Chapter 4

Explain why some electrolyte solutions are strongly conducting, whereas others are weakly conducting.

Define the terms strong electrolyte and weak electrolyte. Give an example of each.

Explain the terms soluble and insoluble. Use the solubility rules to write the formula of an insoluble ionic compound.

What are the advantages and disadvantages of using a molecular equation to represent an ionic reaction?

What is a net ionic equation? What is the value in using a net ionic equation? Give an example.

What are the major types of chemical reactions? Give a brief description and an example of each.

Describe in words how you would prepare pure crystalline \(\mathrm{AgCl}\) and \(\mathrm{NaNO}_{3}\) from solid \(\mathrm{AgNO}_{3}\) and solid \(\mathrm{NaCl}\).

Give an example of a neutralization reaction. Label the acid, base, and salt.

Give an example of a polyprotic acid and write equations for the successive neutralizations of the acidic hydrogen atoms of the acid molecule to produce a series of salts.

Why must oxidation and reduction occur together in a reaction?

Give an example of a displacement reaction. What is the oxidizing agent? What is the reducing agent?

Describe how the amount of sodium hydroxide in a mixture can be determined by titration with hydrochloric acid of known molarity.

What is the net ionic equation for the following molecular equation? $$ \mathrm{HF}(a q)+\mathrm{KOH}(a q) \longrightarrow \mathrm{KF}(a q)+\mathrm{H}_{2} \mathrm{O}(I) $$ Hydrofluoric acid, HF, is a molecular substance and a weak electrolyte. a. \(\mathrm{H}^{+}(a q)+\mathrm{OH}^{-}(a q) \longrightarrow \mathrm{H}_{2} \mathrm{O}(l)\) b. \(\mathrm{H}^{+}(a q)+\mathrm{KOH}(a q) \longrightarrow \mathrm{K}^{+}(a q)+\mathrm{H}_{2} \mathrm{O}(l)\) c. \(\mathrm{HF}(a q)+\mathrm{KOH}(a q) \longrightarrow \mathrm{K}^{+}(a q)+\mathrm{F}^{2}(a q)\) d. \(\mathrm{HF}(a q)+\mathrm{K}^{+}(a q)+\mathrm{OH}^{-}(a q) \longrightarrow \mathrm{KF}(a q)+\mathrm{H}_{2} \mathrm{O}(I)\) e. \(\mathrm{HF}(a q)+\mathrm{OH}^{-}(a q) \longrightarrow \mathrm{F}^{-}(a q)+\mathrm{H}_{2} \mathrm{O}(l)\)

An aqueous sodium hydroxide solution mixed with an aqueous magnesium nitrate solution yields which of the following products? a. magnesium hydroxide \((a q)\) b. magnesium dihydroxide(s) c. magnesium hydroxide \((s)\) d. dimagnesium hydroxide \((s)\) e. sodium nitrate \((l)\)

Which of the following compounds would produce the highest concentration of \(\mathrm{Cl}^{-}\) ions when \(0.10 \mathrm{~mol}\) of each is placed in separate beakers containing equal volumes of water? a. \(\mathrm{NaCl}\) b. \(\mathrm{PbCl}_{2}\) c. \(\mathrm{HClO}_{4}\) d. \(\mathrm{MgCl}_{2}\) e. \(\mathrm{HCl}\)

In an aqueous \(0.10 \mathrm{M} \mathrm{HNO}_{2}\) solution \(\left(\mathrm{HNO}_{2}\right.\) is a weak electrolyte), which of the following would you expect to see in the highest concentration? a. \(\mathrm{H}_{3} \mathrm{O}^{+}\) b. \(\mathrm{NO}_{2}^{-}\) c. \(\mathrm{H}^{+}\) d. \(\mathrm{HNO}_{2}\) e. \(\mathrm{OH}^{-}\)

a. Ammonia, \(\mathrm{NH}_{3}\), is a weak electrolyte. It forms ions in solution by reacting with water molecules to form the ammonium ion and hydroxide ion. Write the balanced chemical reaction for this process, including state symbols. b. From everyday experience you are probably aware that table sugar (sucrose), \(\mathrm{C}_{12} \mathrm{H}_{22} \mathrm{O}_{11}\), is soluble in water. When sucrose dissolves in water, it doesn't form ions through any reaction with water. It just dissolves without forming ions, so it is a nonelectrolyte. Write the chemical equation for the dissolving of sucrose in water. c. Both \(\mathrm{NH}_{3}\) and \(\mathrm{C}_{12} \mathrm{H}_{22} \mathrm{O}_{11}\) are soluble molecular compounds, yet they behave differently in aqueous solution. Briefly explain why one is a weak electrolyte and the other is a nonelectrolyte. d. Hydrochloric acid, \(\mathrm{HCl}\), is a molecular compound that is a strong electrolyte. Write the chemical reaction of \(\mathrm{HCl}\) with water. e. Compare the ammonia reaction with that of hydrochloric acid. Why are both of these substances considered electrolytes? f. Explain why \(\mathrm{HCl}\) is a strong electrolyte and ammonia is a weak electrolyte. g. Classify each of the following substances as either ionic or molecular. \(\mathrm{KCl} \mathrm{NH}_{3} \mathrm{CO}_{2} \mathrm{MgBr}_{2} \mathrm{HCl} \mathrm{Ca}(\mathrm{OH})_{2} \mathrm{PbS} \mathrm{HC}_{2} \mathrm{H}_{3} \mathrm{O}_{2}\) h. For those compounds above that you classified as ionic, use the solubility rules to determine which are soluble. i. The majority of ionic substances are solids at room temperature. Describe what you would observe if you placed a soluble ionic compound and an insoluble ionic compound in separate beakers of water. j. Write the chemical equation(s), including state symbols, for what happens when each soluble ionic compound that you identified above is placed in water. Are these substances reacting with water when they are added to water? k. How would you classify the soluble ionic compounds: strong electrolyte, weak electrolyte, or nonelectrolyte? Explain your answer. L. Sodium chloride, \(\mathrm{NaCl}\), is a strong electrolyte, as is hydroiodic acid, HI. Write the chemical equations for what happens when these substances are added to water. m. Are \(\mathrm{NaCl}\) and HI strong electrolytes because they have similar behavior in aqueous solution? If not, describe, using words and equations, the different chemical process that takes place in each case.

You need to perform gravimetric analysis of a water sample in order to determine the amount of \(\mathrm{Ag}^{+}\) present. a. List three aqueous solutions that would be suitable for mixing with the sample to perform the analysis. b. Would adding \(\mathrm{KNO}_{3}(a q)\) allow you to perform the analysis? c. Assume you have performed the analysis and the silver solid that formed is moderately soluble. How might this affect your analysis results?

In this problem you need to draw two pictures of solutions in beakers at different points in time. Time zero \((t=0)\) will be the hypothetical instant at which the reactants dissolve in the solution (if they dissolve) before they react. Time after mixing \((t>0)\) will be the time required to allow sufficient interaction of the materials. For now, we assume that insoluble solids have no ions in solution and do not worry about representing the stoichiometric amounts of the dissolved ions. Here is an example: Solid \(\mathrm{NaCl}\) and solid \(\mathrm{AgNO}_{3}\) are added to a beaker containing \(250 \mathrm{~mL}\) of water.Note that we are not showing the \(\mathrm{H}_{2} \mathrm{O}\), and we are representing only the ions and solids in solution. Using the same conditions as the example (adding the solids to \(\mathrm{H}_{2} \mathrm{O}\) ), draw pictures of the following: a. Solid lead(II) nitrate and solid ammonium chloride at \(t=0\) and \(t>0\) b. \(\mathrm{FeS}(s)\) and \(\mathrm{NaNO}_{3}(s)\) at \(t=0\) and \(t>0\) c. Solid lithium iodide and solid sodium carbonate at \(t=0\) and \(t>0\)

You come across a beaker that contains water, aqueous ammonium acetate, and a precipitate of calcium phosphate. a. Write the balanced molecular equation for a reaction between two solutions containing ions that could produce this solution. b. Write the complete ionic equation for the reaction in part a. c. Write the net ionic equation for the reaction in part a.

Three acid samples are prepared for titration by \(0.01 M\) \(\mathrm{NaOH}:\) 1\. Sample 1 is prepared by dissolving \(0.01\) mol of \(\mathrm{HCl}\) in \(50 \mathrm{~mL}\) of water. 2\. Sample 2 is prepared by dissolving \(0.01\) mol of \(\mathrm{HCl}\) in \(60 \mathrm{~mL}\) of water. 3\. Sample 3 is prepared by dissolving \(0.01\) mol of \(\mathrm{HCl}\) in \(70 \mathrm{~mL}\) of water. a. Without performing a formal calculation, compare the concentrations of the three acid samples (rank them from highest to lowest). b. When the titration is performed, which sample, if any, will require the largest volume of the \(0.01 M \mathrm{NaOH}\) for neutralization?

Would you expect a precipitation reaction between an ionic compound that is an electrolyte and an ionic compound that is a nonelectrolyte? Justify your answer.

Try and answer the following questions without using a calculator. a. A solution is made by mixing \(1.0 \mathrm{~L}\) of \(0.5 \mathrm{M} \mathrm{NaCl}\) and \(0.5 \mathrm{~L}\) of \(1.0 \mathrm{M} \mathrm{CaCl}_{2}\). Which ion is at the highest concentration in the solution? b. Another solution is made by mixing \(0.50 \mathrm{~L}\) of \(1.0 \mathrm{M} \mathrm{KBr}\) and \(0.50 \mathrm{~L}\) of \(1.0 \mathrm{M} \mathrm{K}_{3} \mathrm{PO}_{4}\). What is the concentration of each ion in the solution?

If one mole of the following compounds were each placed into separate beakers containing the same amount of water, rank the \(\mathrm{Cl}^{-}(a q)\) concentrations from highest to lowest (some may be equivalent): \(\mathrm{KCl}, \mathrm{AlCl}_{3}, \mathrm{PbCl}_{2}, \mathrm{NaCl}, \mathrm{HCl}, \mathrm{NH}_{3}, \mathrm{KOH}\), and \(\mathrm{HCN}\).

Using solubility rules, predict the solubility in water of the following ionic compounds. a. \(\mathrm{PbS}\) b. \(\mathrm{AgNO}_{3}\) c. \(\mathrm{Na}_{2} \mathrm{CO}_{3}\) d. \(\mathrm{CaI}_{2}\)

Using solubility rules, predict the solubility in water of the following ionic compounds. a. \(\mathrm{Al}(\mathrm{OH})_{3}\) b. \(\mathrm{Li}_{3} \mathrm{P}\) C. \(\mathrm{NH}_{4} \mathrm{Cl}\) d. \(\mathrm{NaOH}\)

Using solubility rules, decide whether the following ionic solids are soluble or insoluble in water. If they are soluble, indicate what ions you would expect to be present in solution. a. \(\mathrm{AgBr}\) b. \(\mathrm{Li}_{2} \mathrm{SO}_{4}\) c. \(\mathrm{Ca}_{3}\left(\mathrm{PO}_{4}\right)_{2}\) d. \(\mathrm{Na}_{2} \mathrm{CO}_{3}\)

Using solubility rules, decide whether the following ionic solids are soluble or insoluble in water. If they are soluble, indicate what ions you would expect to be present in solution. a. \(\left(\mathrm{NH}_{4}\right)_{2} \mathrm{SO}_{4}\) b. \(\mathrm{BaCO}_{3}\) c. \(\mathrm{PbSO}_{4}\) d. \(\mathrm{Ca}\left(\mathrm{NO}_{3}\right)_{2}\)

Write net ionic equations for the following molecular equations. \(\mathrm{HBr}\) is a strong electrolyte. a. \(\mathrm{HBr}(a q)+\mathrm{KOH}(a q) \longrightarrow \mathrm{KBr}(a q)+\mathrm{H}_{2} \mathrm{O}(l)\) b. \(\mathrm{AgNO}_{3}(a q)+\mathrm{NaBr}(a q) \longrightarrow \mathrm{AgBr}(s)+\operatorname{NaNO}_{3}(a q)\) c. \(\mathrm{K}_{2} \mathrm{~S}(a q)+2 \mathrm{HBr}(a q) \longrightarrow 2 \mathrm{KBr}(a q)+\mathrm{H}_{2} \mathrm{~S}(g)\) d. \(\mathrm{NaOH}(a q)+\mathrm{NH}_{4} \mathrm{Br}(a q)\) \( \mathrm{NaBr}(a q)+\mathrm{NH}_{3}(g)+\mathrm{H}_{2} \mathrm{O}(l) \)

Write net ionic equations for the following molecular equations. \(\mathrm{HBr}\) is a strong electrolyte. a. \(\mathrm{HBr}(a q)+\mathrm{NH}_{3}(a q) \longrightarrow \mathrm{NH}_{4} \mathrm{Br}(a q)\) b. \(2 \mathrm{HBr}(a q)+\mathrm{Ba}(\mathrm{OH})_{2}(a q) \longrightarrow 2 \mathrm{H}_{2} \mathrm{O}(l)+\operatorname{BaBr}_{2}(a q)\) c. \(\mathrm{Pb}\left(\mathrm{NO}_{3}\right)_{2}(a q)+2 \mathrm{NaBr}(a q) \longrightarrow \mathrm{PbBr}_{2}(s)+2 \mathrm{NaNO}_{3}(a q)\) d. \(\mathrm{MgCO}_{3}(s)+\mathrm{H}_{2} \mathrm{SO}_{4}(a q)\) $$ \mathrm{MgSO}_{4}(a q)+\mathrm{H}_{2} \mathrm{O}(l)+\mathrm{CO}_{2}(g) $$

Lead(II) nitrate solution and sodium sulfate solution are mixed. Crystals of lead(II) sulfate come out of solution, leaving a solution of sodium nitrate. Write the molecular equation and the net ionic equation for the reaction.

Potassium carbonate solution reacts with aqueous hydrobromic acid to give a solution of potassium bromide, carbon dioxide gas, and water. Write the molecular equation and the net ionic equation for the reaction.

Write the molecular equation and the net ionic equation for each of the following aqueous reactions. If no reaction occurs, write \(N R\) after the arrow. a. \(\mathrm{FeSO}_{4}+\mathrm{NaCl} \longrightarrow\) b. \(\mathrm{Na}_{2} \mathrm{CO}_{3}+\mathrm{MgBr}_{2} \longrightarrow\) c. \(\mathrm{MgSO}_{4}+\mathrm{NaOH} \longrightarrow\) d. \(\mathrm{NiCl}_{2}+\mathrm{NaBr} \longrightarrow\)

Write the molecular equation and the net ionic equation for each of the following aqueous reactions. If no reaction occurs, write \(N R\) after the arrow. a. \(\mathrm{AgNO}_{3}+\mathrm{NaI} \longrightarrow\) b. \(\mathrm{Ba}\left(\mathrm{NO}_{3}\right)_{2}+\mathrm{K}_{2} \mathrm{SO}_{4} \longrightarrow\) c. \(\mathrm{Mg}\left(\mathrm{NO}_{3}\right)_{2}+\mathrm{K}_{2} \mathrm{SO}_{4} \longrightarrow\) d. \(\mathrm{CaCl}_{2}+\mathrm{Al}\left(\mathrm{NO}_{3}\right)_{3} \longrightarrow\)

For each of the following, write molecular and net ionic equations for any precipitation reaction that occurs. If no reaction occurs, indicate this. a. Solutions of barium nitrate and lithium sulfate are mixed. b. Solutions of sodium bromide and calcium nitrate are mixed. c. Solutions of aluminum sulfate and sodium hydroxide are mixed. d. Solutions of calcium bromide and sodium phosphate are mixed.

For each of the following, write molecular and net ionic equations for any precipitation reaction that occurs. If no reaction occurs, indicate this. a. Zinc chloride and sodium sulfide are dissolved in water. b. Sodium sulfide and calcium chloride are dissolved in water. C. Magnesium sulfate and potassium bromide are dissolved in water. d. Magnesium sulfate and potassium carbonate are dissolved in water.

Classify each of the following as a strong or weak acid or base. a. HF b. \(\mathrm{KOH}\) c. \(\mathrm{HClO}_{4}\) d. HIO

Classify each of the following as a strong or weak acid or base. a. \(\mathrm{NH}_{3}\) b. HCNO c. \(\mathrm{Sr}(\mathrm{OH})_{2}\) d. HI

Complete and balance each of the following molecular equations (in aqueous solution); include phase labels. Then, for each, write the net ionic equation. a. \(\mathrm{NaOH}+\mathrm{HNO}_{3} \longrightarrow\) b. \(\mathrm{HCl}+\mathrm{Ba}(\mathrm{OH})_{2} \longrightarrow\) c. \(\mathrm{HC}_{2} \mathrm{H}_{3} \mathrm{O}_{2}+\mathrm{Ca}(\mathrm{OH})_{2} \longrightarrow\) d. \(\mathrm{NH}_{3}+\mathrm{HNO}_{3} \longrightarrow\)

Complete and balance each of the following molecular equations (in aqueous solution); include phase labels. Then, for each, write the net ionic equation. a. \(\mathrm{Al}(\mathrm{OH})_{3}+\mathrm{HCl} \longrightarrow\) b. \(\mathrm{HBr}+\mathrm{Sr}(\mathrm{OH})_{2} \longrightarrow\) c. \(\mathrm{Ba}(\mathrm{OH})_{2}+\mathrm{HC}_{2} \mathrm{H}_{3} \mathrm{O}_{2} \longrightarrow\) d. \(\mathrm{HNO}_{3}+\mathrm{KOH} \longrightarrow\)

For each of the following, write the molecular equation, including phase labels. Then write the net ionic equation. Note that the salts formed in these reactions are soluble. a. the neutralization of hydrobromic acid with calcium hydroxide solution b. the reaction of solid aluminum hydroxide with nitric acid c. the reaction of aqueous hydrogen cyanide with calcium hydroxide solution d. the neutralization of lithium hydroxide solution by aqueous hydrogen cyanide

For each of the following, write the molecular equation, including phase labels. Then write the net ionic equation. Note that the salts formed in these reactions are soluble. a. the neutralization of lithium hydroxide solution by aqueous perchloric acid b. the reaction of barium hydroxide solution and aqueous nitrous acid c. the reaction of sodium hydroxide solution and aqueous nitrous acid d. the neutralization of aqueous hydrogen cyanide by aqueous strontium hydroxide

Complete the right side of each of the following molecular equations. Then write the net ionic equations. Assume all salts formed are soluble. Acid salts are possible. a. \(2 \mathrm{KOH}(a q)+\mathrm{H}_{3} \mathrm{PO}_{4}(a q) \longrightarrow\) b. \(3 \mathrm{H}_{2} \mathrm{SO}_{4}(a q)+2 \mathrm{Al}(\mathrm{OH})_{3}(s) \longrightarrow\) c. \(2 \mathrm{HC}_{2} \mathrm{H}_{3} \mathrm{O}_{2}(a q)+\mathrm{Ca}(\mathrm{OH})_{2}(a q) \longrightarrow\) d. \(\mathrm{H}_{2} \mathrm{SO}_{3}(a q)+\mathrm{NaOH}(a q) \longrightarrow\)

Complete the right side of each of the following molecular equations. Then write the net ionic equations. Assume all salts formed are soluble. Acid salts are possible. a. \(\mathrm{Ca}(\mathrm{OH})_{2}(a q)+2 \mathrm{H}_{2} \mathrm{SO}_{4}(a q) \longrightarrow\) b. \(2 \mathrm{H}_{3} \mathrm{PO}_{4}(a q)+\mathrm{Ca}(\mathrm{OH})_{2}(a q) \longrightarrow\) c. \(\mathrm{NaOH}(a q)+\mathrm{H}_{2} \mathrm{SO}_{4}(a q) \longrightarrow\) d. \(\mathrm{Sr}(\mathrm{OH})_{2}(a q)+2 \mathrm{H}_{2} \mathrm{CO}_{3}(a q) \longrightarrow\)

Write molecular and net ionic equations for the successive neutralizations of each acidic hydrogen of sulfurous acid by aqueous calcium hydroxide. \(\mathrm{CaSO}_{3}\) is insoluble; the acid salt is soluble.

Write molecular and net ionic equations for the successive neutralizations of each acidic hydrogen of phosphoric acid by calcium hydroxide solution. \(\mathrm{Ca}_{3}\left(\mathrm{PO}_{4}\right)_{2}\) is insoluble; assume that the acid salts are soluble.

The following reactions occur in aqueous solution. Complete and balance the molecular equations using phase labels. Then write the net ionic equations. a. \(\mathrm{CaS}+\mathrm{HBr} \longrightarrow\) b. \(\mathrm{MgCO}_{3}+\mathrm{HNO}_{3} \longrightarrow\) C. \(\mathrm{K}_{2} \mathrm{SO}_{3}+\mathrm{H}_{2} \mathrm{SO}_{4} \longrightarrow\)

The following reactions occur in aqueous solution. Complete and balance the molecular equations using phase labels. Then write the net ionic equations. a. \(\mathrm{BaCO}_{3}+\mathrm{HNO}_{3} \longrightarrow\) b. \(\mathrm{K}_{2} \mathrm{~S}+\mathrm{HCl} \longrightarrow\) c. \(\mathrm{CaSO}_{3}(s)+\mathrm{HI} \longrightarrow\)

Write the molecular equation and the net ionic equation for the reaction of solid iron(II) sulfide and hydrochloric acid. Add phase labels.

Write the molecular equation and the net ionic equation for the reaction of solid barium carbonate and hydrogen bromide in aqueous solution. Add phase labels.