In simple terms
A friendly intro before the formal notes — no formulas yet.
The Proton Dance
Acids and bases are defined by their ability to donate or accept protons, respectively. This simple exchange governs everything from the acidity of a solution (pH) to the function of biological buffers.
Imagine a ballroom where dancers represent molecules and a special hat represents a proton (H⁺). An 'acid' dancer is one who gives their hat to another dancer. A 'base' dancer is one who accepts a hat. A buffer solution is like a large group of dancers who are very good at quickly passing the hat back and forth amongst themselves, ensuring that the number of dancers left without a hat at any one time remains almost constant, even if new dancers join or leave the floor.
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Brønsted–Lowry: acid proton donor, base proton acceptor.
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Ka, pKa, pH — weak acid partial dissociation.
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Buffer resists pH change — weak acid + conjugate base.
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Titration curves — equivalence point pH depends on acid/base strength.
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Key formulas
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$pH = -log_{10}[H^+(aq)]$
$For a weak acid HA: HA(aq) \rightleftharpoons H^+(aq) + A^-(aq) \ K_a = \frac{[H^+(aq)][A^-(aq)]}{[HA(aq)]}$
$pH = pK_a + log_{10}\left(\frac{[A^-]}{[HA]}\right)$
Full topic notes
Formal explanation with the rigour you need for the exam.
1. The Brønsted–Lowry Theory
The Brønsted–Lowry theory provides a more general definition of acids and bases than the Arrhenius theory. It focuses on the transfer of a proton, which is simply a hydrogen ion, H⁺. An acid is a proton donor, and a base is a proton acceptor. When an acid donates a proton, the species that remains is its conjugate base. When a base accepts a proton, it forms its conjugate acid. Together, they form a conjugate acid-base pair.
Consider the reaction of ammonia with water: . Here, water donates a proton to ammonia. Therefore, acts as an acid and acts as a base. The ammonium ion, , is the conjugate acid of , and the hydroxide ion, , is the conjugate base of . Substances like water that can act as both an acid and a base are called amphoteric.
2. pH, Ka, and pKa: Quantifying Acidity
The pH scale provides a convenient way to express the concentration of hydrogen ions in a solution. It is a logarithmic scale, meaning a change of one pH unit represents a tenfold change in .
pH = -log_{10}[H^+(aq)]
Strong acids, like HCl, fully dissociate in water, so for a $0.1 \ mol \ dm^{-3}$ solution of HCl, $[H^+] = 0.1 \ mol \ dm^{-3}$. Weak acids, like ethanoic acid (), only partially dissociate. This partial dissociation is an equilibrium, described by the acid dissociation constant, Ka.
A larger Ka value indicates a greater extent of dissociation and therefore a stronger acid.
pKa is used for convenience: $pK_a = -log_{10}(K_a)$.
A smaller pKa value indicates a stronger acid.
The ionic product of water, $K_w = [H^+][OH^-] = 1.0 \times 10^{-14} \ mol^2 \ dm^{-6}$ at 298 K, links the concentrations of H⁺ and OH⁻ in any aqueous solution.
3. Buffer Solutions
A buffer solution is a remarkable chemical system that resists changes in pH when small quantities of acid or alkali are added. Acidic buffers are typically made from a weak acid and its conjugate base (usually from a salt). For example, a mixture of ethanoic acid () and sodium ethanoate (). The solution contains a large reservoir of both the weak acid and its conjugate base.
$pH = pK_a + log_{10}\left(\frac{[A^-]}{[HA]}\right)$
Adding acid (H⁺): The excess H⁺ ions are removed by reacting with the conjugate base: . The equilibrium shifts to the left.
Adding alkali (OH⁻): The added OH⁻ ions are neutralised by the weak acid: . The equilibrium shifts to the right.
In both cases, the large reservoirs of HA and A⁻ ensure the change in the ratio is small, thus the pH change is minimal.
The pH of a buffer is determined by the pKa of the weak acid and the ratio of the concentrations of the conjugate base and weak acid.
When asked to explain how a buffer works, you must write two separate equations: one showing the reaction with added H⁺ and one showing the reaction with added OH⁻. State that there are large reservoirs of both the weak acid and its conjugate base.
4. Acid-Base Titration Curves
A titration curve is a graph of pH against the volume of titrant added. The shape of the curve provides valuable information about the strengths of the acid and base involved. Key features include the initial pH, the buffer region (for weak species), the equivalence point (steepest part of the curve), and the final pH.
Strong Acid - Strong Base: Equivalence point at pH 7. Steep vertical section.
Weak Acid - Strong Base: Initial pH is higher (weak acid). A buffer region exists before the equivalence point. Equivalence point is above pH 7 due to the hydrolysis of the conjugate base formed ().
Strong Acid - Weak Base: Equivalence point is below pH 7 due to the hydrolysis of the conjugate acid formed ().
Indicator Choice: A suitable indicator must have a colour change range (pKin) that falls entirely within the steep vertical section of the titration curve.
When sketching titration curves, always label your axes (pH on y-axis, Volume of titrant/cm³ on x-axis). Mark the volume and pH at the equivalence point. For weak acid/base titrations, also indicate the approximate pH at the half-equivalence point, where .
Worked examples
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Calculate the pH of a $0.0500 \ mol \ dm^{-3}$ solution of propanoic acid at 298 K. The for propanoic acid is $1.35 \times 10^{-5} \ mol \ dm^{-3}$.
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Write the Ka expression:
A buffer solution is made by dissolving 12.3 g of sodium ethanoate () in 250 cm³ of $0.800 \ mol \ dm^{-3}$ ethanoic acid. Calculate the pH of the resulting buffer solution. ( of ; for $CH_3COOH = 1.75 \times 10^{-5} \ mol \ dm^{-3}$)
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Calculate moles of conjugate base (ethanoate):
How it all connects
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Glossary
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Quick check
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Revision flashcards
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What is a Brønsted–Lowry acid?
A proton (H⁺) donor.
Key takeaways
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A larger Ka value indicates a greater extent of dissociation and therefore a stronger acid.
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pKa is used for convenience: $pK_a = -log_{10}(K_a)$.
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A smaller pKa value indicates a stronger acid.
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The ionic product of water, $K_w = [H^+][OH^-] = 1.0 \times 10^{-14} \ mol^2 \ dm^{-6}$ at 298 K, links the concentrations of H⁺ and OH⁻ in any aqueous solution.
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