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## Q. 4.12

A precipitate forms when aqueous solutions of ammonium sulfate and barium chloride are mixed. Write the net ionic equation for the reaction. Is this the same net ionic equation as in Sample Exercise 4.11?

## Verified Solution

Collect and Organize Mixing two solutions causes an insoluble compound to precipitate. Using the names of the reactants, we are asked to identify the precipitate. We then need to write the net ionic equation and compare it with the one from the previous Sample Exercise.

Analyze The two salts are in solution, so they both must dissociate in water. This is consistent with the solubility rules in Table 4.4

 Table 4.4 Solubility Rules for Common Ionic Compounds in Water All compounds containing the following ions are soluble:  Cations: Group 1 ions (alkali metals) and $NH_{4}^{+}$  Anions: $NO_{3}^{-}$ and $CH_{3}COO^{-}$  (acetate) Compounds containing the following anions are soluble except as noted: $Cl^{-}, Br^{-}$, and $I^{-}$, except those of $Ag^{+}, Cu^{+}, Hg_{2}^{2+}$, and $Pb^{2+}$ $SO_{4}^{2-}$, except those of $Ba^{2+}, Ca^{2+}, Hg_{2}^{2+}, Pb^{2+}$, and $Sr^{2+}$ Insoluble compounds include the following: All hydroxides except those of group 1 cations and $Ca(OH)_{2}, Sr(OH)_{2}$, and $Ba(OH)_{2}$ All sulfides except those of group 1 cations and $NH_{4}^{+}, Cas, SrS$. and $Bas$ All carbonates except those of group 1 cations and $NH_{4}^{+}$ All phosphates except those of group 1 cations and $NH_{4}^{+}$ Most fluorides, though not those of group 1 cations and $NH_{4}^{+}$

or Table 4.5.

 TABLE 4.5 Solubility Matrix for Common Ionic Compounds in Water CATION Group 1 $or NH_{4}^{+}$ $Ca^{2+} or Sr^{2+} or Ba^{2+}$ $Cu^{+} or Ag^{+}$ $Hg_{2}^{2+} or Pb^{2+}$ ANION $NO_{3}^{-} or CH_{3}COO^{-}$ Soluble Soluble Soluble Soluble $Cl^{-} or Br^{-} or I^{-}$ Soluble Soluble Insoluble Insoluble $OH^{-} or S^{2-}$ Soluble Soluble Insoluble Insoluble $SO_{4}^{2-}$ Soluble Insoluble Soluble Insoluble $F^{-} or CO_{3}^{2-} or PO_{4}^{3-}$ Soluble Insoluble Insoluble Insoluble

We need to determine which combinations of the ions produce insoluble solids. First, we need to write correct formulas of the salts. Then we can check Table 4.4 or Table 4.5 to see which of the possible combinations of ions are insoluble in water.

Solve

Solution 1 consists of ammonium ions and sulfate ions: $NH_{4} ^{+}(aq)$ and $SO_{4}^{2-}(aq)$
Solution 2 consists of barium ions and chloride ions: $Ba^{2+}(aq)$ and $Cl^{-}(aq)$

The new combinations are $BaSO_{4}$ and $NH_{4}Cl$. Table 4.4 indicates that all ammonium salts are soluble, so the ammonium and chloride ions remain dissolved as solvated ions. Barium sulfate is insoluble, however, and that salt precipitates. The overall ionic equation describes the species in solution:

$\sout{2 NH_{4} ^{+}(aq)} +SO_{4}^{2-}(aq) + Ba^{2+}(aq) + \sout{2 Cl^{-}(aq)} → \sout{2 NH_{4} ^{+}(aq)} + \sout{2 Cl^{-}(aq)} + BaSO_{4}(s)$

We can simplify the overall ionic equation by eliminating the spectator ions to yield the net ionic equation, which describes the formation of the precipitate:

$Ba^{2+}(aq)+SO_{4}^{2-}(aq) → BaSO_{4}(s)$

This net ionic equation is not the same as for the reaction between barium hydroxide and sulfuric acid in Sample Exercise 4.11, even though both reactions form the same precipitate.

Think About It To identify whether a precipitation reaction occurs when you mix two solutions containing ions, check whether an insoluble compound is formed when the ions in solution collide with one another. In this exercise, sulfate ions originally in the ammonium sulfate solution collide with barium ions originally in the barium chloride solution, forming insoluble $BaSO_{4}$. The other ions, $NH_{4}^{+}$ and $Cl^{-}$, remain in solution.