Question

Write net ionic equations for each of the following observations. (a) When concentrated \(\mathrm{CaCl}_{2}(\mathrm{aq})\) is added to \(\mathrm{Na}_{2} \mathrm{HPO}_{4}(\mathrm{aq}),\) a white precipitate forms that is \(38.7 \%\) Ca by mass. (b) When a piece of dry ice, \(\mathrm{CO}_{2}(\mathrm{s}),\) is placed in a clear dilute solution of limewater \(\left[\mathrm{Ca}(\mathrm{OH})_{2}(\mathrm{aq})\right]\), bubbles of gas evolve. At first, a white precipitate forms, but then it redissolves.

Short answer

The net ionic equations are, for (a): \(\mathrm{Ca}^{2+} + \mathrm{HPO}_{4}^{2-} \rightarrow \mathrm{CaHPO}_{4}\), and for (b): \(\mathrm{CO}_{2} + \mathrm{Ca}^{2+} + 2\mathrm{OH}^- \rightarrow \mathrm{CaCO}_{3}\) and \(\mathrm{CaCO}_{3} \rightarrow \mathrm{Ca}^{2+} + 2\mathrm{HCO}_{3}^-.

Step-by-step solution

Predict product compounds & Balance equation

For (a), we hypothesize from observations that the products are CaHPO4 (a precipitate) and NaCl. The balanced chemical equation for the reaction: \(\mathrm{CaCl}_{2} + \mathrm{Na}_{2} \mathrm{HPO}_{4} \rightarrow \mathrm{CaHPO}_{4} + 2\mathrm{NaCl}\). In (b), when \(\mathrm{CO}_{2}\) is added to \(\mathrm{Ca} (\mathrm{OH})_{2}\), \(\mathrm{CaCO}_{3}\) initially forms as a precipitate, with water generated as a byproduct. Next, \(\mathrm{CaCO}_{3}\) reacts with excess \(\mathrm{CO}_{2}\) and water to form \(\mathrm{Ca(HCO}_{3})_{2}\), a soluble compound that causes the precipitate to redissolve resulting in: \(\mathrm{CO}_{2} + \(\mathrm{Ca} (\mathrm{OH})_{2} \rightarrow \mathrm{CaCO}_{3} + \mathrm{H}_{2}\mathrm{O}\) and \(\mathrm{CaCO}_{3} + \mathrm{CO}_{2}+ \mathrm{H}_{2}\mathrm{O} \rightarrow \mathrm{Ca(HCO}_{3})_{2}\).

Write the total ionic equation

This is achieved by breaking all the strong electrolytes into their ions: For (a) the equation is \(\mathrm{Ca}^{2+} + 2\mathrm{Cl}^- + 2\mathrm{Na}^+ + \mathrm{HPO}_{4}^{2-} \rightarrow \mathrm{CaHPO}_{4} + 2\mathrm{Na}^+ + 2\mathrm{Cl}^-\). As for (b), the two reactions can be written as \(\mathrm{CO}_{2} + \mathrm{Ca}^{2+} + 2\mathrm{OH}^- \rightarrow \mathrm{CaCO}_{3} + 2\mathrm{OH}^-\) and \(\mathrm{CaCO}_{3} + \mathrm{CO}_{2}+ \mathrm{H}_{2}\mathrm{O} \rightarrow \mathrm{Ca}^{2+} + 2\mathrm{HCO}_{3}^-\).

Write the net ionic equation

This is achieved by canceling out the spectator ions: For (a) we cancel the \(\mathrm{Na}^{+}\) and \(\mathrm{Cl}^{-}\) ions to get the net ionic equation as \(\mathrm{Ca}^{2+} + \mathrm{HPO}_{4}^{2-} \rightarrow \mathrm{CaHPO}_{4}\). As for (b) the first reaction is written as \(\mathrm{CO}_{2} + \mathrm{Ca}^{2+} + 2\mathrm{OH}^- \rightarrow \mathrm{CaCO}_{3}\) (since there are no spectator ions), but in the second reaction, we cancel the \(\mathrm{H}_{2}\mathrm{O}\) and the \(\mathrm{CO}_{2}\) on both sides to give \(\mathrm{CaCO}_{3} \rightarrow \mathrm{Ca}^{2+} + 2\mathrm{HCO}_{3}^-\).

Key concepts

Chemical Reactions

Chemical reactions are the processes in which substances, known as reactants, transform to produce new substances, called products. Each reaction involves a rearrangement of atoms and a change in chemical properties. When writing chemical equations, it is vital to represent this transformation succinctly and accurately. It's helpful to first identify the reactants and predict the products based on known chemical behavior. For instance, when an aqueous solution of calcium chloride reacts with an aqueous solution of sodium phosphate, a double displacement reaction occurs resulting in the formation of a solid precipitate and another salt in solution. In this context, balancing the equation is a critical step that ensures the same number of each type of atom appears on both sides of the equation, obeying the Law of Conservation of Mass.

Understanding the types of chemical reactions, such as synthesis, decomposition, single-replacement, double-replacement, and combustion, can greatly help in predicting the products of a reaction.

Precipitation Reactions

Precipitation reactions occur when two aqueous solutions combine to form one or more insoluble products, known as precipitates. To predict a precipitate in a reaction, one must be familiar with the solubility rules for common ionic compounds. For instance, most chlorides are soluble, but some, like silver chloride, are not. Also, most sulfates are soluble, but calcium sulfate can form a precipitate under certain conditions. These principles guide chemists to anticipate when and what kind of solid will form in a reaction.

Precipitation is a crucial process in chemistry labs and various industrial applications. It aids in the purification of substances, wastewater treatment, and even in medical diagnostics. Writing net ionic equations, like those outlined in the exercise, is essential for focusing on the essence of these reactions by eliminating spectator ions, which do not participate in the actual precipitation process.

Aqueous Solutions

Aqueous solutions are mixtures where water acts as the solvent. They play a central role in chemical reactions, especially in the context of ionic compounds. When such compounds dissolve in water, they dissociate into their constituent ions, which then interact with other ions in the solution. This can lead to a wide variety of chemical phenomena, including the ability to conduct electricity.

In the context of net ionic equations, understanding how substances dissociate in aqueous solutions is key. For instance, knowing that sodium chloride dissolves completely in water to form sodium and chloride ions is fundamental in eliminating spectator ions in net ionic equations. Additionally, the concentration of the solution often influences the reaction's direction and products. For example, the use of concentrated solutions can aid the formation of precipitates which might not occur in dilute solutions. Therefore, the study of aqueous solutions and their properties, such as molarity and acidity, is critical for the accurate prediction and representation of chemical reactions.