1. What are acids and bases?
There are several ways to define acids and bases. Two important models are the Arrhenius and Brønsted–Lowry definitions.
Arrhenius acid
A substance that increases the concentration of H⁺ ions in aqueous solution. More precisely, these protons exist mainly as hydrated ions.
Arrhenius base
A substance that increases the concentration of OH⁻ ions in aqueous solution.
Brønsted–Lowry acid
A proton donor. It transfers an H⁺ ion to another species.
Brønsted–Lowry base
A proton acceptor. It accepts an H⁺ ion from an acid.
Examples
- HCl acts as an acid in water.
- NaOH dissociates to provide OH⁻ ions in water.
- NH₃ acts as a Brønsted–Lowry base by accepting a proton from water.
Ammonia in water
Ammonia accepts a proton from water. In this reaction, water acts as the acid and ammonia acts as the base.
2. Conjugate acid–base pairs
When a Brønsted–Lowry acid donates a proton, it becomes its conjugate base. When a base accepts a proton, it becomes its conjugate acid.
HA/A⁻ is one conjugate pair, while B/BH⁺ is the other.
Conjugate pairs differ by one proton. A stronger acid has a weaker conjugate base, and a weaker acid has a stronger conjugate base, considered in the same solvent and under comparable conditions.
3. Strong and weak acids and bases
Strength describes how extensively an acid or base ionises or reacts with the solvent. It is not the same as concentration.
| Type | Behaviour in water | Example |
|---|---|---|
| Strong acid | Ionises nearly completely under ordinary dilute aqueous conditions. | HCl |
| Weak acid | Ionises only partially and establishes an equilibrium. | CH₃COOH |
| Strong base | Dissociates extensively to provide basic ions in water. | NaOH |
| Weak base | Reacts partially with water to establish an equilibrium. | NH₃ |
Important distinction: A concentrated weak acid can contain more acid per litre than a dilute strong acid. Strength describes ionisation behaviour; concentration describes the amount present per volume.
4. Understanding pH
pH is a logarithmic measure related to hydrogen-ion activity. In introductory calculations for dilute aqueous solutions, it is commonly approximated using the molar concentration of H⁺.
pH
For this introductory formula, [H⁺] is expressed in mol L⁻¹ and treated as a numerical concentration relative to the standard concentration.
pOH
For dilute aqueous solutions, pOH is similarly calculated from the hydroxide-ion concentration.
Relationship between pH and pOH
Here pKᵥ is the negative logarithm of the ionic product of water under the specified conditions. At 25°C, pKᵥ is approximately 14.00.
| At 25°C, approximately | Classification |
|---|---|
| pH below 7 | Acidic |
| pH equal to 7 | Neutral |
| pH above 7 | Basic |
Neutral pH is temperature-dependent. A pH of 7 is neutral at approximately 25°C, but neutrality does not correspond to pH 7 at every temperature.
Worked example: pH of a strong acid
Suppose a dilute HCl solution has a concentration of 0.001 mol L⁻¹. Assuming complete ionisation:
- [H⁺] ≈ 0.001 = 10⁻³ mol L⁻¹.
- pH = −log₁₀(10⁻³).
- Therefore, pH ≈ 3.
5. Ionisation of water
Water undergoes a small amount of self-ionisation. The equilibrium can be represented as:
Kᵥ is the ionic product of water in the usual introductory concentration convention. Its value changes with temperature.
In neutral water, the activities of H₃O⁺ and OH⁻ are equal. At 25°C, each ion has a concentration of approximately 1.0 × 10⁻⁷ mol L⁻¹ in pure water.
6. Acid and base dissociation constants
Weak acids and bases establish equilibria in water. Their dissociation constants describe the positions of these equilibria.
Acid dissociation constant
A larger Kₐ generally indicates a stronger acid in the same solvent and under the same conditions.
Base dissociation constant
A larger Kᵦ generally indicates a stronger base under comparable conditions.
pK values
For a conjugate acid–base pair in water, KₐKᵦ = Kᵥ under the corresponding standard equilibrium conventions. Therefore, pKₐ + pKᵦ = pKᵥ.
7. Neutralisation reactions
Neutralisation commonly occurs when an acid reacts with a base. In many aqueous reactions involving a strong acid and a strong base, hydrogen ions and hydroxide ions form water.
Net ionic equation
This is a simplified representation; in water, the proton is present mainly as a hydrated ion.
Example
Hydrochloric acid reacts with sodium hydroxide to form sodium chloride and water.
Neutralisation does not always produce a solution with pH 7. The final pH depends on acid and base strengths, their amounts, concentration, and temperature.
8. Indicators and titration
Acid–base indicators are substances that change colour over a particular pH range. They help estimate the acidity of a solution or identify the endpoint of a titration.
| Indicator | Acidic colour | Basic colour |
|---|---|---|
| Litmus | Red | Blue |
| Phenolphthalein | Colourless | Pink |
| Methyl orange | Red | Yellow |
Indicator transition ranges are not identical to the entire acidic or basic pH range. Choose an indicator whose transition range is appropriate for the titration's equivalence region.
Acid–base titration calculation
For a reaction with stoichiometric coefficients a for the acid and b for the base, this relation applies at equivalence when M is molar concentration and V is volume in consistent units. It assumes the stated reaction stoichiometry.
For a 1:1 reaction, such as HCl with NaOH, the relationship reduces to MₐVₐ = MᵦVᵦ.
9. Buffer solutions
A buffer solution resists large pH changes when small amounts of acid or base are added. A common buffer contains a weak acid and its conjugate base, or a weak base and its conjugate acid.
Henderson–Hasselbalch equation
This approximation is useful for a buffer made from a weak acid and its conjugate base when the concentrations are suitable approximations to their activities.
- Adding a small amount of acid is countered mainly by the conjugate base.
- Adding a small amount of base is countered mainly by the weak acid.
- Buffers have a limited capacity; large additions can overwhelm them.
10. Key ideas to remember
- Brønsted–Lowry acids donate protons; bases accept protons.
- Strong and weak describe ionisation behaviour, not concentration.
- pH is logarithmic: a one-unit decrease corresponds to a tenfold increase in hydrogen-ion activity.
- At 25°C, pH + pOH is approximately 14 for dilute aqueous solutions using the usual convention.
- Kₐ and Kᵦ describe acid and base equilibria.
- Neutralisation commonly forms water and a salt, but the final pH is not always 7.
- Indicators have specific transition ranges.
- Buffers resist pH changes but have limited capacity.