A chemist titrates phosphoric acid (H₃PO₄) with NaOH and observes two clear inflection points on the potentiometric curve, near pH 4.6 and pH 9.8, but no clear third inflection. Which equivalence points are analytically useful for quantifying phosphoric acid?
AAll three are equally useful; the third inflection is there but requires more sensitive equipment to detect
BOnly the first, near pH 4.6, because the subsequent inflections represent contamination by carbonate
CThe first and second inflection points are analytically useful; the third proton loss is too gradual and buffered near neutral/alkaline pH to give a reliable endpoint
DNone — polyprotic acids cannot be quantified by titration because overlapping Ka values prevent distinct equivalence points
For phosphoric acid, the three Ka values are sufficiently spaced (Ka1 ≈ 10⁻², Ka2 ≈ 10⁻⁷, Ka3 ≈ 10⁻¹²) that the first two proton losses produce distinct, quantifiable equivalence points. However, the third pKa (~12.4) means the final deprotonation occurs under extremely alkaline conditions where CO₂ absorption from air and carbonate interference make endpoint detection unreliable. Recognizing that not all equivalence points are analytically accessible — and choosing the useful ones — is a core skill in quantitative titration.
Question 2 Multiple Choice
A student needs to determine the total acid content of a dark-red berry juice. When they try to use phenolphthalein indicator, the color change is invisible against the sample's natural color. What is the best approach?
ADilute the sample 10-fold so the color is pale enough for the indicator to be seen
BSwitch to a different colored indicator, such as methyl orange, that changes at a lower pH
CUse potentiometric endpoint detection — monitor pH with a glass electrode and locate the equivalence point from the first derivative of the pH-vs-volume curve
DFilter the juice through activated charcoal to remove color before titrating with phenolphthalein
Potentiometric detection is the correct solution for colored, turbid, or opaque samples where visual indicators are obscured. A glass pH electrode continuously monitors the pH as titrant is added; the equivalence point appears as a sharp spike in the first derivative (dpH/dV), which is detectable regardless of sample color or opacity. Option A (dilution) may actually worsen endpoint detection by flattening the pH change near the equivalence point, and switching indicator colors (Option B) doesn't solve the fundamental problem that the sample color obscures any color transition.
Question 3 True / False
A heavily buffered sample requires more titrant than an unbuffered sample of the same total acid content, and the resulting titration curve is flatter near the equivalence point, demanding more precise endpoint location.
TTrue
FFalse
Answer: True
Buffers resist pH change by consuming added base through their conjugate acid components. A highly buffered sample contains a reservoir of protons that must be neutralized before the pH changes sharply — this 'uses up' titrant in the buffer region and produces a gradual pH transition near the equivalence point rather than a steep inflection. This is why potentiometric detection and careful derivative analysis are especially important for buffered samples: a visual indicator might miss or bracket the equivalence point over a large volume range.
Question 4 True / False
Diluting a sample before titration typically improves analytical accuracy by making the pH change at the equivalence point sharper and easier to detect.
TTrue
FFalse
Answer: False
Dilution actually flattens the pH change at the equivalence point, making it harder to detect precisely — especially for weak acid/base systems. The equivalence point pH depends on the concentration of the analyte and product; dilution shifts the equilibrium and reduces the magnitude of the pH spike. For concentrated strong acid/strong base systems, dilution has minimal effect, but for weak acid/base systems and buffered samples, dilution worsens endpoint detection. The misconception reverses the actual effect.
Question 5 Short Answer
Why is rigorous uncertainty evaluation essential in pharmaceutical acid-base titrations, and what are the major sources of uncertainty that must be propagated?
Think about your answer, then reveal below.
Model answer: Pharmaceutical assay regulations (e.g., ±2% of labeled content) impose tight accuracy requirements. Multiple measurement steps each introduce error: the concentration of the standardized titrant (which depends on primary standard purity, weighing precision, and volumetric flask calibration); buret volume readings (graduated uncertainty per reading, doubled because two readings — initial and final — are needed); and the mass of the sample (balance precision). Each source contributes to the combined uncertainty, which is propagated through the stoichiometric calculation. If the total combined uncertainty exceeds the regulatory limit, the method fails even if individual readings appear precise. Demonstrating compliance requires documenting and quantifying each source, not just reporting a single measurement.
The key insight is that analytical chemistry distinguishes accuracy (closeness to truth) from precision (reproducibility), and that regulatory compliance requires demonstrated uncertainty bounds — not just a plausible result. A single measurement, however careful, cannot establish compliance without replication and propagation analysis.