7. Equilibria

A reversible reaction can form products from reactants and also regenerate reactants from products. It is represented with a double arrow when both directions are relevant.

The forward and reverse reactions may have different rates. In a closed system, the composition can settle at a dynamic equilibrium where both continue.

N₂O₄(g) ⇌ 2NO₂(g) allows colourless N₂O₄ to form brown NO₂ and brown NO₂ to recombine. The observed colour depends on the equilibrium composition.

Reversible does not mean the reaction stops or that the amounts become equal. It means both directions occur under the conditions.

When a system at equilibrium is disturbed, it shifts in the direction that opposes the disturbance, partially restoring the balance.

For concentration changes, the system consumes an added reactant or replaces a removed product. For pressure changes in gases, it favours the side with fewer gas molecules; temperature must be treated as heat in the equation.

For N₂ + 3H₂ ⇌ 2NH₃, adding H₂ shifts right, while increasing pressure also shifts right because four gas moles become two.

A catalyst changes how quickly equilibrium is reached, not the equilibrium position. “Opposes” does not mean the disturbance is completely cancelled.

A change in concentration, pressure or temperature disturbs an equilibrium. The system responds in the direction that reduces the effect of that change.

Write the balanced equation first. For gases, compare total gas moles; for temperature, treat heat as a reactant in an endothermic direction or a product in an exothermic direction.

For 2SO₂ + O₂ ⇌ 2SO₃, adding oxygen shifts right and increasing pressure shifts right because three gas moles become two.

The equilibrium shift is not a complete cancellation and a catalyst does not change the final composition. It only speeds both directions.

For aA + bB ⇌ cC + dD, Kc = [C]^c[D]^d / ([A]^a[B]^b), using equilibrium concentrations in mol dm⁻³.

Products appear in the numerator, reactants in the denominator, and equation coefficients become powers. Pure solids and liquids are omitted because their effective concentration is constant.

For N₂ + 3H₂ ⇌ 2NH₃, Kc = [NH₃]²/([N₂][H₂]³). Do not use initial concentrations unless the question states they are equilibrium values.

Kc is not the same as a simple product-minus-reactant ratio. Coefficients become indices, and the expression depends on the balanced equation.

Mole fraction is the moles of one gas divided by total moles. For an ideal gas mixture, partial pressure pᵢ = xᵢP_total, where xᵢ is the mole fraction.

The partial pressures add to the total pressure. Use mole fraction rather than mass fraction because gas behaviour in the ideal model is tied to particle amount.

A mixture with 2 mol N₂ and 1 mol O₂ has x(N₂)=2/3. At total pressure 300 kPa, p(N₂)=200 kPa.

Partial pressure is not the pressure of a separated sample at the same volume. It is the component’s contribution within the mixture.

For gaseous reactions, Kp is written like Kc but uses equilibrium partial pressures. Coefficients in the balanced equation become powers.

Omit solids and liquids, and use a consistent pressure convention. The numerical value of Kp depends on the chosen pressure units unless the standard-state convention is fixed.

For N₂(g)+3H₂(g)⇌2NH₃(g), Kp = p(NH₃)²/(p(N₂)p(H₂)³).

Do not insert total pressure for every species. Each p value is that gas’s equilibrium partial pressure, and the reaction equation controls the powers.

To calculate Kc or Kp, first write the correct expression, then substitute equilibrium concentrations or partial pressures with the required powers and units.

Keep brackets grouped, evaluate powers before division, and check whether omitted phases are pure solids or liquids. A large or small value describes composition, not reaction speed.

If Kc=[C]²/([A][B]) and [A]=0.20, [B]=0.10, [C]=0.30 mol dm⁻³, then Kc=0.09/(0.02)=4.5.

Do not use concentrations at the start of the reaction unless they are explicitly equilibrium values. K is defined at equilibrium.

At equilibrium, the concentrations or partial pressures of all species are related by the balanced equation and the value of K. They are not required to be equal.

Use an ICE-style change table or stoichiometric ratios to express unknown equilibrium amounts in terms of one variable, then substitute into K and solve.

For A ⇌ 2B, if A decreases by x, B increases by 2x. The equilibrium values are [A]₀−x and [B]₀+2x, not independent guesses.

Dynamic equilibrium means forward and reverse rates are equal, not concentrations. The composition depends on K and the starting conditions.

For a given reaction, changing temperature changes the equilibrium constant because it changes the energy balance. Changing concentration or pressure shifts the equilibrium position but leaves K unchanged at that temperature.

Use Le Chatelier for the direction of a shift. For temperature, treat heat as a reactant or product and identify whether the forward reaction is endothermic or exothermic.

For N₂ + 3H₂ ⇌ 2NH₃ + heat, raising temperature shifts left and lowers K; adding nitrogen shifts right but does not change K.

A catalyst changes the rates of both directions, not K or the final equilibrium composition. Pressure changes K only indirectly if temperature also changes.

The Haber process makes ammonia: N₂(g)+3H₂(g)⇌2NH₃(g), an exothermic equilibrium with fewer gas moles on the product side.

High pressure favours ammonia, lower temperature favours yield but slows the reaction, and an iron catalyst increases rate without changing equilibrium position. Industry chooses a compromise condition and recycles unreacted gases.

Raising pressure improves equilibrium yield but increases equipment cost. A moderate temperature with an iron catalyst gives an economically useful throughput rather than the maximum possible yield.

The catalyst does not move equilibrium right. “Compromise temperature” means balancing rate and yield, not ignoring the exothermic nature of the reaction.

An acid is a proton donor and a base is a proton acceptor in the Brønsted–Lowry model. In water, acids form H₃O⁺ while bases remove protons from water or another donor.

The model focuses on proton transfer, so the same species can act as an acid in one reaction and a base in another.

HCl + H₂O → H₃O⁺ + Cl⁻: HCl donates H⁺ and water accepts it. NH₃ accepts H⁺ from water, so NH₃ is a base.

A base is not defined only as a substance containing OH⁻. Ammonia is a Brønsted base even though its formula has no hydroxide ion.

Common acids include hydrochloric acid HCl, nitric acid HNO₃ and sulfuric acid H₂SO₄. Common alkalis are soluble bases such as sodium hydroxide NaOH, potassium hydroxide KOH and aqueous ammonia NH₃(aq).

In water, acids increase hydronium concentration; soluble alkalis provide or generate hydroxide ions. Formula and state determine what particles are present.

HNO₃(aq) dissociates to H⁺/H₃O⁺ and NO₃⁻; NaOH(aq) provides Na⁺ and OH⁻. Their neutralisation produces water and a salt.

“Strong” and “concentrated” are different. Strength describes ionisation; concentration describes amount per volume.

A conjugate acid is formed when a base accepts a proton; a conjugate base is formed when an acid donates one. The members of a conjugate pair differ by exactly H⁺.

Label both proton transfers in the equation, not just the reactant called “acid” in everyday language.

In NH₃ + H₂O ⇌ NH₄⁺ + OH⁻, NH₃/NH₄⁺ is a conjugate base/acid pair and H₂O/OH⁻ is a conjugate acid/base pair.

Conjugate does not mean equal strength or identical charge. It means the formulas differ by one proton under the reaction conditions.

A strong acid dissociates almost completely in water; a strong base reacts or dissociates to produce hydroxide essentially completely. Weak acids and bases establish an equilibrium with substantial undissociated species.

Strength is an equilibrium property, not a measure of how much solution is present. A dilute strong acid can have lower concentration than a concentrated weak acid while still being more completely ionised.

HCl is strong, so [H₃O⁺] is close to its analytical concentration. Ethanoic acid is weak, so only part of CH₃COOH ionises and CH₃COOH ⇌ H⁺ + CH₃COO⁻ must be considered.

Do not use pH alone to label strength without knowing concentration. Concentration and degree of ionisation are separate ideas.

A weak acid or base ionises only partially in water, so the species and its ions coexist at equilibrium. The equilibrium position depends on the acid/base strength and concentration.

Write a reversible equation and use the conjugate pair to track proton transfer. A weak acid does not mean the solution contains no hydronium; it means most solute remains un-ionised.

Ethanoic acid follows CH₃COOH + H₂O ⇌ H₃O⁺ + CH₃COO⁻. Adding water changes concentrations and pH, but the acid remains a weak equilibrium system.

Weak is not the same as dilute. Strength describes degree of ionisation; dilution describes how much solute is present per volume.

A strong acid or base ionises essentially completely in water; a weak one establishes an equilibrium with a significant fraction un-ionised.

To compare strength, compare the ionisation equilibrium, not just the measured pH. A concentrated weak acid can have a lower pH than a very dilute strong acid while still being weaker.

0.001 mol dm⁻³ HCl is strong and nearly fully ionised, whereas 0.10 mol dm⁻³ ethanoic acid is weak and only partly ionised. Their pH values cannot alone define strength.

Strong does not mean concentrated, and weak does not mean harmless. Keep degree of ionisation and amount per volume separate.

Neutralisation is a reaction in which an acid and a base react so that proton transfer produces water. In the simplest ionic form, H₃O⁺ + OH⁻ → 2H₂O.

The other ions remain in solution as the salt’s ions. Write the molecular equation, then identify the net ionic proton-transfer step when useful.

HCl(aq)+NaOH(aq)→NaCl(aq)+H₂O(l). The net reaction is H⁺(aq)+OH⁻(aq)→H₂O(l), while Na⁺ and Cl⁻ are spectators.

Neutralisation does not always produce pH 7: a weak acid/strong base mixture or excess reagent can leave the solution acidic or basic.

A salt forms when the acidic proton is replaced by a metal ion or ammonium ion. The salt remains in solution as ions, and some ions can react with water by hydrolysis.

The parent acid and base determine whether the salt solution is approximately neutral, acidic or basic. Do not infer pH from the word “salt” alone.

NaCl from strong acid/strong base is approximately neutral. CH₃COONa contains acetate, the conjugate base of a weak acid, so acetate can accept protons from water and make the solution basic.

Salt formation is not the same as complete removal of every proton from every acid. Polyprotic acids and weak conjugate ions need separate analysis.

A titration curve plots pH against the volume of acid or alkali added. Its shape reflects the relative strengths of the reactants and the equivalence point where stoichiometric neutralisation is complete.

Strong/strong curves have a steep jump centred near pH 7; weak/strong combinations shift the equivalence pH because the conjugate ion hydrolyses. Half-equivalence can identify a weak acid’s pKa.

Adding NaOH to ethanoic acid gives a buffer region before the steep rise and an equivalence point above pH 7, unlike HCl/NaOH.

Equivalence point is not always pH 7 and is not the same as the point where an indicator happens to change colour.

An indicator is suitable when its transition range lies within the sharp pH change near the equivalence point. The endpoint colour change should therefore occur close to the stoichiometric endpoint.

Choose from the expected curve, not from the acid name alone. Strong acid/strong base allows a broad choice; weak acid/strong base needs an indicator changing in the basic region.

Phenolphthalein is suitable for a weak acid–strong base titration because its transition is above pH 7; methyl orange is better suited to a strong acid–weak base curve.

An indicator does not measure exact equivalence by magic. If its transition range falls outside the steep section, the endpoint error can be significant.