Difficulty: Intermediate to Advanced | Prerequisites: Ka/pKa concepts, ICE tables, acid-base equilibria fundamentals
This material builds on everything from the Ka/pKa unit. Buffers are one of the most exam-heavy topics in CHEM 202: you need to know how they resist pH changes, how to pick the right acid for a target pH, and how to use the Henderson-Hasselbalch equation. Titration curves and equivalence-point behaviour tie the whole acid-base unit together. Lewis acid-base theory broadens the definition beyond proton transfer, and salt hydrolysis explains why "neutral" salts can shift solution pH.
Buffers resist pH changes by neutralising added acid or base using a weak acid/conjugate base pair. To choose the right buffer acid, match its pKa to the target pH. Lewis acids accept electron pairs (metal cations, electron-deficient molecules like BCl3), Lewis bases donate them (anions with lone pairs, NH3). At the equivalence point of a titration, the pH depends on which type of acid and base were mixed: weak acid + strong base gives pH above 7, strong acid + strong base gives pH of exactly 7.
Buffer
A solution that resists changes in pH when small amounts of acid or base are added. It contains a weak acid and its conjugate base (or a weak base and its conjugate acid) in comparable concentrations.
Think of it as a chemical shock absorber for pH.
Henderson-Hasselbalch equation
The equation pH = pKa + log([A⁻]/[HA]), which relates the pH of a buffer to the pKa of its weak acid and the ratio of conjugate base to acid concentrations.
In simple terms, this is the shortcut formula for buffer pH problems. If [A⁻] = [HA], the log term is zero and pH = pKa.
Equivalence point
The point in a titration where the moles of titrant exactly equal the moles of analyte. All of the original acid (or base) has been neutralised.
Think of it as the moment you have added exactly enough base to react with all the acid, or vice versa.
Half-equivalence point
The point in a titration where exactly half the acid has been neutralised. At this point, [HA] = [A⁻], so pH = pKa.
In simple terms, this is where you can read the pKa directly off the titration curve.
Lewis acid
A species that can accept a pair of electrons. Metal cations (Cu2+, Fe2+, Al3+) and electron-deficient molecules (BCl3) are Lewis acids.
Lewis base
A species that can donate a pair of electrons. Anions with lone pairs (Cl⁻, OH⁻) and neutral molecules with lone pairs (NH3) are Lewis bases.
Salt hydrolysis
The reaction of a salt's ions with water to produce acidic or basic solutions. Whether a salt solution is acidic, basic, or neutral depends on the strength of the parent acid and base.
Titration curve
A graph of pH versus volume of titrant added. Its shape tells you about the acid's strength, the equivalence point, and the buffer region.
A buffer made from equal amounts of ammonia (NH3) and ammonium chloride (NH4Cl) uses the equilibrium:
NH4+(aq) + H2O(l) ⇌ NH3(aq) + H3O+(aq)
When you add a strong acid:
The added H3O+ reacts with NH3 (the base component)
The equilibrium shifts to the left
[NH4+] increases, [NH3] decreases
The pH drops slightly, but much less than it would without the buffer
When you add a strong base:
The added OH⁻ reacts with NH4+ (the acid component)
The equilibrium shifts to the right
[NH4+] decreases, [NH3] increases
The pH rises slightly, but again much less than it would in plain water
A buffer works most effectively when pH ≈ pKa. To prepare a buffer at a target pH, pick the acid whose pKa is closest to that pH.
Example: to make a buffer at pH 4.5, convert each candidate's Ka to pKa:
Acid | Ka | pKa | Closest to 4.5? |
|---|---|---|---|
HOC6H4OCOOH | 1.0 x 10^(-3) | 3.00 | No |
C6H4(COOH)2 | 2.9 x 10^(-4) | 3.54 | No |
CH3COOH | 1.8 x 10^(-5) | 4.74 | Yes (closest) |
C5H5O5COOH | 4.0 x 10^(-6) | 5.40 | No |
HBrO | 2.3 x 10^(-9) | 8.64 | No |
Acetic acid (CH3COOH) is the correct choice because 4.74 is nearest to 4.5.
The Lewis definition is broader than Bronsted-Lowry. It does not require proton transfer.
Lewis acids accept an electron pair:
Metal cations: Cu2+, Fe2+, Al3+ (they are electron-deficient and can accept lone pairs from ligands)
Electron-deficient molecules: BCl3 (boron has only 6 electrons, can accommodate 2 more)
Lewis bases donate an electron pair:
Anions with lone pairs: Cl⁻, OH⁻
Neutral molecules with lone pairs: NH3
Quick identification rule: if the species has an incomplete octet or a positive charge on a metal, it is likely a Lewis acid. If it has lone pairs available to share, it is a Lewis base.
Titration type | pH at equivalence point | Why |
|---|---|---|
Weak acid + strong base | Greater than 7.0 | The conjugate base of the weak acid remains in solution and is itself basic |
Strong acid + strong base | Equal to 7.0 | Neither ion hydrolyses; only water is present |
Weak base + strong acid | Less than 7.0 | The conjugate acid of the weak base remains in solution and is itself acidic |
When a salt dissolves, its ions may react with water:
Anions that are conjugate bases of weak acids make the solution basic: HCO3⁻, CO3 2⁻, SO4 2⁻, CH3COO⁻, HCOO⁻
Cations that are conjugate acids of weak bases (or small, highly charged metal ions) can make the solution acidic: NH4+, Al3+, Fe2+
Spectator ions have no effect on pH: alkali metal cations (Na+, K+), alkaline earth cations (Ca2+, Mg2+ in many contexts), halide anions (Cl⁻, Br⁻), and nitrate (NO3⁻)
Henderson-Hasselbalch equation:
\text{pH} = \text{pK}_a + \log\left(\frac{[\text{A}^-]}{[\text{HA}]}\right)When [A⁻] = [HA], the log term equals zero, so pH = pKa. This is the condition at the half-equivalence point of a weak acid/strong base titration.
Useful rearrangement: if you need the ratio of base to acid for a given pH:
\frac{[\text{A}^-]}{[\text{HA}]} = 10^{(\text{pH} - \text{pK}_a)}Blood is buffered at pH 7.4 by the carbonic acid/bicarbonate system (H2CO3/HCO3⁻). Without this buffer, even mild metabolic changes would cause dangerous pH swings. This is the same Henderson-Hasselbalch logic applied in a biological context.
Titrations are a routine analytical technique in quality control: a food chemist titrates a juice sample with NaOH to determine its total acid content, and a pharmaceutical lab titrates drug formulations to verify active-ingredient concentration.
Students often assume the equivalence point of any titration is at pH 7. It is only 7.0 for a strong acid/strong base titration. For weak acid/strong base, the equivalence point is above 7.
A buffer does not prevent pH from changing at all. It limits the change. If you add a very large amount of acid or base, the buffer will be overwhelmed and the pH will shift significantly.
Students sometimes confuse Lewis and Bronsted-Lowry definitions. Every Bronsted acid is a Lewis acid, but the reverse is not always true. Metal cations like Cu2+ and Fe2+ are Lewis acids but they do not donate protons.
When identifying salt effects on pH, students forget that many common ions are spectators. Na+, K+, Cl⁻, NO3⁻, and the halides do not affect pH. Only ions derived from weak acids or weak bases (or small, highly charged metals) cause hydrolysis.
⚠️ Buffer problems using Henderson-Hasselbalch are a near-certainty. Be comfortable plugging in pKa, [A⁻], and [HA] and solving for pH, or solving for the ratio given a target pH.
⚠️ Expect a question asking you to select the best buffer acid for a target pH from a list of Ka values. Convert each to pKa and pick the one closest to the target.
⚠️ Know the equivalence point pH rules cold. Weak acid + strong base = pH above 7. Strong acid + strong base = pH exactly 7. This is a quick conceptual question that appears frequently.
⚠️ Be able to identify Lewis acids and Lewis bases from a list of species. Look for electron-deficient species (Lewis acids) and lone-pair donors (Lewis bases).
⚠️ Salt hydrolysis questions may ask you to predict whether a salt solution is acidic, basic, or neutral. Trace each ion back to its parent acid or base.
⚠️ Know how to read a titration curve: identify the equivalence point (steepest part of the curve), the half-equivalence point (halfway up the buffer region, where pH = pKa), and the buffer region (the relatively flat portion before equivalence).
True or False: A buffer made from NH3 and NH4Cl will have its pH increase when a strong acid is added.
Fill in the blank: At the half-equivalence point of a weak acid titration, pH = ______.
True or False: BCl3 is a Lewis base because it has three chlorine atoms with lone pairs.
Fill in the blank: When a weak acid is titrated with a strong base, the pH at the equivalence point is ______ 7.0.
True or False: A solution of NaCl in water has a pH of exactly 7.0 (at 25°C).
Q: A buffer is prepared from 0.20 M acetic acid (pKa = 4.74) and 0.10 M sodium acetate. What is the pH of the buffer?
A: pH = 4.74 + log(0.10/0.20) = 4.74 + log(0.5) = 4.74 + (-0.30) = 4.44.
Q: You need a buffer at pH 9.0. Which acid would you choose: one with Ka = 1.0 x 10^(-5) (pKa = 5.0) or one with Ka = 1.0 x 10^(-9) (pKa = 9.0)?
A: The acid with pKa = 9.0. A buffer works best when pKa ≈ target pH.
Q: From this list, identify all Lewis acids and all Lewis bases: BCl3, Cu2+, Cl⁻, Fe2+, NH3, OH⁻, Al3+.
A: Lewis acids: BCl3 (incomplete octet on B), Cu2+, Fe2+, Al3+ (all electron-deficient metal cations). Lewis bases: Cl⁻, NH3, OH⁻ (all have lone pairs to donate).
Q: A solution of sodium acetate (NaCH3COO) in water will be acidic, basic, or neutral?
A: Basic. Na+ is a spectator ion. CH3COO⁻ is the conjugate base of the weak acid acetic acid, so it hydrolyses to produce OH⁻, raising the pH above 7.
Q: During a titration of HCl (strong acid) with NaOH (strong base), what is the pH at the equivalence point?
A: Exactly 7.0. Both ions produced (Na+ and Cl⁻) are spectators and do not hydrolyse.
Q: In a weak acid/strong base titration, where on the curve can you read the pKa of the acid?
A: At the half-equivalence point. This is the point where half the acid has been neutralised, so [HA] = [A⁻] and pH = pKa.
Buffer chemistry connects directly to the Ka/pKa material: you cannot use Henderson-Hasselbalch without knowing pKa, and the buffer region of a titration curve is where the acid and conjugate base coexist.
Lewis acid-base theory extends into coordination chemistry and organic reaction mechanisms (nucleophilic attack is a Lewis base donating electrons to a Lewis acid electrophile).
Salt hydrolysis ties back to conjugate acid-base pairs: predicting whether a salt is acidic or basic requires tracing each ion to the acid or base it came from.
Related Terms / Search Tags
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