Question

Using the general solubility rules given in Table 4.1, name three reagents that would form precipitates with each of the following ions in aqueous solution. Write the net ionic equation for each of your suggestions. a. chloride ion b. calcium ion c. iron(III) ion d. sulfate ion e. mercury(I) ion, \(\mathrm{Hg}_{2}{ }^{2+}\) f. silver ion

Short answer

a. Chloride ion: Ag⁺, Pb²⁺, Hg₂²⁺ Net ionic equations: \( Ag^+ + Cl^- \rightarrow AgCl↓ \) \( Pb^{2+} + 2Cl^- \rightarrow PbCl_2↓ \) \( Hg_2^{2+} + 2Cl^- \rightarrow Hg_2Cl_2↓\) b. Calcium ion: CO₃²⁻, PO₄³⁻, C₂O₄²⁻ Net ionic equations: \( Ca^{2+} + CO_3^{2-} \rightarrow CaCO_3↓ \) \( 3Ca^{2+} + 2PO_4^{3-} \rightarrow Ca_3(PO_4)_2↓ \) \( Ca^{2+} + C_2O_4^{2-} \rightarrow CaC_2O_4↓\) c. Iron(III) ion: OH⁻, PO₄³⁻, S²⁻ Net ionic equations: \( Fe^{3+} + 3OH^- \rightarrow Fe(OH)_3↓ \) \( 2Fe^{3+} + 3PO_4^{3-} \rightarrow Fe_2(PO_4)_3↓ \) \( Fe^{3+} + 3S^{2-} \rightarrow Fe_2S_3↓\) d. Sulfate ion: Ba²⁺, Pb²⁺, Ca²⁺ Net ionic equations: \( Ba^{2+} + SO_4^{2-} \rightarrow BaSO_4↓ \) \( Pb^{2+} + SO_4^{2-} \rightarrow PbSO_4↓ \) \( Ca^{2+} + SO_4^{2-} \rightarrow CaSO_4↓\) e. Mercury(I) ion, Hg₂²⁺: Cl⁻, Br⁻, I⁻ Net ionic equations: \( Hg_2^{2+} + 2Cl^- \rightarrow Hg_2Cl_2↓ \) \( Hg_2^{2+} + 2Br^- \rightarrow Hg_2Br_2↓ \) \( Hg_2^{2+} + 2I^- \rightarrow Hg_2I_2↓\) f. Silver ion: Cl⁻, Br⁻, I⁻ Net ionic equations: \( Ag^+ + Cl^- \rightarrow AgCl↓ \) \( Ag^+ + Br^- \rightarrow AgBr↓ \) \( Ag^+ + I^- \rightarrow AgI↓\)

Step-by-step solution

Identify possible reagents forming precipitates with chloride ions

Using the general solubility rules, we can find that silver, lead, and mercury(I) ions will form precipitates with chloride ions. These are Ag⁺, Pb²⁺, and Hg₂²⁺ ions.

Write the net ionic equations for precipitation reactions

\ \[ Ag^+ + Cl^- \rightarrow AgCl↓ \] \[ Pb^{2+} + 2Cl^- \rightarrow PbCl_2↓ \] \[ Hg_2^{2+} + 2Cl^- \rightarrow Hg_2Cl_2↓\] #b. Calcium ion#

Identify possible reagents forming precipitates with calcium ions

Using the general solubility rules, we can find that carbonate, phosphate, and oxalate ions will form precipitates with calcium ions. These are CO₃²⁻, PO₄³⁻, and C₂O₄²⁻ ions.

Write the net ionic equations for precipitation reactions

\ \[ Ca^{2+} + CO_3^{2-} \rightarrow CaCO_3↓ \] \[ 3Ca^{2+} + 2PO_4^{3-} \rightarrow Ca_3(PO_4)_2↓ \] \[ Ca^{2+} + C_2O_4^{2-} \rightarrow CaC_2O_4↓\] #c. Iron(III) ion#

Identify possible reagents forming precipitates with iron(III) ions

Using the general solubility rules, we can find that hydroxide, phosphate, and sulfide ions will form precipitates with iron(III) ions. These are OH⁻, PO₄³⁻, and S²⁻ ions.

Write the net ionic equations for precipitation reactions

\ \[ Fe^{3+} + 3OH^- \rightarrow Fe(OH)_3↓ \] \[ 2Fe^{3+} + 3PO_4^{3-} \rightarrow Fe_2(PO_4)_3↓ \] \[ Fe^{3+} + 3S^{2-} \rightarrow Fe_2S_3↓\] #d. Sulfate ion#

Identify possible reagents forming precipitates with sulfate ions

Using the general solubility rules, we can find that barium, lead, and calcium ions will form precipitates with sulfate ions. These are Ba²⁺, Pb²⁺, and Ca²⁺ ions.

Write the net ionic equations for precipitation reactions

\ \[ Ba^{2+} + SO_4^{2-} \rightarrow BaSO_4↓ \] \[ Pb^{2+} + SO_4^{2-} \rightarrow PbSO_4↓ \] \[ Ca^{2+} + SO_4^{2-} \rightarrow CaSO_4↓\] #e. Mercury(I) ion, Hg₂²⁺

Identify possible reagents forming precipitates with mercury(I) ions

Using the general solubility rules, we can find that chloride, bromide, and iodide ions will form precipitates with mercury(I) ions. These are Cl⁻, Br⁻, and I⁻ ions.

Write the net ionic equations for precipitation reactions

\ \[ Hg_2^{2+} + 2Cl^- \rightarrow Hg_2Cl_2↓ \] \[ Hg_2^{2+} + 2Br^- \rightarrow Hg_2Br_2↓ \] \[ Hg_2^{2+} + 2I^- \rightarrow Hg_2I_2↓\] #f. Silver ion#

Identify possible reagents forming precipitates with silver ions

Using the general solubility rules, we can find that chloride, bromide, and iodide ions will form precipitates with silver ions. These are Cl⁻, Br⁻, and I⁻ ions.

Write the net ionic equations for precipitation reactions

\ \[ Ag^+ + Cl^- \rightarrow AgCl↓ \] \[ Ag^+ + Br^- \rightarrow AgBr↓ \] \[ Ag^+ + I^- \rightarrow AgI↓\]

Key concepts

Precipitation Reactions

In chemistry, precipitation reactions occur when two solutions containing soluble salts are mixed, and an insoluble salt is formed as a result. This process occurs because of the limited solubility of certain ion combinations. When ions combine to create a compound that does not dissolve in water, a solid known as a precipitate forms and falls out of the solution. For instance, when calcium ions \(Ca^{2+}\) and carbonate ions \(CO_3^{2-}\) combine, calcium carbonate \(CaCO_3\) is formed, which is a white solid precipitate.
Precipitation reactions are essential in various environmental and industrial processes, such as water softening and mineral extraction. Additionally, they are used in laboratory analyses to isolate specific ions or to determine the concentration of a particular substance. Understanding solubility rules is crucial to predict which ions will form precipitates, allowing the chemist to control the outcome of these reactions.

Net Ionic Equations

Net ionic equations simplify chemical equations by only showing the ions and molecules directly involved in the reaction. They exclude spectator ions, which are ions that do not participate in the chemical change and remain unchanged in the solution. For example, in the reaction between silver ions \(Ag^+\) and chloride ions \(Cl^-\) to form silver chloride \(AgCl\), the net ionic equation would be: \[ Ag^+ + Cl^- \rightarrow AgCl\downarrow \]
In this equation, the silver and chloride ions are the only species changing their state, forming an insoluble salt, while any other ions (like sodium or nitrate) are omitted. Net ionic equations are valuable for emphasizing the actual chemical change taking place in a reaction and help beginners focus on the key components involved.

This approach aids students in understanding the fundamentals of chemical reactions and gaining clarity on which ions interact to form precipitates.

Aqueous Solutions

An aqueous solution is a solution where water is the solvent. In such a solution, substances dissolve in water, resulting in free-moving ions that can engage in chemical reactions. Most reactions in water, such as precipitation reactions, depend on the properties of aqueous solutions to facilitate ion interaction.
Water's polarity enables it to efficiently dissolve ionic substances. When, for example, table salt \(NaCl\) dissolves, its ions are separated and distributed in the water. This separation allows them to interact with other solutes dissolved in the solution.Aqueous solutions are fundamental in the study of chemistry because they represent the environment in which most biochemical processes occur. This understanding allows scientists and students to model and predict reactions that are analogous to those occurring in living organisms.

Chemical Reagents

Chemical reagents are substances used to cause a desired chemical reaction. In the context of precipitation reactions, reagents are the ions in solution that combine to form the insoluble product. Choosing the right reagents is essential to drive the reaction towards the formation of the desired precipitate.
This selection process is informed by the solubility rules, which guide chemists in predicting which combinations of ions will lead to precipitate formation. Common chemical reagents include acids, bases, and salts, among others. An understanding of chemical reagents is crucial for effective experimental design, as it impacts the reaction pathway and product yield.
In educational settings, working with reagents helps students grasp the practical aspects of chemistry and the mechanisms by which chemical reactions occur. This foundational knowledge is applied beyond academia, such as in pharmaceuticals and manufacturing, where precise reaction outcomes are necessary.