Glycine titration

Adding sodium hydroxide to acidified glycine successively removes its two titratable protons. Two broad buffer solution regions appear around the two pKa values, separated by a steep pH rise near the isoelectric point. This is useful evidence for two separated acid-base chemistry steps even though the changing chemical species are invisible. In the approximation that added hydroxide only changes glycine protonation, $b=1-\bar{z}$ is the added base amount of substance per initial amount of substance of glycine. Then $b\approx 0.5$, $1$ and $1.5$ correspond to $\mathrm{p}{K}_{a1}$, $\mathrm{pI}$ and $\mathrm{p}{K}_{a2}$; endpoints are not exactly pure chemical species. ^[ref]

For the solid-curve model, start with $n_G=0.00200\mathrm{mol}$ of glycine in $V_0=0.0200\mathrm{L}$ and exactly $n_G$ of fully dissociated hydrochloric acid; this fixes the chloride inventory, but does not force every glycine molecule to remain $G^{+}$. Let the sodium hydroxide concentration be $C_B=0.100\mathrm{mol}{\mathrm{L}}^{-1}$ and $b=n_{\mathrm{NaOH}}/n_G$. Assume volume is additive and no other ions are present; stipulate $K_w=10^{-14}$ for self-ionization of water at a temperature of $25{}^{\circ }\mathrm{C}$. ThenBalancing all positive and negative electric charges requiresSolving this equation gives the solid curve. It remains finite at $b=0$ and beyond $b=2$, where ignoring free hydronium and hydroxide would fail. At the first pKa, the full model gives $b\approx 0.436$, rather than $0.5$: appreciable free hydronium makes that half-equivalent shortcut inaccurate. At the second pKa, $b\approx 1.502$. Known acid inventory, dilution, and thermodynamic activity assumptions are therefore part of any quantitative interpretation, not optional corrections inferred from the curve's shape.

The Amrita Virtual Lab procedure uses $0.1\mathrm{mol}{\mathrm{L}}^{-1}$ glycine, hydrochloric acid and sodium hydroxide, two burettes, a pH meter, beaker and stirrer. A pipette transfers $20\mathrm{mL}$ of glycine into a $100\mathrm{mL}$ beaker. Calibrate the meter with standard buffer solutions at pH $4$, $7$ and $10$, then record the starting pH. Add hydrochloric acid in $0.3\mathrm{mL}$ increments, recording pH after each addition until $1.6$. Rinse the electrode in distilled water, take a separate fresh $20\mathrm{mL}$ portion, and add sodium hydroxide in $0.3\mathrm{mL}$ increments until pH $12.5$. Its observed readout is paired added volume and pH, plotted as an acid-base titration curve; it does not measure each glycine form directly. ^[ref]

These are two titrations from the starting glycine chemical solution, rather than the single acidified-sample protocol assumed by the solid curve in glycine titration. Retain each portion's initial volume and the signed added acid/base amount of substance; do not concatenate raw burette readings and assume the result begins at $b=0$. For this procedure's $0.00200\mathrm{mol}$ of glycine, each $0.3\mathrm{mL}$ addition supplies $0.000030\mathrm{mol}$ of acid or base, a $0.015$ increment relative to the glycine inventory. This calculated resolution is not a reported measurement.

Useful additional controls are a fresh repeat portion, independently checked acid/base concentrations, a water blank, and a calibration-check buffer solution after the run. Stir without striking the glass electrode, keep its sensing region immersed and wait for stable readings; record temperature and background salts. Reduce addition size near a steep pH change. These are suggested controls for testing the model, beyond the published steps. A water blank can reveal contamination but must not be subtracted directly from pH, which is a logarithm.

The standards at pH $4$, $7$ and $10$ do not bracket either endpoint of this protocol. Extreme readings therefore require independently checked instrument performance, and alkaline samples are sensitive to atmospheric carbon dioxide uptake. Electrode drift, incomplete mixing, volume error, dilution and changing ionic strength can shift apparent pKa values. A titration constrains the acid/base model; it cannot establish chirality or separate every microscopic protonation form. The two original diagrams show calculated predictions to compare with measured points, not reconstructed experimental data.