Chemistry

Problem

0.875\,\text{g}

sample of a commercial fertilizer is dissolved in water and titrated with

0.200\,\text{M}

NaOH(aq)

. The titration requires

35.60\,\text{mL}

of the base to reach the stoichiometric point. The only acid-reactive component in the fertilizer is assumed to be ammonium dihydrogen phosphate,

NH_4H_2PO_4

. No other acid–base reactions occur. (a) Write the balanced molecular equation for the complete neutralization of

NH_4H_2PO_4

with aqueous sodium hydroxide. (b) Calculate the number of moles of

NaOH

delivered from the buret. (c) Determine the mass of

NH_4H_2PO_4

that reacted. (d) Calculate the percent by mass of

NH_4H_2PO_4

in the fertilizer sample.

Intuitive explanation

This is an acid–base titration where you're using a known base concentration and volume to back-calculate the amount of acid in the original sample. The key is recognizing that ammonium dihydrogen phosphate is diprotic (it has two replaceable hydrogens: one from $NH_4^+$ and one from $H_2PO_4^-$ ), so each mole of $NH_4H_2PO_4$ consumes 2 moles of $NaOH$ . Use the titration data to find moles of base, convert to moles of the phosphate compound, then to grams and finally mass percentage.

Intuitive explanation

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