactions are occurring at the same rate, so there is no net change in concentration.”

Le Chatelier’s Principle

Le Chatelier’s Principle states that if a system at dynamic equilibrium is subjected to a change in conditions (temperature, pressure, or concentration), the position of equilibrium shifts in the direction that opposes (partially counteracts) the imposed change.

Le Chatelier’s principle is a qualitative predictive tool — it tells you the direction of the shift, but not the magnitude or the time required to re-establish equilibrium.

Effect of Concentration Changes

Increasing the concentration of a reactant: equilibrium shifts to the right (forward direction) to consume some of the added reactant, producing more products.
Increasing the concentration of a product: equilibrium shifts to the left (reverse direction) to consume some of the added product, regenerating more reactants.
Decreasing the concentration of a reactant (e.g., by removing it): equilibrium shifts to the left to replenish the removed reactant.
Decreasing the concentration of a product (e.g., by removing it): equilibrium shifts to the right to replenish the removed product.

The principle of continuously removing a product to drive a reversible reaction towards completion is exploited industrially — for example, in the Haber process, ammonia is liquefied and removed from the reaction vessel as it forms, pulling the equilibrium to the right and improving the overall yield.

Effect of Pressure Changes (Gaseous Systems Only)

Pressure changes only affect equilibria involving gases where there is a difference in the total number of gas molecules on the two sides of the equation.

Increasing pressure: equilibrium shifts towards the side with fewer gas molecules (moles of gas) to reduce the pressure.
Decreasing pressure: equilibrium shifts towards the side with more gas molecules to increase the pressure.

Example: N₂(g) + 3H₂(g) ⇌ 2NH₃(g)

Reactant side: 1 + 3 = 4 moles of gas

Product side: 2 moles of gas

Increasing pressure shifts equilibrium to the right (fewer gas molecules), favouring ammonia production.

Decreasing pressure shifts equilibrium to the left (more gas molecules), favouring the decomposition of ammonia.

If the number of gas molecules is the same on both sides (e.g., H₂(g) + I₂(g) ⇌ 2HI(g), where both sides have 2 moles of gas), pressure changes have no effect on the equilibrium position.

Effect of Temperature Changes

The effect of temperature depends on whether the forward reaction is exothermic or endothermic.

Increasing temperature: equilibrium shifts in the endothermic direction (to absorb the added heat).
Decreasing temperature: equilibrium shifts in the exothermic direction (to release heat and counteract the cooling).

Example: N₂(g) + 3H₂(g) ⇌ 2NH₃(g) ΔH = −92 kJ mol⁻¹ (exothermic forward)

Increasing temperature shifts equilibrium to the left (endothermic reverse direction), reducing ammonia yield.

Decreasing temperature shifts equilibrium to the right (exothermic forward direction), increasing ammonia yield — but at the cost of a slower rate of reaction, which is why the Haber process uses a compromise temperature of approximately 450°C.

Effect of a Catalyst

A catalyst provides an alternative reaction pathway with a lower activation energy. Crucially, it lowers the activation energy equally for both the forward and reverse reactions. As a result:

A catalyst speeds up both the forward and reverse reactions to the same extent.
A catalyst does not shift the position of equilibrium.
A catalyst does not change the equilibrium concentrations or the value of (K_c).
A catalyst does allow equilibrium to be reached more quickly.

Exam tip: A question asking “How does adding a catalyst affect the yield of ammonia in the Haber process?” tests precisely this distinction — the correct answer is that the catalyst has no effect on the yield, but it allows the equilibrium yield to be achieved in a shorter time, making the process economically viable.

Summary Table of Le Chatelier Effects

Change Imposed	Equilibrium Shifts…
Increase [reactant]	To the right (→)
Increase [product]	To the left (←)
Increase pressure (fewer gas moles on right)	To the right (→)
Increase pressure (more gas moles on right)	To the left (←)
Increase temperature (exothermic forward)	To the left (←)
Increase temperature (endothermic forward)	To the right (→)
Add catalyst	No shift (both rates increase equally)

The Equilibrium Constant ((K_c))

At a given temperature, the ratio of product concentrations to reactant concentrations (each raised to the power of its stoichiometric coefficient) is a constant known as the equilibrium constant, (K_c).

For the general reversible reaction:

[ a\text{A} + b\text{B} \rightleftharpoons c\text{C} + d\text{D} ]

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

Key features of (K_c):

(K_c) is temperature-dependent: changing the temperature changes the value of (K_c).
(K_c) is independent of concentration, pressure, and the presence of a catalyst.
A large (K_c) (much greater than 1) means the equilibrium position lies to the right (products favoured).
A small (K_c) (much less than 1) means the equilibrium position lies to the left (reactants favoured).
Pure solids and pure liquids are omitted from the (K_c) expression because their concentrations are constant.

While quantitative (K_c) calculations are more commonly encountered at A-Level, IGCSE students should understand the qualitative idea that at equilibrium, the ratio of products to reactants is fixed at a given temperature, and that this ratio is what shifts when conditions change.

Industrial Applications

Dynamic equilibrium and Le Chatelier’s principle are not merely academic concepts — they underpin the design of some of the world’s most important industrial chemical processes.

The Haber Process

The Haber process produces ammonia (NH₃) from nitrogen and hydrogen:

[ \text{N}_2(g) + 3\text{H}_2(g) \rightleftharpoons 2\text{NH}_3(g) \quad \Delta H = -92 \text{ kJ mol}^{-1} ]

Pressure: 200 atm — high pressure favours the forward reaction (4 moles gas → 2 moles gas), increasing yield. Higher pressures would give an even higher yield but would be uneconomical due to plant construction costs and safety risks.
Temperature: ~450°C — a compromise. Low temperature favours the exothermic forward reaction (higher equilibrium yield), but would make the reaction unacceptably slow. High temperature increases the rate but reduces yield. 450°C balances rate and yield.
Catalyst: Finely divided iron — speeds up the attainment of equilibrium without affecting the position. Allows a reasonable rate at the compromise temperature.
Product removal: Ammonia is continuously liquefied and removed, pulling the equilibrium to the right.

The Contact Process

The Contact Process produces sulfuric acid (H₂SO₄), with the key equilibrium step being the oxidation of sulfur dioxide:

[ 2\text{SO}_2(g) + \text{O}_2(g) \rightleftharpoons 2\text{SO}_3(g) \quad \Delta H = -197 \text{ kJ mol}^{-1} ]

Pressure: ~1—2 atm — the equilibrium already lies far to the right at moderate pressure (3 moles gas → 2 moles gas), so high pressure is unnecessary for an acceptable yield.
Temperature: ~450°C — again a compromise between rate (favoured by high temperature) and yield (favoured by low temperature, since the forward reaction is exothermic).
Catalyst: Vanadium(V) oxide (V₂O₅) — allows equilibrium to be reached rapidly at the operating temperature.

Both processes illustrate the same strategic tension between thermodynamics (equilibrium yield, favoured by low temperatures for exothermic reactions) and kinetics (rate of reaction, favoured by high temperatures), resolved through compromise conditions and the use of catalysts.

Experimental Demonstration

A classic classroom demonstration of dynamic equilibrium uses the reaction between bismuth(III) chloride (BiCl₃) and water:

[ \text{BiCl}_3(aq) + \text{H}_2\text{O}(l) \rightleftharpoons \text{BiOCl}(s) + 2\text{HCl}(aq) ]

Adding water shifts equilibrium to the right, producing a white precipitate of BiOCl.
Adding concentrated HCl shifts equilibrium to the left, dissolving the precipitate and reforming a clear solution.
The system can be cycled repeatedly between the two states by alternating additions of water and HCl, visually demonstrating the reversibility and the response of the equilibrium to concentration changes.

Another common demonstration uses the chromate-dichromate equilibrium:

[ 2\text{CrO}_4^{2-}(aq) + 2\text{H}^+(aq) \rightleftharpoons \text{Cr}_2\text{O}_7^{2-}(aq) + \text{H}_2\text{O}(l) ] (yellow chromate) ⇌ (orange dichromate)

Adding acid (H⁺) shifts equilibrium to the right, turning the solution orange.
Adding alkali (OH⁻, which removes H⁺) shifts equilibrium to the left, turning the solution yellow.
The colour change is instantaneous and reversible, providing a clear visual indicator of equilibrium shifts.