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 d. sulfate ion b. calcium ion e. mercury(I) ion, \(\mathrm{Hg}_{2}^{2+}\) c. iron(III) ion f. silver ion

Short answer

For each given ion: Chloride ion (Cl-): 1. Silver nitrate (AgNO3) forms AgCl(s) with net ionic equation: \(Ag^+ + Cl^- \rightarrow AgCl(s)\) 2. Lead(II) nitrate (Pb(NO3)2) forms PbCl2(s) with net ionic equation: \(Pb^{2+} + 2Cl^- \rightarrow PbCl_2(s)\) 3. Mercury(II) nitrate (Hg(NO3)2) forms HgCl2(s) with net ionic equation: \(Hg^{2+} + 2Cl^- \rightarrow HgCl_2(s)\) Sulfate ion (SO4^2-): 1. Barium chloride (BaCl2) forms BaSO4(s) with net ionic equation: \(Ba^{2+} + SO_4^{2-} \rightarrow BaSO_4(s)\) 2. Calcium chloride (CaCl2) forms CaSO4(s) with net ionic equation: \(Ca^{2+} + SO_4^{2-} \rightarrow CaSO_4(s)\) 3. Lead(II) nitrate (Pb(NO3)2) forms PbSO4(s) with net ionic equation: \(Pb^{2+} + SO_4^{2-} \rightarrow PbSO_4(s)\) Calcium ion (Ca^2+): 1. Sodium carbonate (Na2CO3) forms CaCO3(s) with net ionic equation: \(Ca^{2+} + CO_3^{2-} \rightarrow CaCO_3(s)\) 2. Sodium oxalate (Na2C2O4) forms CaC2O4(s) with net ionic equation: \(Ca^{2+} + C_2O_4^{2-} \rightarrow CaC_2O_4(s)\) 3. Sodium sulfate (Na2SO4) forms CaSO4(s) with net ionic equation: \(Ca^{2+} + SO_4^{2-} \rightarrow CaSO_4(s)\) Mercury(I) ion (Hg2^2+): 1. Potassium iodide (KI) forms Hg2I2(s) with net ionic equation: \(Hg_2^{2+} + 2I^- \rightarrow Hg_2I_2(s)\) 2. Potassium bromide (KBr) forms Hg2Br2(s) with net ionic equation: \(Hg_2^{2+} + 2Br^- \rightarrow Hg_2Br_2(s)\) 3. Potassium chloride (KCl) forms Hg2Cl2(s) with net ionic equation: \(Hg_2^{2+} + 2Cl^- \rightarrow Hg_2Cl_2(s)\) Iron(III) ion (Fe^3+): 1. Sodium hydroxide (NaOH) forms Fe(OH)3(s) with net ionic equation: \(Fe^{3+} + 3OH^- \rightarrow Fe(OH)_3(s)\) 2. Sodium sulfide (Na2S) forms Fe2S3(s) with net ionic equation: \(2Fe^{3+} + 3S^{2-} \rightarrow Fe_2S_3(s)\) 3. Sodium phosphate (Na3PO4) forms FePO4(s) with net ionic equation: \(Fe^{3+} + PO_4^{3-} \rightarrow FePO_4(s)\) Silver ion (Ag+): 1. Sodium chloride (NaCl) forms AgCl(s) with net ionic equation: \(Ag^+ + Cl^- \rightarrow AgCl(s)\) 2. Sodium bromide (NaBr) forms AgBr(s) with net ionic equation: \(Ag^+ + Br^- \rightarrow AgBr(s)\) 3. Sodium iodide (NaI) forms AgI(s) with net ionic equation: \(Ag^+ + I^- \rightarrow AgI(s)\)

Step-by-step solution

a. Chloride ion (Cl-)

For the chloride ion, we can predict that the reagents that could form precipitates are: 1. Silver nitrate (AgNO3) - Forms Silver chloride (AgCl) 2. Lead(II) nitrate (Pb(NO3)2) - Forms Lead(II) chloride (PbCl2) 3. Mercury(II) nitrate (Hg(NO3)2) - Forms Mercury(II) chloride (HgCl2) Net ionic equations: 1. \(Ag^+ + Cl^- \rightarrow AgCl(s)\) 2. \(Pb^{2+} + 2Cl^- \rightarrow PbCl_2(s)\) 3. \(Hg^{2+} + 2Cl^- \rightarrow HgCl_2(s)\)

d. Sulfate ion (SO4^2-)

For the sulfate ion, we can predict that the reagents that could form precipitates are: 1. Barium chloride (BaCl2) - Forms Barium sulfate (BaSO4) 2. Calcium chloride (CaCl2) - Forms Calcium sulfate (CaSO4) 3. Lead(II) nitrate (Pb(NO3)2) - Forms Lead(II) sulfate (PbSO4) Net ionic equations: 1. \(Ba^{2+} + SO_4^{2-} \rightarrow BaSO_4(s)\) 2. \(Ca^{2+} + SO_4^{2-} \rightarrow CaSO_4(s)\) 3. \(Pb^{2+} + SO_4^{2-} \rightarrow PbSO_4(s)\)

b. Calcium ion (Ca^2+)

For the calcium ion, we can predict that the reagents that could form precipitates are: 1. Sodium carbonate (Na2CO3) - Forms Calcium carbonate (CaCO3) 2. Sodium oxalate (Na2C2O4) - Forms Calcium oxalate (CaC2O4) 3. Sodium sulfate (Na2SO4) - Forms Calcium sulfate (CaSO4) Net ionic equations: 1. \(Ca^{2+} + CO_3^{2-} \rightarrow CaCO_3(s)\) 2. \(Ca^{2+} + C_2O_4^{2-} \rightarrow CaC_2O_4(s)\) 3. \(Ca^{2+} + SO_4^{2-} \rightarrow CaSO_4(s)\)

e. Mercury(I) ion (Hg2^2+)

For the mercury(I) ion, we can predict that the reagents that could form precipitates are: 1. Potassium iodide (KI) - Forms Mercury(I) iodide (Hg2I2) 2. Potassium bromide (KBr) - Forms Mercury(I) bromide (Hg2Br2) 3. Potassium chloride (KCl) - Forms Mercury(I) chloride (Hg2Cl2) Net ionic equations: 1. \(Hg_2^{2+} + 2I^- \rightarrow Hg_2I_2(s)\) 2. \(Hg_2^{2+} + 2Br^- \rightarrow Hg_2Br_2(s)\) 3. \(Hg_2^{2+} + 2Cl^- \rightarrow Hg_2Cl_2(s)\)

c. Iron(III) ion (Fe^3+)

For the iron(III) ion, we can predict that the reagents that could form precipitates are: 1. Sodium hydroxide (NaOH) - Forms Iron(III) hydroxide (Fe(OH)3) 2. Sodium sulfide (Na2S) - Forms Iron(III) sulfide (Fe2S3) 3. Sodium phosphate (Na3PO4) - Forms Iron(III) phosphate (FePO4) Net ionic equations: 1. \(Fe^{3+} + 3OH^- \rightarrow Fe(OH)_3(s)\) 2. \(2Fe^{3+} + 3S^{2-} \rightarrow Fe_2S_3(s)\) 3. \(Fe^{3+} + PO_4^{3-} \rightarrow FePO_4(s)\)

f. Silver ion (Ag+)

For the silver ion, we can predict that the reagents that could form precipitates are: 1. Sodium chloride (NaCl) - Forms Silver chloride (AgCl) 2. Sodium bromide (NaBr) - Forms Silver bromide (AgBr) 3. Sodium iodide (NaI) - Forms Silver iodide (AgI) Net ionic equations: 1. \(Ag^+ + Cl^- \rightarrow AgCl(s)\) 2. \(Ag^+ + Br^- \rightarrow AgBr(s)\) 3. \(Ag^+ + I^- \rightarrow AgI(s)\)

Key concepts

Precipitate Formation

Precipitates are insoluble solids that form when two solutions react and their respective ions combine to produce an insoluble compound. This process is significant in the realm of chemistry, particularly in aqueous reactions, where solubility rules help predict whether a precipitate will form.

Solubility rules are guidelines developed to predict the solubility of different compounds in water. A common rule is that most sodium, potassium, and ammonium salts are soluble, while carbonate, phosphate, sulfate, and others vary depending on the cation they are paired with.

For instance, when chloride ions are mixed with silver, lead, or mercury ions in solution, the resulting compounds (silver chloride, lead(II) chloride, and mercury(II) chloride, respectively) are insoluble and thus form precipitates. Understanding these rules allows chemists to design reactions that separate components or form new substances by manipulating the conditions to favor precipitate formation.

Improving Comprehension:

To further grasp the concept of precipitate formation, visualize the process as combining the perfect recipe ingredients. When the right ions are mixed in the correct conditions, a solid (the precipitate) emerges, much like adding flour to wet ingredients can turn a liquid batter into a solid cake when baked.

Net Ionic Equations

Net ionic equations are a streamlined representation of a chemical reaction in an aqueous solution. They show only the species that are involved in the actual chemical change, omitting the spectator ions that do not participate in the formation of the precipitate.

In writing a net ionic equation, identify the ions that combine to form the insoluble precipitate and balance the equation to reflect the conservation of mass and charge. For example, when calcium ions (\(Ca^{2+}\)) react with carbonate ions (\(CO_3^{2-}\)), the net ionic equation is \(Ca^{2+} + CO_3^{2-} \rightarrow CaCO_3(s)\), which leads to the formation of solid calcium carbonate, a precipitate.

Why It Matters:

Net ionic equations are essential for understanding the essence of chemical reactions without the clutter of unnecessary information. They allow students and scientists alike to focus on the cause-and-effect relationship between reactants and products, offering a clearer picture of the transformation occurring within a solution.

Aqueous Solution Reactions

Aqueous solution reactions are processes where the reactants are dissolved in water to form a solution. The solvents and solutes in these reactions interact through the ions they release in water, leading to the formation of new compounds or precipitates.

The solubility of the ions in water is a key factor in predicting the outcomes of such reactions. Soluble ionic compounds disassociate into ions, which can then react with other ions in the solution. If the resulting compound has low solubility, it will typically settle out of the solution as a precipitate.

Understanding the Basics:

To fully appreciate the dynamics of aqueous solution reactions, picture the solution as a dance floor where ions mingle. If the right partners find each other (reacting ions leading to an insoluble compound), they exit the dance floor (the solution) by forming a solid (precipitate). The nature of the solute and solvent, concentration, and temperature are all 'music' that influences the 'dance' of ions in solution. Keeping track of the emerging patterns (reaction products) helps chemists control and utilize these reactions for various applications, from industrial synthesis to environmental remediation.