Chemical equilibrium — OCR A-Level Chemistry
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Chemical equilibrium explained
Chemical equilibrium covers 9 specification points, including Dynamic equilibrium and le Chatelier’s principle and (a) explanation that a dynamic equilibrium….
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Dynamic equilibrium and le Chatelier’s principle
A reversible reaction reaches dynamic equilibrium in a closed system when the rates of the forward and reverse reactions become equal and the concentrations of reactants and products remain constant, although both reactions continue. Le Chatelier’s principle states that when a system at equilibrium is subjected to a change in conditions, the position of equilibrium shifts to oppose that change. Raising the temperature shifts equilibrium in the endothermic direction; lowering it shifts equilibrium in the exothermic direction. Increasing the pressure of a gaseous system shifts equilibrium towards the side with fewer moles of gas. Increasing the concentration of a reactant shifts equilibrium towards the products. A catalyst does not shift the position of equilibrium. The equilibrium constant Kc depends only on temperature.
(a) explanation that a dynamic equilibrium exists in a closed system when the rate of the forward reaction is equal to the rate of the reverse reaction and the concentrations of reactants and products do not change
A dynamic equilibrium is reached only in a closed system, meaning no reactant or product can enter or leave. At this point the forward and reverse reactions continue at the same time, but the rate of the forward reaction equals the rate of the reverse reaction. Because the two rates match, the concentrations of reactants and products stay constant, even though particles are still reacting in both directions. The word dynamic signals that both reactions keep occurring; the word equilibrium signals the balance of rates. For example, in a sealed vessel where hydrogen and iodine form hydrogen iodide, the forward and reverse rates become equal and the amounts of each substance stop changing, though the reactions have not stopped.
(b) le Chatelier’s principle and its application for homogeneous equilibria to deduce qualitatively the effect of a change in temperature, pressure or concentration on the position of equilibrium
Le Chatelier's principle states that when a system at equilibrium is disturbed, the position of equilibrium shifts to oppose the change. Apply it only to homogeneous equilibria, where all species are in the same phase. For a temperature change, heating shifts equilibrium in the endothermic direction and cooling shifts it in the exothermic direction. For a pressure change, increasing pressure shifts equilibrium towards the side with fewer moles of gas, and decreasing pressure shifts it towards the side with more moles of gas. For a concentration change, adding a reactant shifts equilibrium to the right to use it up, while removing a product shifts it to the right to replace it. Deduce the direction qualitatively rather than calculating values.
(c) explanation that a catalyst increases the rate of both forward and reverse reactions in an equilibrium by the same amount resulting in an unchanged position of equilibrium
A catalyst provides an alternative reaction pathway of lower activation energy. In a system at equilibrium the same pathway is available to the forward and the reverse reaction, so both rates rise by the same factor and the two rates remain equal to each other. Because the rates stay equal, the concentrations of reactants and products at equilibrium do not change, so the position of equilibrium is unchanged. The catalyst therefore does not alter the equilibrium yield or the value of the equilibrium constant; it only shortens the time taken to reach equilibrium. For example, adding iron to the Haber process speeds the approach to equilibrium but does not shift the balance between nitrogen, hydrogen and ammonia at a given temperature and pressure.
(d) the techniques and procedures used to investigate changes to the position of equilibrium for changes in concentration and temperature.
Investigating equilibrium shifts requires a reversible system with a visible or measurable change, such as the colour of an indicator or of a transition-metal complex ion. To test a change in concentration, prepare the equilibrium mixture, record the initial appearance or reading, then add a small volume of a solution containing one of the species and observe or measure the new position after mixing. To test a change in temperature, place portions of the same mixture in water baths at different temperatures, allow each to reach thermal equilibrium, and compare the colour or reading. Control variables such as total volume, concentrations of other species and observation time so that only the chosen factor changes. Interpret each result using Le Chatelier's principle, identifying the direction of the shift.
(e) explanation of the importance to the chemical industry of a compromise between chemical equilibrium and reaction rate in deciding the operational conditions
Industrial processes such as the Haber process must balance two competing demands. A low temperature shifts the equilibrium position towards the products because the forward reaction is exothermic, raising the equilibrium yield, but it also slows both the forward and reverse reactions, so equilibrium is reached very slowly. A high pressure shifts this equilibrium towards the products because there are fewer gas moles on the product side, and it also increases the rate by raising gas concentration, but high pressure demands costly, strong plant and carries safety risks. Industry therefore chooses a compromise temperature and pressure that gives an acceptable yield at an acceptable rate, often with a catalyst to speed attainment of equilibrium without changing the position. The chosen conditions maximise economic return, not theoretical yield.
The equilibrium constant, Kc
For a homogeneous equilibrium aA + bB ⇌ cC + dD, the equilibrium constant Kc is defined by Kc = [C]^c[D]^d / [A]^a[B]^b, where each concentration is the equilibrium concentration in mol dm⁻³ and each exponent is the stoichiometric coefficient from the balanced equation. Pure solids and pure liquids are omitted because their concentration is effectively constant. Kc depends only on temperature; changing concentration, pressure or adding a catalyst does not change its value at a fixed temperature. A large Kc indicates the equilibrium lies towards the products, and a small Kc indicates it lies towards the reactants.
(f) expressions for the equilibrium constant, Kc , for homogeneous reactions and calculations of the equilibrium constant, Kc , from provided equilibrium concentrations
For a homogeneous equilibrium all species share one phase, so their concentrations can appear together in a single Kc expression. Take the balanced equation, for example N₂(g) + 3H₂(g) ⇌ 2NH₃(g). Write products over reactants, each concentration raised to its stoichiometric coefficient: Kc = [NH₃]² / ([N₂][H₂]³). Pure solids and pure liquids are omitted because their concentration is effectively constant. To calculate Kc, substitute the equilibrium concentrations given, keeping units consistent (mol dm⁻³), then evaluate. Units of Kc follow from the powers: here (mol dm⁻³)² / (mol dm⁻³ × (mol dm⁻³)³) = mol⁻² dm⁶. Check the arithmetic and the power on each species before selecting an answer.
(g) estimation of the position of equilibrium from the magnitude of Kc.
The size of Kc indicates where a homogeneous equilibrium lies. A large Kc, well above 1, means the products dominate at equilibrium, so the position lies to the right. A small Kc, well below 1, means reactants dominate, so the position lies to the left. A Kc close to 1 means appreciable amounts of both reactants and products are present. Because Kc depends only on temperature for a given reaction, the magnitude is compared at a stated temperature. For example, if Kc for an esterification is 4, both sides are significantly populated; if Kc is 1 × 10⁻⁵, the equilibrium lies strongly towards the reactants. Estimate the position by comparing the value with 1 and describing which side predominates, rather than claiming the reaction goes to completion.
Your focus
- Describe the features of dynamic equilibrium in a closed system.
- Apply Le Chatelier’s principle to predict the direction of shift when temperature, pressure or concentration changes.
- Explain why a catalyst does not alter the position of equilibrium.
Show all 27 objectives
- Explain why a dynamic equilibrium needs a closed system
- Describe the balance of forward and reverse rates at equilibrium
- State that reactant and product concentrations stay constant at equilibrium
- State Le Chatelier's principle and apply it to homogeneous equilibria
- Deduce the direction of shift caused by a temperature, pressure or concentration change
- Explain why a catalyst does not change the position of equilibrium
- Explain how a catalyst lowers activation energy for both forward and reverse reactions.
- Describe why equal increases in forward and reverse rates leave the position of equilibrium unchanged.
- Distinguish between the rate of reaching equilibrium and the position of equilibrium.
- Describe a procedure to investigate the effect of concentration change on equilibrium position.
- Describe a procedure to investigate the effect of temperature change on equilibrium position.
- Interpret observed shifts using Le Chatelier's principle while controlling variables.
- Explain how temperature and pressure affect both the equilibrium position and the rate of reaction
- Justify a chosen compromise of temperature and pressure for a named industrial process
- Evaluate the economic and safety factors that shape operational conditions in the chemical industry
- Construct the Kc expression for a given homogeneous equilibrium
- Explain how temperature and other changes affect the value of Kc
- Interpret the size of Kc in terms of the position of equilibrium
- Construct a Kc expression for a given homogeneous equilibrium.
- Explain why pure solids and liquids are excluded from Kc.
- Calculate Kc from supplied equilibrium concentrations and state its units.
- Interpret the magnitude of Kc to estimate the position of equilibrium.
- Distinguish the position of equilibrium from reaction rate.
- Explain that Kc depends on temperature for a given reaction.
Chemical equilibrium exam tips
Marking Points
- Defines dynamic equilibrium as the state in a closed system where forward and reverse reaction rates are equal and concentrations remain constant while both reactions continue.
- States Le Chatelier’s principle as the shift of equilibrium position to oppose a change in conditions.
- Explains that increasing temperature shifts equilibrium in the endothermic direction and decreasing temperature shifts it in the exothermic direction.
- Explains that increasing pressure shifts equilibrium towards the side with fewer moles of gas.
- Explains that increasing the concentration of a reactant shifts equilibrium towards the products.
- States that a catalyst does not change the position of equilibrium and that Kc depends only on temperature.
- States that equilibrium requires a closed system so no substances enter or leave
- Explains that the forward and reverse reactions both continue at equilibrium
- Identifies that the rate of the forward reaction equals the rate of the reverse reaction
- States that the concentrations of reactants and products remain constant
- Distinguishes dynamic equilibrium from a reaction that has simply stopped
- States that the position of equilibrium shifts to oppose the change applied
- Deduces that raising temperature shifts equilibrium in the endothermic direction
- Deduces that increasing pressure shifts equilibrium towards the side with fewer moles of gas
- Deduces that adding a reactant or removing a product shifts equilibrium to the right
- Applies the principle only to homogeneous equilibria where all species share one phase
- States that a catalyst lowers the activation energy by providing an alternative reaction pathway.
- Explains that the same lowered activation energy applies to both the forward and the reverse reaction.
- Concludes that both forward and reverse rates increase by the same amount so the rates remain equal.
- States that the position of equilibrium and the equilibrium yield are unchanged by a catalyst.
- Recognises that a catalyst only reduces the time taken to reach equilibrium.
- Selects a reversible system that shows a visible or measurable change in position of equilibrium.
- Records an initial observation or measurement before changing concentration or temperature.
- Adds a solution containing one species to change concentration while keeping other variables controlled.
- Uses water baths at different temperatures to investigate the effect of temperature change.
- Compares results and interprets each shift using Le Chatelier's principle.
- Explains that a low temperature increases the equilibrium yield for an exothermic forward reaction but decreases the rate at which equilibrium is reached
- Explains that a high pressure increases both the equilibrium yield, where the product side has fewer gas moles, and the reaction rate, but raises plant cost and safety demands
- Identifies a compromise temperature and pressure that balances an acceptable yield against an acceptable rate of attainment of equilibrium
- Explains that a catalyst increases the rate of attainment of equilibrium without altering the equilibrium position or the value of Kc
- Links the chosen operational conditions to economic factors such as energy cost, plant cost and throughput rather than to maximum theoretical yield alone
- Writes the Kc expression with product concentrations in the numerator and reactant concentrations in the denominator, each raised to its stoichiometric coefficient
- States that concentrations in the Kc expression are equilibrium concentrations measured in mol dm⁻³
- Explains that the value of Kc depends only on temperature and is unchanged by concentration, pressure or a catalyst
- Omits pure solids and pure liquids from the Kc expression because their concentration is effectively constant
- Interprets a large Kc as an equilibrium position towards the products and a small Kc as one towards the reactants
- Writes the product concentration raised to its stoichiometric coefficient over the reactant concentrations each raised to their coefficients.
- Omits pure solids and pure liquids from the Kc expression because their concentration is constant.
- Substitutes the supplied equilibrium concentrations into the expression and evaluates Kc correctly.
- Derives the units of Kc from the powers in the expression, for example mol⁻² dm⁶ for the ammonia synthesis.
- Distinguishes a homogeneous equilibrium, where all species share one phase, from a heterogeneous one.
- States that a large Kc, much greater than 1, places the equilibrium towards the products.
- States that a small Kc, much less than 1, places the equilibrium towards the reactants.
- Recognises that a Kc near 1 indicates comparable amounts of reactants and products.
- Links the magnitude of Kc to the relative amounts of species at equilibrium rather than to reaction rate.
- Notes that Kc varies with temperature for a given reaction.
Examiner Tips
- 💡Check whether the system is closed and whether all species are gaseous before applying pressure reasoning.
- 💡Count the moles of gas on each side before deciding the direction of a pressure-induced shift.
- 💡Link temperature changes to the endothermic or exothermic direction rather than to the sign of enthalpy alone.
- 💡Remember that Kc changes only with temperature, not with pressure or concentration changes.
- 💡Check that the option you choose mentions both equal rates and constant concentrations, since either alone is incomplete
- 💡Reject any option that says the reaction has stopped or that amounts are equal
- 💡Look for the phrase closed system, because equilibrium cannot be maintained if substances escape
- 💡Count moles of gas on each side before deciding the direction of a pressure shift
- 💡Match the temperature change to the endothermic or exothermic direction using the enthalpy change sign
- 💡Check that the option names the correct direction of shift rather than only describing the disturbance
- 💡Link the answer explicitly to activation energy and to both reaction directions rather than saying only that a catalyst speeds up a reaction.
- 💡Use the phrase 'by the same amount' to show that the two rates stay equal.
- 💡Distinguish clearly between rate of reaching equilibrium and position of equilibrium.
- 💡Check that any statement about yield or equilibrium constant is consistent with an unchanged position.
- 💡Name the reversible system and the observable change you will monitor.
- 💡State the variable you change and list the variables you keep constant.
- 💡Describe the procedure in ordered steps, including mixing and equilibration.
- 💡Interpret each result with Le Chatelier's principle rather than only describing colour changes.
- 💡Name a real industrial process, such as the Haber process, and quote its approximate operating temperature and pressure to anchor the compromise argument
- 💡Use Le Chatelier's principle to justify the equilibrium shift and collision theory to justify the rate change, keeping the two arguments distinct
- 💡Conclude explicitly that the compromise maximises economic return, since examiners reward the link between chemistry and industrial decision-making
- 💡Balance the equation first, then transfer each coefficient directly into the exponent of the Kc expression
- 💡Check that every species included is aqueous or gaseous, and delete any solid or liquid before finalising the expression
- 💡When interpreting a value, compare it with 1 to decide whether products or reactants dominate at equilibrium
- 💡Write the balanced equation first, then build the expression term by term from the coefficients.
- 💡Check each exponent against the coefficient in the equation before substituting numbers.
- 💡Work out the units separately from the numerical value so an error in one does not hide an error in the other.
- 💡Compare the numerical Kc with 1 to decide which side predominates.
- 💡Use the words 'predominates' or 'lies to the left/right' rather than 'complete'.
- 💡Remember Kc is temperature-dependent, so quote the temperature when comparing values.
Common Mistakes
- States that at equilibrium the forward and reverse reactions have stopped; correct this by stating that both reactions continue at equal rates.
- Claims that a catalyst shifts the equilibrium position; correct this by stating that a catalyst only speeds the attainment of equilibrium.
- Applies pressure changes to systems with equal moles of gas on both sides and expects a shift; correct this by noting that no shift occurs when the number of gaseous moles is equal.
- Associates a temperature increase with the exothermic direction; correct this by linking a temperature rise to the endothermic direction.
- Claims that the reaction stops at equilibrium; correct this by stating that both reactions continue at equal rates
- States that concentrations of reactants and products become equal; correct this by saying the concentrations stay constant, not necessarily equal
- Describes equilibrium as possible in an open system; correct this by requiring a closed system
- Says only the forward reaction occurs; correct this by naming both forward and reverse reactions running together
- Assumes a catalyst shifts the position of equilibrium; correct this by stating a catalyst only speeds up attainment of equilibrium
- Applies pressure reasoning to equilibria with equal moles of gas on both sides; correct this by noting no shift occurs in that case
- Reverses the temperature rule; correct this by linking heating to the endothermic direction
- Applies the principle to a heterogeneous mixture as if all species were in one phase; correct this by restricting it to homogeneous equilibria
- Claims that a catalyst shifts the equilibrium towards the products; correct this by stating that the position of equilibrium is unchanged.
- Assumes a catalyst speeds only the forward reaction; correct this by noting that both directions are accelerated equally.
- States that a catalyst changes the value of the equilibrium constant; correct this by explaining that only temperature changes that value.
- Believes a catalyst increases the equilibrium yield; correct this by explaining that yield at equilibrium is unaffected and only the rate of approach changes.
- Changes more than one variable at once; correct this by altering only concentration or only temperature in each test.
- Omits a control or initial reading; correct this by recording the starting appearance or measurement before the change.
- Fails to allow the mixture to reach the new temperature; correct this by waiting until thermal equilibrium is established.
- Describes observations without interpreting them; correct this by linking each change to the direction of equilibrium shift.
- States that a catalyst increases the equilibrium yield; correct this by explaining that a catalyst speeds attainment of equilibrium but leaves the position and Kc unchanged
- Claims that the lowest possible temperature is always chosen because it gives the highest yield; correct this by explaining that the rate would be too slow and the process uneconomic
- Assumes that high pressure is always used because it raises yield; correct this by explaining that plant strength, cost and safety limit the pressure chosen
- Treats equilibrium yield and reaction rate as independent; correct this by explaining that changing temperature or pressure usually affects both in opposing ways
- Raises concentrations to the wrong powers; correct this by using the stoichiometric coefficients from the balanced equation as exponents
- Includes a pure solid or pure liquid in the expression; correct this by omitting species whose concentration is effectively constant
- States that adding a catalyst changes Kc; correct this by explaining that a catalyst changes only the rate of attainment, not the value of Kc
- Confuses Kc with the rate constant; correct this by defining Kc as a ratio of equilibrium concentrations at a given temperature
- Inverts the expression by placing reactants over products; write products over reactants.
- Raises concentrations to the wrong power by ignoring stoichiometric coefficients; apply each coefficient as the exponent.
- Includes a solid or liquid in the expression; omit pure solids and liquids.
- Uses initial rather than equilibrium concentrations; substitute the equilibrium values supplied.
- Treats a large Kc as a fast reaction; relate Kc to position, not rate.
- Claims a large Kc means the reaction goes to completion; describe products as predominating at equilibrium.
- Compares Kc values measured at different temperatures as if directly comparable; compare at the same temperature.
- Confuses the direction of the shift with the size of Kc; state which side predominates.