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Chapter 17: Problem 55

Sodium phosphate, \(\mathrm{Na}_{3} \mathrm{PO}_{4},\) is made commercially by first neutralizing phosphoric acid with sodium carbonate to obtain \(\mathrm{Na}_{2} \mathrm{HPO}_{4}\). The \(\mathrm{Na}_{2} \mathrm{HPO}_{4}\) is further neutralized to \(\mathrm{Na}_{3} \mathrm{PO}_{4}\) with \(\mathrm{NaOH}\) (a) Write net ionic equations for these reactions. (b) \(\mathrm{Na}_{2} \mathrm{CO}_{3}\) is a much cheaper base than is \(\mathrm{NaOH}\) Why do you suppose that \(\mathrm{NaOH}\) must be used as well as \(\mathrm{Na}_{2} \mathrm{CO}_{3}\) to produce \(\mathrm{Na}_{3} \mathrm{PO}_{4} ?\)

Short Answer

Expert verified
The net ionic reactions are:(a) \(\mathrm{H}_{3}\mathrm{PO}_{4} + \mathrm{Na}_{2}\mathrm{CO}_{3} \rightarrow 2\mathrm{Na}_{2}\mathrm{HPO}_{4} + \mathrm{CO}_{2} + \mathrm{H}_{2}\mathrm{O}\) (b) \(\mathrm{Na}_{2}\mathrm{HPO}_{4} + \(\mathrm{NaOH} \rightarrow 2\mathrm{Na}_{3} \mathrm{PO}_{4} + \mathrm{H}_{2}\mathrm{O}\). The use of both \(\mathrm{Na}_{2} \mathrm{CO}_{3}\) and \(\mathrm{NaOH}\) in the production of \(\mathrm{Na}_{3} \mathrm{PO}_{4}\) is necessary as \(\mathrm{Na}_{2} \mathrm{CO}_{3}\) is only capable of providing two \(Na^{+}\) ions and the third \(Na^{+}\) is provided by \(\mathrm{NaOH}\).

Step by step solution

01

Writing net ionic equation for \(\mathrm{Na}_{3} \mathrm{PO}_{4}\) production

The reaction of phosphoric acid with sodium carbonate can be represented as follows: \(\mathrm{H}_{3}\mathrm{PO}_{4} + \mathrm{Na}_{2}\mathrm{CO}_{3} \rightarrow \mathrm{Na}_2\mathrm{HPO}_{4} + \mathrm{CO}_{2} + \mathrm{H}_{2}\mathrm{O}\). When this reaction equation is balanced, it becomes: \(\mathrm{H}_{3}\mathrm{PO}_{4} + \mathrm{Na}_{2}\mathrm{CO}_{3} \rightarrow 2\mathrm{Na}_{2}\mathrm{HPO}_{4} + \mathrm{CO}_{2} + \mathrm{H}_{2}\mathrm{O}\). This equation excludes spectator ions hence, this is a net ionic equation.
02

Writing net ionic equation for \(\mathrm{Na}_{2}\mathrm{HPO}_{4}\) conversion to \(\mathrm{Na}_{3} \mathrm{PO}_{4}\)

This reaction neutralize \(\mathrm{Na}_{2}\mathrm{HPO}_{4}\) to \(\mathrm{Na}_{3} \mathrm{PO}_{4}\) with \(\mathrm{NaOH}\) and represented as \(\mathrm{Na}_{2}\mathrm{HPO}_{4} + \(\mathrm{NaOH} \rightarrow \mathrm{Na}_{3} \mathrm{PO}_{4} + \mathrm{H}_{2}\mathrm{O}\). After balancing this equation it becomes: \(\mathrm{Na}_{2}\mathrm{HPO}_{4} + \(\mathrm{NaOH} \rightarrow 2\mathrm{Na}_{3} \mathrm{PO}_{4} + \mathrm{H}_{2}\mathrm{O}\). Spectator ions are already excluded from this equation, thus this also represents the net ionic equation.
03

Explanation of \(\mathrm{Na}_{2} \mathrm{CO}_{3}\) and \(\mathrm{NaOH}\) usage

\(\mathrm{Na}_{2} \mathrm{CO}_{3}\) is only capable of providing two \(Na^{+}\) ions while the manufacturing of \(\mathrm{Na}_{3} \mathrm{PO}_{4}\) requires three \(Na^{+}\) ions. Therefore, another base with \(Na^{+}\) ion is needed. So, \(\mathrm{NaOH}\) is used in this reaction as it provides an extra \(Na^{+}\) needed to fully neutralize phosphate ion \(\mathrm{(PO_{4}^{-3})}\) to produce \(\mathrm{Na}_{3} \mathrm{PO}_{4}\).

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Most popular questions from this chapter

\(\begin{array}{lll}\text { Given } & 1.00 & \mathrm{L}\end{array}\) of a solution that is \(0.100 \mathrm{M}\) \(\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{COOH}\) and \(0.100 \mathrm{M} \mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{COO}\) (a) Over what pH range will this solution be an effective buffer? (b) What is the buffer capacity of the solution? That is, how many millimoles of strong acid or strong base can be added to the solution before any significant change in pH occurs?

Explain whether the equivalence point of each of the following titrations should be below, above, or at pH 7: (a) \(\mathrm{NaHCO}_{3}(\text { aq) titrated with } \mathrm{NaOH}(\mathrm{aq}) ; \text { (b) } \mathrm{HCl}(\mathrm{aq})\) titrated with \(\mathrm{NH}_{3}(\mathrm{aq}) ;\) (c) KOH(aq) titrated with HI(aq).

Is a solution of sodium dihydrogen citrate, \(\mathrm{NaH}_{2} \mathrm{Cit}\) likely to be acidic, basic, or neutral? Explain. Citric \(\mathrm{acid}, \mathrm{H}_{3} \mathrm{Cit}, \mathrm{is} \mathrm{H}_{3} \mathrm{C}_{6} \mathrm{H}_{5} \mathrm{O}_{7}\)

What stoichiometric concentration of the indicated substance is required to obtain an aqueous solution with the pH value shown: (a) \(\mathrm{Ba}(\mathrm{OH})_{2}\) for \(\mathrm{pH}=11.88 ;(\mathrm{b})\) \(\mathrm{CH}_{3} \mathrm{COOH}\) in \(0.294 \mathrm{M} \mathrm{NaCH}_{3} \mathrm{COO}\) for \(\mathrm{pH}=4.52 ?\)

Calculate the pH of the buffer formed by mixing equal volumes \(\left[\mathrm{C}_{2} \mathrm{H}_{5} \mathrm{NH}_{2}\right]=1.49 \mathrm{M} \quad\) with \(\quad\left[\mathrm{HClO}_{4}\right]=\) 1.001 M. \(K_{\mathrm{b}}=4.3 \times 10^{-4}\)
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