How far? — OCR A-Level Chemistry
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How far? explained
How far? covers 10 specification points, including Equilibrium and (a) use of the terms mole fraction and partial….
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Equilibrium
This section introduces chemical equilibrium as a dynamic state in which the forward and reverse reactions proceed at equal rates in a closed system, so concentrations remain constant though not equal. Read it alongside the earlier rates work: equilibrium is reached, not static, and the position can be shifted by changing concentration, pressure or temperature. Link it to the equilibrium constant Kc, its expression in terms of concentrations, and how its value depends only on temperature. Use the reading to build the idea that a catalyst speeds attainment of equilibrium without changing the position or the value of Kc.
(a) use of the terms mole fraction and partial pressure
Mole fraction and partial pressure describe the composition of a gaseous equilibrium mixture. The mole fraction of a component is the number of moles of that component divided by the total number of moles of all components present; the fractions in a mixture sum to 1. Partial pressure is the pressure a gas would exert if it alone occupied the container, and for an ideal gas it equals its mole fraction multiplied by the total pressure: p(A) = x(A) × p(total). For the equilibrium aA(g) + bB(g) ⇌ cC(g) + dD(g), the equilibrium constant Kp is written using partial pressures, each raised to its stoichiometric coefficient. Worked example: a mixture contains 2 mol of A and 6 mol of B, total pressure 200 kPa; x(A) = 2/8 = 0.25, so p(A) = 0.25 × 200 kPa = 50 kPa, and p(B) = 0.75 × 200 kPa = 150 kPa.
(b) calculation of quantities present at equilibrium, given appropriate data
Equilibrium quantities are found by tabulating initial moles, the change in moles and the moles at equilibrium, using the stoichiometry of the balanced equation. For H₂(g) + I₂(g) ⇌ 2HI(g), starting with 1.0 mol H₂ and 1.0 mol I₂ in a vessel, if 0.60 mol of H₂ reacts then 0.60 mol of I₂ also reacts and 2 × 0.60 = 1.20 mol of HI forms, leaving 0.40 mol H₂, 0.40 mol I₂ and 1.20 mol HI, a total of 2.00 mol. Mole fractions follow by dividing each equilibrium amount by the total, and partial pressures by multiplying each mole fraction by the total pressure. The same table supports calculation of Kc from concentrations (moles divided by volume) or Kp from partial pressures.
(c) the techniques and procedures used to determine quantities present at equilibrium
To find how much of each substance is present once a homogeneous or heterogeneous system has reached equilibrium, you freeze the composition and then measure it. A common procedure is titration: quench a measured sample so the reaction stops, then titrate the quenched portion against a standard solution and use the result to calculate the amount of one species. From the balanced equation and the known starting amounts you then work out the amounts of the other species by stoichiometry, remembering that solids are not included in the amounts used for Kc. Techniques must be chosen so that sampling does not disturb the equilibrium: withdraw a small aliquot, stop the reaction quickly, and keep temperature constant because Kc depends on temperature. Repeating the measurement at intervals until the titre is constant confirms equilibrium has been reached.
(d) expressions for Kc and Kp for homogeneous and heterogeneous equilibria (see also 3.2.3 f)
An equilibrium constant expression is built from the balanced equation. For Kc, each concentration is raised to the power of its stoichiometric coefficient, products over reactants. For Kp, each partial pressure is used in the same way. In a homogeneous equilibrium all species are in the same phase, so all appear in the expression. In a heterogeneous equilibrium a pure solid or pure liquid has constant activity and is omitted, so only gaseous or dissolved species appear. For example, for CaCO₃(s) ⇌ CaO(s) + CO₂(g), Kc = [CO₂] and Kp = p(CO₂). Partial pressure is calculated as mole fraction multiplied by total pressure, and units of Kc and Kp follow from the powers used.
(e) calculations of Kc and Kp, or related quantities, including determination of units (see also 3.2.3 f)
This criterion covers numerical work with equilibrium constants. For Kc, write the expression as products over reactants with each concentration raised to its stoichiometric coefficient, substitute equilibrium concentrations in mol dm⁻³, and evaluate. For Kp, use partial pressures in Pa or kPa, each raised to its coefficient, and evaluate. Determine units by cancelling the units of the numerator against the denominator; for example, for H₂(g) + I₂(g) ⇌ 2HI(g) the units cancel completely, so Kc is dimensionless, whereas for N₂O₄(g) ⇌ 2NO₂(g) Kc has units of mol dm⁻³. Related quantities include calculating an equilibrium concentration or partial pressure from a known K value, and converting between mole fraction and partial pressure using p = mole fraction × total pressure. Always check the arithmetic and the direction of any rearrangement.
(f) (i) the qualitative effect on equilibrium constants of changing temperature for exothermic and endothermic reactions
This criterion concerns how temperature changes affect the value of an equilibrium constant, treated qualitatively rather than numerically. For an exothermic forward reaction, raising the temperature shifts the position of equilibrium to the left, so the equilibrium constant decreases; lowering the temperature increases it. For an endothermic forward reaction, raising the temperature shifts equilibrium to the right, so the equilibrium constant increases; lowering the temperature decreases it. The reasoning links to Le Chatelier's principle: the system opposes the temperature change by favouring the direction that absorbs or releases heat. Note that a catalyst changes the rate of attainment but does not change the value of the equilibrium constant, and pressure changes do not alter K for a given reaction at constant temperature.
(ii) the constancy of equilibrium constants with changes in concentration, pressure or in the presence of a catalyst
The equilibrium constant Kc depends only on temperature for a given reaction. Changing the concentration of a species shifts the position of equilibrium, but the ratio of product concentrations to reactant concentrations, each raised to its stoichiometric power, returns to the same value once equilibrium is re-established. Changing pressure by compressing a gaseous system also shifts the position of equilibrium, yet Kc is unchanged because the ratio is unaffected once new equilibrium concentrations are reached. Adding a catalyst speeds the forward and reverse reactions equally, so equilibrium is reached sooner but Kc stays the same. Only a change in temperature alters Kc, because it changes the value of ΔG° and hence the position of equilibrium.
(g) explanation of how an equilibrium constant controls the position of equilibrium on changing concentration, pressure and temperature
The equilibrium constant expresses the ratio of products to reactants at equilibrium, and the system adjusts the position of equilibrium so that this ratio is restored whenever it is disturbed. Increasing the concentration of a reactant makes the reaction quotient smaller than Kc, so the forward reaction is favoured until the ratio equals Kc again. Raising the pressure of a gaseous system by reducing volume favours the side with fewer gas molecules, shifting position while Kc stays fixed at constant temperature. Changing temperature alters the value of Kc itself: for an exothermic forward reaction, raising temperature lowers Kc and shifts equilibrium towards reactants, whereas for an endothermic reaction the opposite occurs.
(h) application of the above principles in 5.1.2 How far? for Kc, Kp to other equilibrium constants, where appropriate (see also 5.1.3 c etc.).
This criterion asks you to transfer the equilibrium-constant principles met for Kc and Kp to other equilibrium constants where appropriate. The core principles are: write the expression from the balanced equation, raise each concentration or partial pressure to its stoichiometric coefficient, and place products over reactants. For Kc you use equilibrium concentrations in mol dm⁻³; for Kp you use equilibrium partial pressures in Pa. The same logic extends to constants such as Ka for weak acids, where Ka = [H⁺][A⁻]/[HA], and to Kw = [H⁺][OH⁻]. Recognise that a constant's form follows its defining equation, that its value depends on temperature only, and that units depend on the powers in the expression.
Your focus
- Explain equilibrium as a dynamic state with equal forward and reverse rates.
- Write and interpret an expression for Kc.
- Describe how changing conditions affects the position of equilibrium and the value of Kc.
Show all 31 objectives
- Define mole fraction and partial pressure accurately
- Calculate the partial pressure of a component from its mole fraction and the total pressure
- Construct a Kp expression using partial pressures with correct stoichiometric powers
- Construct an ICE table for a gaseous equilibrium from given initial data
- Calculate equilibrium moles using the stoichiometry of the balanced equation
- Determine Kc or Kp from equilibrium quantities with correct units
- Describe a procedure that determines the amount of a species present at equilibrium.
- Convert a measured amount of one species into amounts of the others using the balanced equation.
- Explain how constant titre and constant temperature are used to confirm equilibrium.
- Construct Kc and Kp expressions from a balanced equation.
- Distinguish homogeneous from heterogeneous equilibria and omit pure solids and liquids.
- Calculate partial pressures and determine the units of an equilibrium constant.
- Construct Kc and Kp expressions from a balanced equation.
- Calculate a value of Kc or Kp from equilibrium data and state its units.
- Use a known equilibrium constant to determine an unknown concentration or partial pressure.
- Predict qualitatively how changing temperature affects the equilibrium constant for exothermic and endothermic reactions.
- Explain the predicted effect using Le Chatelier's principle.
- Distinguish the effect of temperature from that of catalysts and pressure on the equilibrium constant.
- Explain why Kc is constant when concentration, pressure or a catalyst is changed.
- Distinguish between a shift in equilibrium position and a change in the value of Kc.
- Identify temperature as the only factor that changes Kc for a given reaction.
- Explain how changes in concentration and pressure shift equilibrium while Kc stays constant
- Explain how temperature changes the value of Kc and the position of equilibrium
- Apply the reaction quotient compared with Kc to predict the direction of shift
- Construct an equilibrium expression for a given balanced equation, applying stoichiometric powers.
- Distinguish the concentration basis of Kc from the partial-pressure basis of Kp.
- Apply the same principles to write Ka and Kw expressions.
- Explain why an equilibrium constant's value depends only on temperature.
How far? exam tips
Marking Points
- Defines mole fraction as the number of moles of a component divided by the total number of moles of all components in the mixture
- States that the mole fractions of all components in a mixture add up to 1
- Defines partial pressure as the pressure a gas would exert if it alone occupied the container, equal to its mole fraction times the total pressure
- Calculates a partial pressure by multiplying the mole fraction by the total pressure, for example 0.25 × 200 kPa = 50 kPa
- Applies partial pressures in the Kp expression, raising each to its stoichiometric coefficient from the balanced equation
- Sets out initial moles, change in moles and equilibrium moles in a table consistent with the balanced equation
- Applies the stoichiometric ratio so that the change in each species is in proportion to its coefficient in the equation
- Calculates the total number of moles at equilibrium by adding the equilibrium amounts of all species
- Converts equilibrium moles to concentration by dividing by the volume, or to partial pressure by multiplying mole fraction by total pressure
- Substitutes the equilibrium values into the correct Kc or Kp expression to obtain the required quantity
- Identifies that a sample must be quenched or otherwise stopped before titration so the equilibrium composition is fixed.
- Uses the balanced equation and known initial amounts to convert the measured amount of one species into the amounts of the others by stoichiometry.
- Explains that the titre becoming constant on repeated sampling shows the system has reached equilibrium.
- States that temperature must be held constant because the value of Kc depends on temperature.
- Writes products over reactants with each concentration or partial pressure raised to its stoichiometric coefficient.
- Omits pure solids and pure liquids from heterogeneous expressions because their activity is constant.
- Calculates a partial pressure as mole fraction multiplied by total pressure before substituting into Kp.
- Derives the units of Kc or Kp from the powers in the expression.
- Writes the Kc expression with product concentrations in the numerator and reactant concentrations in the denominator, each raised to its stoichiometric coefficient.
- Substitutes equilibrium concentrations in mol dm⁻³ into the Kc expression and evaluates the numerical value correctly.
- Writes the Kp expression using partial pressures and raises each partial pressure to its stoichiometric coefficient.
- Determines the units of Kc or Kp by cancelling the units of the numerator against those of the denominator.
- Calculates a partial pressure from mole fraction and total pressure using p = mole fraction × total pressure.
- Rearranges the equilibrium expression to find an unknown equilibrium concentration or partial pressure from a known K value.
- States that raising the temperature decreases the equilibrium constant for an exothermic forward reaction.
- States that raising the temperature increases the equilibrium constant for an endothermic forward reaction.
- Explains the effect by linking the shift in equilibrium position to Le Chatelier's principle opposing the temperature change.
- Recognises that a catalyst does not change the value of the equilibrium constant.
- States that Kc depends only on temperature for a given reaction.
- Explains that changing concentration shifts the equilibrium position but leaves Kc unchanged once equilibrium is re-established.
- Explains that changing pressure shifts the position of a gaseous equilibrium but leaves Kc unchanged.
- Explains that a catalyst speeds both forward and reverse reactions equally, so equilibrium is reached sooner but Kc is unchanged.
- Identifies temperature as the only factor among those listed that changes the value of Kc.
- Explains that the reaction quotient is compared with Kc to decide the direction of shift
- States that increasing reactant concentration shifts equilibrium forward to restore the Kc ratio
- Explains that a pressure increase favours the side with fewer moles of gas at constant Kc
- States that changing temperature changes the value of Kc and so moves the position of equilibrium
- Writes the equilibrium expression with product terms in the numerator and reactant terms in the denominator, each raised to its stoichiometric coefficient.
- Uses equilibrium concentrations in mol dm⁻³ for Kc and equilibrium partial pressures in Pa for Kp, keeping the two conventions distinct.
- Transfers the same construction to Ka for a weak acid, writing Ka = [H⁺][A⁻]/[HA] from the dissociation equation.
- States that the numerical value of an equilibrium constant changes with temperature and is unaffected by concentration or pressure changes at constant temperature.
- Determines the units of a constant by cancelling the concentration or pressure units raised to the powers in the expression.
Examiner Tips
- 💡Read the section with the Kc expression beside it so the link between the dynamic model and the constant is clear.
- 💡Note which factors change the position of equilibrium and which change only the rate.
- 💡Check that any worked example uses a closed system, since equilibrium requires no exchange with surroundings.
- 💡Check that your mole fractions sum to 1 before using them, which catches arithmetic slips early
- 💡Keep units consistent, converting all pressures to the same unit such as kPa before substituting into Kp
- 💡Write the Kp expression from the balanced equation first, then substitute partial pressures, so powers are not forgotten
- 💡Draw the ICE table before any calculation so every quantity has a clear place
- 💡Check the total moles at equilibrium by adding the final column, which exposes arithmetic errors
- 💡State units for Kc or Kp where required, and check whether the question asks for moles, concentration or partial pressure
- 💡Link each procedural step to why it is done, not just what is done, so the technique is justified.
- 💡Show the stoichiometric conversion explicitly from the measured species to the others.
- 💡Keep temperature fixed throughout and say so when explaining why Kc is constant.
- 💡Write the balanced equation first, then build the expression term by term.
- 💡Check each species' phase and delete pure solids and liquids before finalising.
- 💡Work out units by cancelling the powers in the expression.
- 💡Write the balanced equation first, then build the expression directly from its coefficients.
- 💡Show the unit cancellation explicitly so the final units are clearly justified.
- 💡Check whether the question gives initial or equilibrium amounts and convert to equilibrium values before substituting.
- 💡Identify whether the forward reaction is exothermic or endothermic before predicting the effect on K.
- 💡Link each prediction explicitly to the direction of the equilibrium shift.
- 💡Remember that only temperature changes the value of the equilibrium constant for a given reaction.
- 💡Link each factor to whether it changes the equilibrium position and whether it changes the value of Kc, treating these as separate questions.
- 💡Use the expression for Kc to show why concentration and pressure changes cancel out at the new equilibrium.
- 💡State explicitly that temperature is the only variable that alters Kc, and be ready to explain why in terms of ΔG°.
- 💡Compare the reaction quotient with Kc to justify the direction of any shift
- 💡Count moles of gas on each side before predicting the effect of a pressure change
- 💡Link temperature effects to whether the forward reaction is exothermic or endothermic
- 💡Write the balanced equation first, then build the expression term by term from the coefficients.
- 💡Check units by cancelling powers of mol dm⁻³ or Pa before selecting an answer.
- 💡For Ka and Kw questions, derive the expression from the defining dissociation or self-ionisation equation rather than memorising blindly.
- 💡Eliminate options that change a constant's value with concentration or pressure at fixed temperature.
Common Mistakes
- Describes equilibrium as a state where reactions stop; recognise it as dynamic with forward and reverse rates equal.
- Assumes concentrations of reactants and products are equal at equilibrium; understand they are constant but generally unequal.
- Believes a catalyst shifts the equilibrium position; note it only speeds the approach to equilibrium.
- Thinks Kc changes with concentration or pressure; recognise Kc depends only on temperature.
- Divides by the number of components rather than by the total number of moles; divide the moles of the component by the total moles of all components
- Treats partial pressure as a fixed fraction of 100 kPa regardless of the stated total pressure; multiply the mole fraction by the actual total pressure given
- Omits the stoichiometric powers when writing Kp; raise each partial pressure to the coefficient shown in the balanced equation
- Applies the change in moles equally to every species regardless of coefficient; scale each change by the coefficient in the balanced equation
- Forgets to subtract the reacted amount from the initial amount; subtract the change from the initial moles to obtain equilibrium moles
- Uses initial rather than equilibrium concentrations in the Kc expression; substitute the equilibrium values after the table is complete
- Titrates the equilibrium mixture directly without quenching; stop the reaction first so the measured amount reflects the equilibrium composition.
- Includes the amount of a pure solid when calculating Kc; omit solids and use only species present in the expression.
- Assumes the titre is constant from the first sample; repeat sampling until successive titres agree before treating the system as at equilibrium.
- Includes a pure solid such as CaCO₃ in the expression; omit pure solids and liquids and keep only gaseous or dissolved species.
- Raises concentrations to the wrong power; use the stoichiometric coefficient from the balanced equation.
- Uses total pressure instead of partial pressure in Kp; multiply mole fraction by total pressure first.
- Places reactants in the numerator and products in the denominator; write products over reactants.
- Omits the stoichiometric powers in the equilibrium expression; raise each species to its coefficient.
- Uses initial concentrations instead of equilibrium concentrations in the Kc expression; substitute equilibrium values.
- States that Kc always has units; cancel the units and recognise when the expression is dimensionless.
- Claims that raising temperature always increases the equilibrium constant; distinguish exothermic from endothermic forward reactions.
- States that a catalyst increases the equilibrium constant; recognise that a catalyst only speeds attainment of equilibrium.
- Confuses the effect of pressure with the effect of temperature on K; note that K depends on temperature for a given reaction.
- Assumes the equilibrium constant changes when concentration changes; recognise that K is constant at a fixed temperature.
- Claims that adding a catalyst increases Kc; correct this by stating that a catalyst speeds both directions equally and leaves Kc unchanged.
- Assumes that changing concentration changes Kc; correct this by explaining that the ratio returns to the same value at the new equilibrium position.
- States that changing pressure changes Kc; correct this by explaining that compression shifts position but the equilibrium ratio is unchanged.
- Believes Kc is fixed regardless of temperature; correct this by stating that temperature is the one factor that does change Kc.
- Assumes pressure change alters Kc; recognise that pressure shifts position while Kc stays constant at fixed temperature
- Treats temperature change like a concentration change; understand that temperature changes the value of Kc itself
- Ignores the mole ratio of gases when predicting pressure effects; compare total moles of gas on each side before deciding the shift
- Places reactants over products in the expression; write products in the numerator and reactants in the denominator.
- Uses initial concentrations instead of equilibrium concentrations in a Kc expression; substitute the equilibrium values.
- Omits the power on a species whose coefficient is greater than one; raise each term to its stoichiometric coefficient.
- Claims the value of Kc changes when concentration is altered at constant temperature; state that only temperature changes the value.