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Loading home…A conjugate acid–base pair differs by one proton (H⁺).
The acid donates a proton; its conjugate base is what remains.
The base accepts a proton; its conjugate acid is the protonated form.
Example: NH₃ (base) and NH₄⁺ (conjugate acid) form a pair.
In water, H₂O/H₃O⁺ and H₂O/OH⁻ are also conjugate pairs.
Every Brønsted–Lowry acid–base reaction involves two such pairs.
A Lewis acid accepts a pair of electrons.
A Lewis base donates a pair of electrons.
BF₃ and AlCl₃ are common Lewis acids (empty orbitals).
NH₃ and H₂O are Lewis bases (lone pairs on N or O).
The Lewis concept covers reactions without proton transfer.
Example: BF₃ + :NH₃ → F₃B←NH₃ (adduct).
Kw = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴ at 25 °C.
pH = −log[H⁺] and pOH = −log[OH⁻].
At 25 °C, pH + pOH = 14.
Neutral water has [H⁺] = [OH⁻] = 10⁻⁷ mol dm⁻³, so pH = 7.
Acidic solutions have pH < 7; basic solutions have pH > 7.
Kw changes with temperature, so the neutral pH is 7 only near 25 °C.
The common ion effect is the suppression of ionization of a weak electrolyte by adding a strong electrolyte that shares an ion.
Example: CH₃COOH ⇌ H⁺ + CH₃COO⁻; adding CH₃COONa increases [CH₃COO⁻].
Le Chatelier’s principle shifts the equilibrium left, lowering [H⁺].
Buffers use this idea: a weak acid plus its salt (common anion).
Solubility of sparingly soluble salts also falls if a common ion is added.
The effect is important in analytical chemistry and buffer design.
Salt hydrolysis is reaction of a salt’s ions with water, changing pH.
Salts of strong acid + strong base (e.g. NaCl) do not hydrolyse; pH ≈ 7.
Salts of weak acid + strong base (e.g. CH₃COONa) give basic solutions.
Salts of strong acid + weak base (e.g. NH₄Cl) give acidic solutions.
Salts of weak acid + weak base may be acidic, basic, or nearly neutral.
Hydrolysis explains why many “neutral-looking” salts are not pH 7.
Ksp is the solubility product constant for a sparingly soluble ionic solid.
For MxAy(s) ⇌ x Mⁿ⁺ + y Aᵐ⁻, Ksp = [Mⁿ⁺]ˣ[Aᵐ⁻]ʸ (solid omitted).
A larger Ksp means greater molar solubility (for the same stoichiometry type).
Comparing ion product Q with Ksp predicts precipitation.
Common ions decrease solubility via Le Chatelier’s principle.
Ksp is used in qualitative analysis and water hardness chemistry.
Arrhenius defined an acid as a substance that increases [H⁺] in water and a base as one that increases [OH⁻] in water.
Examples: HCl → H⁺ + Cl⁻; NaOH → Na⁺ + OH⁻.
Limitation: Arrhenius applies mainly in aqueous solution and cannot classify NH₃ well without invoking OH⁻ from water.
Brønsted–Lowry defines an acid as a proton donor and a base as a proton acceptor.
This covers aqueous and non-aqueous proton transfer, e.g. HCl + NH₃ → NH₄⁺ + Cl⁻.
Conjugate pairs appear naturally: acid/conjugate base differ by one H⁺.
Lewis defines an acid as an electron-pair acceptor and a base as an electron-pair donor.
BF₃, Al³⁺ and H⁺ are Lewis acids; NH₃ and OH⁻ are Lewis bases.
Lewis is the broadest concept and includes adduct formation without proton transfer.
All Arrhenius acids/bases fit Brønsted–Lowry in water; Brønsted acids/bases fit Lewis when H⁺ transfer is viewed via electron pairs.
School answers should give definitions, one example each, and one clear limitation of Arrhenius.
State that “strong” vs “weak” refers to extent of ionization, not to which definition is used.
Indicators and pH still rely on [H⁺]/[OH⁻] in water.
Titration theory uses Brønsted proton transfer most often.
Lewis ideas appear when discussing BF₃ complexes and metal cations.
Use consistent symbols H⁺ or H₃O⁺ when discussing aqueous acids.
A good comparison in clear points beats vague statements.
Always link each definition to a concrete reaction equation.
Mention that water is amphoteric in the Brønsted sense.
Conclude by ranking scope: Arrhenius < Brønsted–Lowry < Lewis.
For a strong monoprotic acid of concentration C, [H⁺] ≈ C if fully ionized.
Then pH = −log C (e.g. 0.01 M HCl → [H⁺] = 0.01 → pH = 2).
For a strong base such as 0.1 M NaOH, [OH⁻] ≈ 0.1, pOH = 1, pH = 13 at 25 °C.
For a weak acid HA with Ka, set up HA ⇌ H⁺ + A⁻ and Ka = [H⁺][A⁻]/[HA].
If α is small, [H⁺] ≈ √(Ka C), then pH = −log[H⁺].
Check the approximation: α = [H⁺]/C should be ≪ 1 (often < 0.05).
For a weak base B with Kb, [OH⁻] ≈ √(Kb C), then pOH and pH = 14 − pOH.
Relate Ka and Kb for a conjugate pair: Ka × Kb = Kw.
Percent ionization = ([H⁺]/C)×100% for weak acids.
Dilution increases percent ionization of a weak acid even as pH rises.
Polyprotic acids need stepwise Ka1, Ka2, …; often Ka1 dominates [H⁺].
Always specify temperature because Kw and Ka/Kb depend on T.
Report answers with sensible significant figures.
Show ICE tables clearly in exam work.
Distinguish strong vs weak before choosing a formula.
Never use √(Ka C) for a strong acid.
Link calculated pH to indicator choice in titrations.
Practise both “find pH from C” and “find C from pH” styles.
For [H⁺] = 1×10⁻⁹, pH = 9 (basic relative to pure water).
State Kw = 1×10⁻¹⁴ when converting pH ↔ pOH at 25 °C.
A buffer resists large pH change when small amounts of acid or base are added.
Typical acidic buffer: weak acid + its salt (conjugate base), e.g. CH₃COOH/CH₃COONa.
Typical basic buffer: weak base + its salt (conjugate acid), e.g. NH₃/NH₄Cl.
Added H⁺ is consumed by the conjugate base (A⁻ + H⁺ → HA).
Added OH⁻ is consumed by the weak acid (HA + OH⁻ → A⁻ + H₂O).
The Henderson–Hasselbalch equation for an acidic buffer is pH = pKa + log([A⁻]/[HA]).
When [A⁻] = [HA], pH = pKa; this is often the most effective region.
Buffer capacity is highest when concentrations of both components are large and similar.
Dilution of a good buffer changes pH little if the ratio [A⁻]/[HA] stays constant.
Blood and many biological fluids are buffered near physiological pH.
In calculations, use equilibrium concentrations (often ≈ analytical concentrations).
Choose a buffer whose pKa is close to the desired pH.
Adding too much strong acid/base exhausts one component and breaks the buffer.
Indicators are not buffers; they are weak acids/bases used in tiny amounts.
Common-ion reasoning explains why the weak acid does not fully ionize.
Exam answers should include definition, composition, mechanism, and the HH equation.
Give one numerical sketch: if pKa = 4.74 and [A⁻]/[HA] = 1, pH = 4.74.
Relate buffers to salt hydrolysis and conjugate pairs.
State limitations: extreme dilution or huge additions defeat buffering.
Mention that buffers keep titration curves flatter before the steep jump.
Ksp is the equilibrium constant for dissolution of a sparingly soluble ionic solid.
For AgCl(s) ⇌ Ag⁺ + Cl⁻, Ksp = [Ag⁺][Cl⁻].
For CaF₂(s) ⇌ Ca²⁺ + 2F⁻, Ksp = [Ca²⁺][F⁻]².
If s is molar solubility of AgCl, then Ksp = s² so s = √Ksp.
For CaF₂, Ksp = (s)(2s)² = 4s³, so s = ∛(Ksp/4).
Stoichiometry must match the formula when converting Ksp ↔ s.
If Q (ion product) > Ksp, precipitation occurs until Q = Ksp.
If Q < Ksp, more solid can dissolve.
A common ion lowers solubility: adding NaCl to AgCl(aq) increases [Cl⁻] and shifts left.
This is used to wash precipitates and in selective precipitation.
Temperature usually changes Ksp; hotter water often dissolves more solid.
Complexation or pH can increase apparent solubility (e.g. amphoteric hydroxides).
Compare salts of the same type carefully; do not rank solubility by Ksp alone across different formulas.
In qualitative analysis, H₂S and OH⁻ control [S²⁻] or [OH⁻] via Ksp ideas.
Always omit pure solids from the Ksp expression.
Show units consistently (mol dm⁻³ for concentrations).
State assumptions: activity ≈ concentration in dilute solution.
Link common-ion effect to Le Chatelier’s principle.
A complete answer defines Ksp, works one solubility example, and explains common ion.
For AB(s), remember Ksp = [A⁺][B⁻].
Titration finds the concentration of an acid or base using a standard solution.
The equivalence point is when stoichiometrically equal amounts of acid and base have reacted.
The end point is the observed indicator colour change; it should match equivalence closely.
For strong acid–strong base, the equivalence pH is about 7 and the jump is sharp.
Both methyl orange and phenolphthalein can work for strong–strong titrations.
For strong acid–weak base, equivalence is acidic; methyl orange is suitable.
For weak acid–strong base, equivalence is basic; phenolphthalein is suitable.
Weak–weak titrations have poor pH jumps; indicators are less reliable.
An indicator is a weak acid/base whose conjugate forms have different colours.
It changes colour over a characteristic pH range (about 2 units).
Choose an indicator whose range lies on the steep part of the titration curve.
Phenolphthalein is colourless in acid and pink in base (approx. pH 8.2–10).
Methyl orange is red in strong acid and yellow in base (approx. pH 3.1–4.4).
Volume readings, rinsing errors, and parallax affect accuracy.
Use a white tile under the flask to see colour change clearly.
Calculate concentration from n = CV and the balanced mole ratio.
Sketching a titration curve helps justify indicator choice.
Buffers appear near the middle of weak-acid titrations.
Salt hydrolysis explains why equivalence pH is not always 7.
Exam answers should define equivalence vs end point and match indicator to titration type.