Closed System

Summary: A system where no matter can enter or leave but energy can be exchanged with the surroundings — essential for dynamic equilibrium. Tags: igcse chemistry Created: 2026-07-18

A closed system is a physical system in which no matter can enter or leave the system boundary, though energy — in the form of heat, light, or work — may still be exchanged with the surroundings. This concept is central to chemistry because many fundamental processes, especially reversible reactions capable of reaching dynamic equilibrium, can only occur as described when the system is closed. If a gaseous product escapes from an open beaker, for instance, the reverse reaction is starved of one of its reactants and cannot proceed, meaning the forward reaction goes to completion instead of establishing equilibrium. The requirement of a closed system distinguishes equilibrium-governed reactions from those that run irreversibly: the Haber process must be carried out in a sealed reactor so that ammonia does not escape before it can decompose back into nitrogen and hydrogen, while the hydration and dehydration of copper(II) sulfate can only reach an equilibrium state when water vapour is trapped inside a sealed container. In the IGCSE syllabus, students are expected to identify whether a given scenario constitutes a closed system and to predict whether equilibrium can be established under those conditions.

Definition

A system, in thermodynamic terms, is the specific part of the universe under consideration — for a chemist, this is typically the contents of a reaction vessel. The system boundary separates it from the surroundings (everything else). Systems are classified based on what can cross this boundary:

• Open system: Both matter and energy can be exchanged with the surroundings. Example: an uncovered beaker of boiling water — steam (matter) escapes and heat (energy) is lost to the air.
• Closed system: Energy can cross the boundary, but matter cannot. Example: a sealed glass flask containing a reversible reaction — heat can flow in or out through the glass walls, but no gases, liquids, or solids can enter or escape.
• Isolated system: Neither matter nor energy can be exchanged with the surroundings. Example (idealised): a perfectly insulated sealed container, or the universe as a whole.

In practice, truly isolated systems do not exist in a chemistry laboratory — some heat exchange is always inevitable. A sealed thermos flask with a stopper comes close but is not perfect. For IGCSE purposes, a closed system is the relevant category.

Matter vs Energy Exchange

The key distinction between system types hinges on whether the boundary permits the passage of matter, energy, both, or neither.

The Energy Exchange in a Closed System

The fact that energy can cross the boundary of a closed system is essential to understanding Le Chatelier’s Principle. When the temperature of the surroundings changes, heat flows into or out of the closed system, and the equilibrium position shifts in response:

• Heating a closed system at equilibrium: If the forward reaction is exothermic, adding heat shifts the equilibrium to the left (favouring the endothermic reverse reaction). The system absorbs some of the added energy by converting products back to reactants.
• Cooling a closed system at equilibrium: If the forward reaction is exothermic, removing heat shifts the equilibrium to the right (favouring the exothermic forward reaction). The system releases energy to counteract the cooling.

If the system were isolated — with no energy exchange possible — temperature changes could not be applied and no equilibrium shift would occur in response to heating or cooling. This is another reason closed (not isolated) systems are the practical focus in IGCSE chemistry.

Why Closed Systems Are Essential for Equilibrium

The necessity of a closed system for dynamic equilibrium follows directly from the definition of equilibrium itself: forward and reverse reactions must occur simultaneously and at equal rates. If matter can escape, one direction is compromised.

Scenario 1: An Open System — Equilibrium Cannot Form

Take the thermal decomposition of calcium carbonate:

${\text{CaCO}}_3(s)\rightleftharpoons \text{CaO}(s)+{\text{CO}}_2(g)$

If this reaction is carried out in an open crucible (as it often is in a school laboratory):

1. Heat is applied and CaCO₃ begins to decompose into CaO and CO₂.
2. The CO₂ gas, being less dense than air, rises and disperses into the atmosphere.
3. The concentration of CO₂ in the vicinity of the solid CaO is effectively zero.
4. The reverse reaction — CaO + CO₂ → CaCO₃ — has no CO₂ to work with and cannot proceed at any meaningful rate.
5. The forward reaction continues until all CaCO₃ is consumed.
6. Result: The reaction goes to completion. No equilibrium is established.

Scenario 2: A Closed System — Equilibrium Can Form

If the same reaction is carried out in a sealed, strong container:

1. Heat is applied and CaCO₃ decomposes, producing CO₂ gas.
2. The CO₂ is trapped inside the container; its concentration builds up.
3. As more CO₂ accumulates, the rate of the reverse reaction (CaO + CO₂ → CaCO₃) increases.
4. Simultaneously, the forward reaction rate decreases as CaCO₃ is consumed.
5. Eventually, forward and reverse rates equalise.
6. Result: Dynamic equilibrium is established. CaCO₃, CaO, and CO₂ coexist at constant concentrations.

Scenario 3: The Haber Process as a Closed System

${\text{N}}_2(g)+3{\text{H}}_2(g)\rightleftharpoons 2{\text{NH}}_3(g)$

In the industrial Haber process, the reaction vessel is a sealed, pressurised unit. Nitrogen and hydrogen are fed in, and the ammonia produced is continuously removed by cooling and liquefaction. While this is strictly a flow system (an open system in the operational sense), the key point is that within the reactor itself, gases cannot simply escape — the vessel is sealed. Furthermore, the concept of a closed system is what allows us to model and understand the equilibrium behaviour:

• If the vessel were open to the atmosphere, the NH₃ gas would escape, the reverse reaction (2NH₃ → N₂ + 3H₂) would be starved, and the forward reaction would dominate — the equilibrium position could never be established.

Scenario 4: Hydrated Salt Equilibria

For the copper(II) sulfate hydration/dehydration equilibrium:

${\text{CuSO}}_4\cdot 5{\text{H}}_2\text{O}(s)\rightleftharpoons {\text{CuSO}}_4(s)+5{\text{H}}_2\text{O}(g)$

The same principle applies to cobalt(II) chloride:

${\text{CoCl}}_2\cdot 6{\text{H}}_2\text{O}(s)\rightleftharpoons {\text{CoCl}}_2(s)+6{\text{H}}_2\text{O}(g)$

In an open system, heating pink CoCl₂·6H₂O irreversibly produces blue CoCl₂ (used in cobalt chloride paper testing). In a closed system, both pink and blue forms can coexist at equilibrium, and the colour shifts with temperature and humidity.

Examples at IGCSE

1. Sealed Container for the Haber Process

The Haber process reactor operates at approximately 450 °C and 200 atm. The vessel is a sealed, thick-walled steel unit capable of withstanding high pressure. The closed nature of the system ensures that:

• Nitrogen and hydrogen gases cannot leak out before reacting.
• The ammonia produced remains in the reaction mixture long enough for the reverse reaction to occur.
• Unreacted N₂ and H₂ are recycled back into the reactor after ammonia is separated out — this recycling would be impossible if the system were not closed.

2. Sealed Container for the Contact Process

The catalytic oxidation of SO₂ to SO₃ takes place in a sealed converter unit at around 450 °C with a vanadium(V) oxide catalyst. The system is closed to prevent:

• Escape of toxic SO₂ gas into the environment.
• Loss of SO₃ product before it can be absorbed into concentrated H₂SO₄.
• Entry of contaminants that might poison the catalyst (e.g., arsenic compounds).

3. Hydrated CuSO₄ / CoCl₂ Experiments — Open vs Closed

A classic IGCSE practical investigation compares what happens when hydrated salts are heated in open versus closed containers:

This experiment directly demonstrates why a closed system is necessary: in the sealed tube, water vapour is physically trapped and available for the reverse reaction.

4. Contrast: An Open Beaker Where CO₂ Gas Escapes

Consider adding hydrochloric acid to marble chips (calcium carbonate) in an open beaker:

${\text{CaCO}}_3(s)+2\text{HCl}(aq)\to {\text{CaCl}}_2(aq)+{\text{H}}_2\text{O}(l)+{\text{CO}}_2(g)$

This reaction is not reversible under these conditions. The CO₂ gas effervesces and escapes into the atmosphere. Even if the reaction were theoretically reversible, the open beaker prevents any reverse reaction from occurring because one of the products (CO₂) is continuously lost from the system. The reaction proceeds until the limiting reagent is consumed — it goes to completion.

Contrast this with the equilibrium:

${\text{CaCO}}_3(s)\rightleftharpoons \text{CaO}(s)+{\text{CO}}_2(g)$

In a sealed container, the CO₂ cannot escape and equilibrium can be established. This is a fundamentally different type of system to the open-beaker acid-carbonate reaction.

5. The NO₂ / N₂O₄ Equilibrium

$2{\text{NO}}_2(g)\rightleftharpoons {\text{N}}_2{\text{O}}_4(g)$

This equilibrium provides a striking visual demonstration of the closed-system requirement. In a sealed gas syringe:

• At room temperature, the mixture is pale brown (equilibrium mixture of both gases).
• Heating shifts the equilibrium to the left (endothermic direction), producing more brown NO₂ — the colour darkens.
• Cooling shifts it to the right (exothermic direction), producing more colourless N₂O₄ — the colour lightens.

If the syringe were open at the nozzle, both gases would escape and no equilibrium could be observed.

Distinguishing Closed from Isolated Systems — An IGCSE Exam Point

Students often confuse closed and isolated systems. The distinction matters:

• A closed system allows energy exchange. Most chemical reactions involve exothermic or endothermic processes, and the heat released or absorbed flows between the system and surroundings. Temperature-dependent equilibrium shifts (Le Chatelier) are observed precisely because energy can cross the boundary of a closed system.
• An isolated system allows neither matter nor energy exchange. In such a system, temperature changes cannot be externally imposed — any temperature change must arise from the reaction itself. For IGCSE, isolated systems are a theoretical concept; the practical focus is on closed systems.

Exam tip: If a question describes a “sealed glass container that can be heated or cooled,” this is a closed system (matter cannot escape, but energy can enter or leave as heat). If it describes a “perfectly sealed and insulated container with no heat exchange possible,” this is an isolated system.

Sources

• Cambridge IGCSE Chemistry Syllabus 0620, Topic 7: Chemical Reactions — Reversible Reactions and Equilibrium
• OpenStax Chemistry 2e, Chapter 5: Thermochemistry (System and Surroundings) and Chapter 13: Equilibrium Concepts
• BBC Bitesize, GCSE Chemistry — Reversible Reactions
• ChemGuide, Jim Clark — “Chemical Equilibria: An Introduction”
• Royal Society of Chemistry, Education in Chemistry — “Equilibria and Closed Systems”
• IGCSE Chemistry Coursebook (Cambridge University Press), Chapters 1 (States of Matter — System Definitions) and 10 (Equilibrium)

Common Misconceptions

See also: Dynamic Equilibrium, Reversible Reactions, Equilibrium, Le Chatelier’s Principle