= Acetone iodination rate-law experiment
<Acetone> reacts with <iodine> to form <iodoacetone>, <iodide>, and a <hydrogen ion>:
$$
\mathrm{CH_3COCH_3}+\mathrm I_2\longrightarrow
\mathrm{CH_3COCH_2I}+\mathrm H^++\mathrm I^-.
$$
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 $2\,{}^\circ\mathrm C$.<Chemistry M01B Laboratory Manual, version 4.0 by Moorpark College Department of Chemistry/Experiment 4: Chemical Kinetics: The Iodination of Acetone>[Chemistry M01B Laboratory Manual, version 4.0 by Moorpark College Department of Chemistry; Version 4.0, Experiment 4, printed pp. 17–18: Procedure and Table 1]{ref}
The stock dilutions give $[\mathrm{acetone}]_0=0.800$, $[\mathrm H^+]_0=0.200$, and $[\mathrm I_2]_0=0.00100\,\mathrm{mol\,L^{-1}}$. 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.<Chemistry M01B Laboratory Manual, version 4.0 by Moorpark College Department of Chemistry/Experiment 4: Chemical Kinetics: The Iodination of Acetone>[Chemistry M01B Laboratory Manual, version 4.0 by Moorpark College Department of Chemistry; Printed pp. 15–16, iodine independence and equations (3)–(7); p. 18, Table 1, mixtures A–D]{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 $m=\ln(r_B/r_A)/\ln2$; 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.
Back to article page
