A student carried out two titrations using an NaOH solution...

A student carried out two titrations using an $\mathrm{NaOH}$ solution of unknown concentration in the buret. In one titration she weighed out $0.2458 \mathrm{g}$ of $\mathrm{KHP}$ (see page 144 ) and transferred it to an Erlenmeyer flask. She then added $20.00 \mathrm{mL}$ of distilled water to dissolve the acid. In the other titration she weighed out $0.2507 \mathrm{g}$ of KHP but added 40.00 mL of distilled water to dissolve the acid. Assuming no experimental error, would she obtain the same result for the concentration of the NaOH solution?

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Key Concepts

Acid-Base Titration

Acid-base titration involves the gradual addition of a solution of known concentration (titrant) to a solution containing an analyte (which can be acidic or basic) until the reaction reaches its equivalence point, where stoichiometrically equivalent amounts of acid and base have reacted. This technique is fundamental in analytical chemistry for accurately determining the concentration of an unknown solution.

Stoichiometry of Neutralization

The stoichiometry of a neutralization reaction is based on the balanced chemical equation showing the reaction between an acid and a base, typically in a 1:1 molar ratio for monoprotic acids such as KHP (potassium hydrogen phthalate) and NaOH. Accurate titration relies on this stoichiometric relationship to calculate the molarity of the unknown solution from the amount of reactant consumed in reaching the equivalence point.

Dilution Effects in Titrations

The addition of water to dissolve the analyte does not affect the titration results because dilution alters the concentration of the analyte but not the total number of moles present. During titration, the reaction is based solely on the moles of reactant available, so varying the solvent volume has no impact on the calculated concentration of the unknown titrant as long as complete dissolution is achieved.

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