Question

Mercury(I) ions \(\left(\mathrm{Hg}_{2}{ }^{2+}\right)\) can be removed from solution by precipitation with \(\mathrm{Cl}^{-}\). Suppose a solution contains aqueous \(\mathrm{Hg}_{2}\left(\mathrm{NO}_{3}\right)_{2}\). Write complete ionic and net ionic equations to show the reaction of aqueous \(\mathrm{Hg}_{2}\left(\mathrm{NO}_{3}\right)_{2}\) with aqueous sodium chloride to form solid \(\mathrm{Hg}_{2} \mathrm{Cl}_{2}\) and aqueous sodium nitrate.

Short answer

Net ionic equation: \(\mathrm{Hg}_2}^{2+}(aq) + 2\mathrm{Cl}^{-}(aq) \rightarrow \mathrm{Hg}_2\mathrm{Cl}_2(s)\)

Step-by-step solution

Write the balanced molecular equation

First write the molecular equation with the reactants and products using their correct chemical formulas. The reactants are aqueous Mercury(I) nitrate \(\mathrm{Hg}_2(\mathrm{NO}_3)_2\) and aqueous sodium chloride \(\mathrm{NaCl}\). The products are solid Mercury(I) chloride \(\mathrm{Hg}_2\mathrm{Cl}_2\) and aqueous sodium nitrate \(\mathrm{NaNO}_3\). The balanced molecular equation is: \[\mathrm{Hg}_2(\mathrm{NO}_3)_2(aq) + 2\mathrm{NaCl}(aq) \rightarrow \mathrm{Hg}_2\mathrm{Cl}_2(s) + 2\mathrm{NaNO}_3(aq)\]

Write the complete ionic equation

Next, separate the aqueous compounds into their respective ions, keeping the solid compound intact. The complete ionic equation is: \[\mathrm{Hg}_2}^{2+}(aq) + 2\mathrm{NO}_3^{-}(aq) + 2\mathrm{Na}^{+}(aq) + 2\mathrm{Cl}^{-}(aq) \rightarrow \mathrm{Hg}_2\mathrm{Cl}_2(s) + 2\mathrm{Na}^{+}(aq) + 2\mathrm{NO}_3^{-}(aq)\]

Write the net ionic equation

Cancel out the spectator ions that do not participate in the reaction. The spectator ions in this case are the sodium \(\mathrm{Na}^{+}\) and nitrate \(\mathrm{NO}_3^{-}\) ions since they appear on both sides. The net ionic equation features only the ions and compounds that participate in the reaction: \[\mathrm{Hg}_2}^{2+}(aq) + 2\mathrm{Cl}^{-}(aq) \rightarrow \mathrm{Hg}_2\mathrm{Cl}_2(s)\]

Key concepts

Chemical Precipitation

Chemical precipitation is a process where a solid, known as the precipitate, forms within a solution as the result of a chemical reaction. When certain ions in solution combine, they can create an insoluble compound that falls out of the solution.

In the provided exercise, the reaction between aqueous Mercury(I) nitrate and sodium chloride results in the formation of solid Mercury(I) chloride. This substance is insoluble in water and, hence, precipitates out. By understanding the conditions which favor the formation of a precipitate, such as the concentration of the ions and the temperature, it's possible to predict and manipulate the outcomes of reactions in both laboratory and industrial settings.

Chemical precipitation is crucial in fields such as water treatment, where harmful ions are removed from water to make it safe for consumption or release into the environment.

Ionic Equations

Ionic equations provide a more detailed view of a chemical reaction by showing the ions involved as opposed to whole compounds. When compounds in a reaction are aqueous, they dissociate into cations and anions, and ionic equations represent this separation.

Dissociated ions are free to interact with each other, and these interactions may lead to the formation of new compounds, which can be either soluble or insoluble in the solution. For the reaction in our exercise, the complete ionic equation breaks down soluble ionic compounds into their constituent ions, highlighting how the Mercury(I) ions combine with chloride ions to form Mercury(I) chloride as a precipitate.

Writing these equations helps students visualize the actual reactive species in the solution and understand the changes that occur during the reaction process. It's a valuable tool in predicting the outcomes and products of the reactions.

Spectator Ions

Spectator ions are ions in a solution that do not participate in the chemical reaction and remain unchanged. They can be identified in a reaction as the ions that appear on both sides of a complete ionic equation.

In the step-by-step solution, sodium \(\mathrm{Na}^{+}\) and nitrate \(\mathrm{NO}_3^{-}\) ions are spectators, as they do not contribute to the formation of the precipitate and remain dissolved in the water. The net ionic equation omits these ions, focusing on the ions that do contribute to the reaction — specifically, the ions that change partners, forming the precipitate. By recognizing spectator ions, students can simplify equations and concentrate on the essence of the reaction.

Understanding spectator ions is important in stoichiometry and predicting reaction products, as they give us deeper insight into which parts of the reaction are active and which are not.

Solubility Rules

Solubility rules are guidelines to predict whether an ionic compound will dissolve in water (being soluble) or form a precipitate (being insoluble).

For example, most chloride salts are soluble, but there are exceptions such as silver, mercury, and lead chlorides, which are not soluble and thus precipitate from solution. In our exercise, Mercury(I) chloride, \(\mathrm{Hg}_2\mathrm{Cl}_2\), forms as a precipitate, illustrating that it is an exception to the general rule for chlorides.

Knowing these rules helps in writing net ionic equations and is essential in making predictions about the outcomes of reactions. For students, memorizing these rules can be a challenging but necessary part of learning chemistry, as it allows for quick identification of the possible reaction products and understanding of how different compounds will behave when mixed in aqueous solutions.