Write net ionic equations to explain the formation of (a) a white precipitate when solutions of calcium sulfate and sodium carbonate are mixed. (b) two different precipitates formed when solutions of iron(III) sulfate and barium hydroxide are mixed.

Precipitation reactions occur when two soluble salt solutions are mixed together to form one or more insoluble products, known as precipitates. In our textbook exercise, two examples are given:

(a) When aqueous solutions of calcium sulfate and sodium carbonate are combined, they react to form a white precipitate of calcium carbonate. This reaction occurs because calcium carbonate is not soluble in water and thus precipitates out of the solution.
(b) In the second reaction, when aqueous solutions of iron(III) sulfate and barium hydroxide mix, iron(III) hydroxide and barium sulfate form. These compounds also precipitate as they are not soluble in water.

Precipitation reactions have numerous practical applications, including water softening, waste treatment, and in analytical chemistry as a means to isolate and purify certain ions.

Understanding solubility rules is essential for predicting the outcome of chemical reactions and writing net ionic equations. Solubility rules are guidelines that help determine whether a compound will dissolve in water (soluble) or not (insoluble).

Some basic rules are as follows:

• Most nitrate (\(NO_3^-\)) salts are soluble.
• Salts containing alkali metal ions (Li+, Na+, K+, Cs+, Rb+) and the ammonium ion (\(NH_4^+\)) are typically soluble.
• Chloride (\(Cl^-\)), bromide (\(Br^-\)), and iodide (\(I^-\)) salts are soluble except when paired with silver (\(Ag^+\)), lead (\(Pb^{2+}\)), or mercury (\(Hg_2^{2+}\)).
• Most sulfate (\(SO_4^{2-}\)) salts are soluble, with exceptions like calcium sulfate (\(CaSO_4\)), barium sulfate (\(BaSO_4\)), and lead sulfate (\(PbSO_4\)).
• Carbonates (\(CO_3^{2-}\)), phosphates (\(PO_4^{3-}\)), sulfides (\(S^{2-}\)), and hydroxides (\(OH^-\)) are generally insoluble, with some exceptions for alkali metals and ammonium.

By applying these rules to the reactants in a chemical equation, one can predict whether a precipitation reaction will occur and what the precipitate will be.

Chemical equations are symbolic representations of chemical reactions. They show the reactants transforming into products, while preserving mass and charge balance. It is crucial that chemical equations be balanced, meaning they have the same number of each type of atom on both sides of the equation.

To further refine our understanding of chemical reactions, we break them down into total ionic and net ionic equations:

Total Ionic Equations

These equations split soluble ionic compounds into their constituent ions, making it easier to identify the species actually involved in the reaction. This is helpful when predicting the formation of precipitates. For insoluble compounds or weak electrolytes, they remain in their molecular form in total ionic equations.

Net Ionic Equations

Net ionic equations eliminate spectator ions, which do not participate in the reaction. What remains are the ions responsible for the formation of the precipitate. Writing net ionic equations focuses on the essentials of a chemical reaction, making them particularly useful in illustrating precipitation reactions, as shown in our textbook exercise examples.