 Acetone iodination rate-law experiment

Acetone reacts with iodine to form iodoacetone, iodide, and a hydrogen ion:In water, ${\mathrm{H}}^{+}$ denotes a solvated hydrogen ion; added hydrochloric acid provides acidic conditions.

Moorpark College's protocol mixes acetone, hydrochloric acid and water from burets in matched test tubes; pipet in iodine, swirl and time colour loss against white paper. Use water, then reacted solution, as references. Duplicate each mixture; keep total volume $5.00\mathrm{mL}$ and run temperatures within $2{}^{\circ }\mathrm{C}$.^[ref]

The stock dilutions give $[\mathrm{acetone}{]}_0=0.800$, $[{\mathrm{H}}^{+}{]}_0=0.200$, and $[{\mathrm{I}}_2{]}_0=0.00100\mathrm{mol}{\mathrm{L}}^{-1}$. B and C separately double acetone and acid; D halves iodine. The manual supplies no measured times; it states zero partial order of reaction in iodine and asks students to determine the other exponents.^[ref]

The endpoint measurement is a time $\tau$ to an operational colour threshold, not direct observation of a reaction intermediate. If the iodine disappearance rate $r$ is approximately constant until that endpoint, $r\approx [{\mathrm{I}}_2{]}_0/\tau$. Halving $[{\mathrm{I}}_2{]}_0$ at fixed other concentrations then approximately halves $\tau$ without changing $r$. A longer or shorter endpoint time alone is not a comparison of reaction rates unless the starting concentration is included. For an observed power-form rate law, doubling only acetone gives a rate ratio $2^m$ and hence $m=\ln (r_B/r_A)/\ln 2$; doubling only acid similarly estimates its own partial order of reaction. Those are inference formulas, not asserted experimental results.

A useful proposed reaction mechanism would explain how iodine is consumed after an earlier transformation sets the throughput, making the observed reaction rate insensitive to iodine over a tested range. The observed zero exponent would contradict treating this net chemical equation as a single elementary reaction with directional rate law $k[\mathrm{acetone}][{\mathrm{I}}_2]$. It would not uniquely prove which reaction intermediate or earlier step is involved. Independently detecting a candidate reaction intermediate would provide another kind of evidence.

The excess-reactant approximation has an explicit scale: the baseline iodine inventory is only $0.00100/0.800=0.125\%$ of the initial acetone inventory, and its net hydrogen ion production is $0.00100/0.200=0.5\%$ of the initial acid inventory, assuming only the stated one-to-one chemical reaction. These are calculated bounds on composition change, not measured reaction orders. Mixing delay, a subjective colour threshold, and temperature drift limit the inference. Repeats and a fixed viewing geometry help; a calibrated absorbance trace could test linearity during conversion instead of relying on an endpoint alone. That extension would require a reference blank and a justified Beer-Lambert law calibration, rather than treating visible colour as an exact concentration. A zero-order reaction approximation in iodine must fail at depletion; it cannot imply consumption after no iodine remains.