?Based on oxidation number consideration, explain why carbon monoxide (CO) is flammable

Balance the following equations and write the corresponding ionic and net ionic equations (if appropriate). (a) \(\mathrm{CH}_{3} \mathrm{COOH}(a q)+\mathrm{KOH}(a q) \rightarrow\) (b) \(\mathrm{H}_{2} \mathrm{CO}_{3}(a q)+\mathrm{NaOH}(a q) \rightarrow\) (c) \(\mathrm{HNO}_{3}(a q)+\mathrm{Ba}(\mathrm{OH})_{2}(a q) \rightarrow\)

Give an example of a combination redox reaction, a decomposition redox reaction, and a displacement redox reaction.

All combustion reactions are redox reactions. True or false? Explain.

What is an oxidation number? How is it used to identify redox reactions? Explain why, except for ionic compounds, oxidation number does not have any physical significance.

(a) Without referring to Figure 4.11. give the oxidation numbers of the alkali and alkaline earth metals in their compounds. (b) Give the highest oxidation numbers that the Groups 3A-7A elements can have.

How is the activity series organized? How is it used in the study of redox reactions?

Use the following reaction to define redox reaction, half-reaction, oxidizing agent, reducing agent: \(4 N a(s)+O_{2}(g) \rightarrow 2 N a_{2} O(s)\)

Is it possible to have a reaction in which oxidation occurs and reduction does not? Explain.

What is the requirement for an element to undergo disproportionation reactions? Name five common elements that are likely to take part in such reactions.

For the complete redox reactions given here, (i) break down each reaction into its half-reactions; (ii) identify the oxidizing agent; (iii) identify the reducing agent. (a) \(2 \mathrm{Sr}+\mathrm{O}_{2} \rightarrow 2 \mathrm{SrO}\) (b) \(2 \mathrm{Li}+\mathrm{H}_{2} \rightarrow 2 \mathrm{LiH}\) (c) \(2 \mathrm{Cs}+\mathrm{Br}_{2} \rightarrow 2 \mathrm{CsBr}\) (d) \(3 \mathrm{Mg}+\mathrm{N}_{2} \rightarrow \mathrm{Mg}_{3} \mathrm{~N}_{2}\)

For the complete redox reactions given here, write the half-reactions and identify the oxidizing and reducing agents. (a) \(4 \mathrm{Fe}+3 \mathrm{O}_{2} \rightarrow 2 \mathrm{Fe}_{2} \mathrm{O}_{3}\) (b) \(\mathrm{Cl}_{2}+2 \mathrm{NaBr} \rightarrow 2 \mathrm{NaCl}+\mathrm{Br}_{2}\) (c) \(\mathrm{Si}+2 \mathrm{~F}_{2} \rightarrow \mathrm{SiF}_{4}\) (d) \(\mathrm{H}_{2}+\mathrm{Cl}_{2} \rightarrow 2 \mathrm{HCl}\)

Arrange the following species in order of increasing oxidation number of the sulfur atom: (a) \(\mathrm{H}_{2} \mathrm{~S}\), (b) \(\mathrm{~S}_{8}\), (c) \(\mathrm{H}_{2} \mathrm{SO}_{4}\), (d) \(\mathrm{~S}^{2-}\), (e) \(\mathrm{HS}\), (f) \(\mathrm{SO}_{2}\), (g) \(\mathrm{SO}_{3}\).

Phosphorus forms many oxoacids. Indicate the oxidation number of phosphorus in each of the following acids: (a) HPO3, (b) H3PO2, (c) H3PO3, (d) H,PO4, (e) H2P2O7, (f) H5P3O10. (a) \(\mathrm{HPO}_{3}\), (b) \(\mathrm{H}_{3} \mathrm{PO}_{2}\), (c) \(\mathrm{H}_{3} \mathrm{PO}_{3}\), (d) \(\mathrm{H}_{3} \mathrm{PO}_{4}\), (e) \(\mathrm{H}_{4} \mathrm{P}_{2} \mathrm{O}_{7}\) (f) \(\mathrm{H}_{5} \mathrm{P}_{3} \mathrm{O}_{10}\)

Give the oxidation number of the underlined atoms in the following molecules and ions: (a) \(\underline{C l }\mathrm {F}\) (b) \(\underline{I} F_{7}\) (c) \(\underline{C} H_{4}\) (d) \(\underline{C}_{2} H_{2}\) (e) \(\underline{C}_{2} H_{4}\) (f) \(\mathrm{K}_{2} \underline{\mathrm{Cr}} \mathrm{O}_{4}\) (g) \(\mathrm{K}_{2} \underline {Cr}_{2} \mathrm{O}_{7}\) (h) \(\mathrm{K \underline {Mn} O}_{4}\) (i) \(\mathrm{NaH \underline {C} O}_{3}\) (j) \(\underline{L i}_{2}\) (k) \(\mathrm{Na \underline I O}_{3}\) (l) \(\mathrm{K \underline {O } _ { 2 }}\) (m) \(\underline{P} F_{6}^{-}\) (n) \(\mathrm{K \underline {Au}} \mathrm{Cl}_{4}\)

Give the oxidation number for the following species: \(H_{2}\), \(S e_{8}\), \(P_{4}\), O, U, \(A s_{4}\), \(B_{12}\).

Give oxidation numbers for the underlined atoms in the following molecules and ions: (a) \(\underline{C s}_{2} \mathrm{O}\) (b) \(\mathrm{Ca \underline I}_{2}\) (c) \(\underline{A l}_{2} \mathrm{O}_{3}\) (d) \(\mathrm{H}_{3} \underline{A s}_{3}\) (e) \(\underline{\mathrm{Ti}} \mathrm{O}_{2}\) (f) \(\underline {Mo} {\mathrm {O}_{4}^{2-}}\) (g) \(\underline{P t} \mathrm{Cl}_{4}^{2-}\) (h) \(\underline{P t} \mathrm{Cl}_{6}^{2-}\) (i) \(\underline{\operatorname{Sn}} F_{2}\) (j) \(\underline{C l} F_{3}\) (k) \(\underline{S b} F_{6}^{-}\)

Give the oxidation numbers of the underlined atoms in the following molecules and ions: (a) \(M g_{3} \underline{N}_{2}\) (b) \(\mathrm{Cs} \underline{O}_{2}\) (c) \(C a \underline{C}_{2}\) (d) \(\underline {C} \mathrm{O}_{3}^{2-}\) (e) \(\underline{C}_{2} \mathrm{O}_{4}^{2-}\) (f) \(\underline{\mathrm{Zn}} \mathrm{O}_{2}^{2-}\) (g) \(N a \underline{B} H_{4}\) (h) \(\underline{W} O_{4}^{2-}\)

Nitric acid is a strong oxidizing agent. State which of the following species is least likely to be produced when nitric acid reacts with a strong reducing agent such as zinc metal, and explain why: \(\mathrm{N}_2\mathrm{O},\mathrm{\ NO},\mathrm{\ NO}_2,\mathrm{\ N}_2\mathrm{O}_4,\mathrm{\ N}_2\mathrm{O}_5,\mathrm{\ NH}_4^+\).

Which of the following metals can react with water? (a) \(\mathrm{Au}\), (b) \(\mathrm{Li}\), (c) \(\mathrm{Hg}\), (d) \(\mathrm{Ca}\), (e) Pt.

On the basis of oxidation number considerations, one of the following oxides would not react with molecular oxygen: NO, \(\mathrm{N}_{2} \mathrm{O}\), \(\mathrm{SO}_{2}\), \(\mathrm{SO}_{3}\), \(\mathrm{P}_{4} \mathrm{O}_{6}\). Which one is it? Why?

Predict the outcome of the reactions represented by the following equations by using the activity series, and balance the equations. (a) Cu (s) + HCl (aq) ? (b) \(I_{2}(s)+N a B r(a q) \rightarrow\) (c) \(M g(s)+\mathrm{CuSO}_{4}(a q) \rightarrow\) (d) \(C l_{2}(g)+K B r(a q) \rightarrow\)

Classify the following redox reactions: (a) \(2 \mathrm{H}_{2} \mathrm{O}_{2} \rightarrow 2 \mathrm{H}_{2} \mathrm{O}+\mathrm{O}_{2}\) (b) \(\mathrm{Mg}+2 \mathrm{AgNO}_{3} \rightarrow \mathrm{Mg}\left(\mathrm{NO}_{3}\right)_{2}+2 \mathrm{Ag}\) (c) \(\mathrm{NH}_{4} \mathrm{NO}_{2} \rightarrow \mathrm{N}_{2}+2 \mathrm{H}_{2} \mathrm{O}\) (d) \(\mathrm{H}_{2}+\mathrm{Br}_{2} \rightarrow 2 \mathrm{HBr}\)

Classify the following redox reactions: (a) \(P_{4}+10 \mathrm{Cl}_{2} \rightarrow 4 \mathrm{PCl}_{5}\) (b) \(2 \mathrm{NO} \rightarrow \mathrm{N}_{2}+\mathrm{O}_{2}\) (c) \(\mathrm{Cl}_{2}+2 \mathrm{KI} \rightarrow 2 \mathrm{KCl}+\mathrm{I}_{2}\) (d) \(3 \mathrm{HNO}_{2} \rightarrow \mathrm{HNO}_{3}+\mathrm{H}_{2} \mathrm{O}+2 \mathrm{NO}\)

Which of the following are redox processes? (a) \(\mathrm{CO}_{2} \rightarrow \mathrm{CO}_{3}^{2-}\) (b) \(\mathrm{VO}_{3} \rightarrow \mathrm{VO}_{2}\) (c) \(\mathrm{SO}_{3} \rightarrow \mathrm{SO}_{4}^{2-}\) (d) \(\mathrm{NO}_{2}^{-} \rightarrow \mathrm{NO}_{3}^{-}\) (e) \(\mathrm{Cr}^{3+} \rightarrow \mathrm{CrO}_{4}^{2-}\)

Of the following, which is most likely to be the strongest oxidizing agent? \(O_{2}\), \(O_{2}^{+}\), \(O_{2}^{-}\), \(O_{2}^{2-}\).

Write the equation for calculating molarity. Why is molarity a convenient concentration unit in chemistry?

Describe the steps involved in preparing a solution of known molar concentration using a volumetric flask.

Describe how you would prepare \(250 \mathrm{~mL}\) of a \(0.707 M \ \mathrm{NaNO}_{3}\) solution.

Calculate the mass of KI in grams required to prepare \(5.00 \times 10^{2}\) mL of a 2.80 M solution.

How many moles of \(\mathrm{MgCl}_{2}\) are present in 60.0 mL of 0.100 M \(\mathrm{MgCl}_{2}\) solution?

How many grams of KOH are present in 35.0 mL of a 5.50 M solution?

Calculate the molarity of each of the following solutions: (a) 29.0 g of ethanol \(\left(\mathrm{C}_{2} \mathrm{H}_{5} \mathrm{OH}\right)\) in 545 mL of solution, (b) 15.4 g of sucrose \(\left(\mathrm{C}_{12} \mathrm{H}_{22} \mathrm{O}_{11}\right)\) in 74.0 mL of solution, (c) 9.00 g of sodium chloride (NaCl) in 86.4 mL of solution.

Calculate the molarity of each of the following solutions: (a) 6.57 g of methanol \(\left(\mathrm{CH}_{3} \mathrm{OH}\right)\) in \(1.50 \times 10^{2}\) mL of solution, (b) 10.4 g of calcium chloride \(\left(\mathrm{CaCl}_{2}\right)\) in \(2.20 \times 10^{2}\) mL of solution, (c) 7.82 g of naphthalene \(\left(\mathrm{C}_{10} \mathrm{H}_{8}\right)\) in 85.2 mL of benzene solution.

Calculate the volume inmL of a solution required to provide the following: (a) 2.14 g of sodium chloride from a 0.270 M solution, (b) 4.30 g of ethanol from a 1.50 M solution, (c) 0.85 g of acetic acid \(\left(\mathrm{CH}_{3} \mathrm{COOH}\right)\) from a 0.30 M solution.

Determine how many grams of each of the following solutes would be needed to make \(2.50 \times 10^{2}\) mL of a 0.100 M solution: (a) cesium iodide (CsI), (b) sulfuric acid (\(\mathrm{H}_{2} \mathrm{SO}_{4}\)), (c) sodium carbonate (\(\mathrm{Na}_{2} \mathrm{CO}_{3}\)), (d) potassium dichromate (\(\mathrm{~K}_{2} \mathrm{Cr}_{2} \mathrm{O}_{7}\)), (e) potassium permanganate (\(\mathrm{KMnO}_{4}\)).

What volume of 0.416 M \(\mathrm{Mg}\left(\mathrm{NO}_{3}\right)_{2}\) should be added to 255 mL of 0.102 M \(\mathrm{KNO}_{3}\) to produce a solution with a concentration of 0.278 M \(\mathrm{NO}_{3}^{-}\) ions? Assume volumes are additive.

Barium hydroxide, often used to titrate weak organic acids, is obtained as the octahydrate. \(\mathrm{Ba}(\mathrm{OH})_{2} \cdot 8 \mathrm{H}_{2} \mathrm{O}\). What mass of \(\mathrm{Ba}(\mathrm{OH})_{2} \cdot 8 \mathrm{H}_{2} \mathrm{O}\) would be required to make 500.0 mL of a solution that is 0.1500 M in hydroxide ions?

Describe the basic steps involved in diluting a solution of known concentration.

Write the equation that enables us to calculate the concentration of a diluted solution. Give units for all the terms.

Describe how to prepare \(1.00 \ \mathrm{~L}\) of \(0.646 \ \mathrm{M} \ \mathrm{HCl}\) solution, starting with a \(2.00 \ \mathrm{M} \ \mathrm{HCl}\) solution.

Water is added to 25.0 mL of a 0.866 M \(\mathrm{KNO}_{3}\) solution until the volume of the solution is exactly 500 mL. What is the concentration of the final solution?

How would you prepare \(60.0 \ \mathrm{~mL}\) of \(0.200 \ \mathrm{M} \ \mathrm{HNO}_{3}\) from a stock solution of \(4.00 \ \mathrm{M} \ \mathrm{HNO}_{3}\)?

You have 505 mL of a 0.125 M HCl solution and you want to dilute it to exactly 0.100 M. How much water should you add? Assume volumes are additive.

A 35.2-mL, 1.66 M \(\mathrm{KMnO}_{4}\) solution is mixed with 16.7 mL of 0.892 M \(\mathrm{KMnO}_{4}\) solution. Calculate the concentration of the final solution.

A 46.2-mL, 0.568 M calcium nitrate \(\left[\mathrm{Ca}\left(\mathrm{NO}_{3}\right)_{2}\right]\) solution is mixed with 80.5 mL of 1.396 M calcium nitrate solution. Calculate the concentration of the final solution.

Describe the basic steps involved in gravimetric analysis. How does this procedure help us determine the identity of a compound or the purity of a compound if its formula is known?

Problem 1P What is an atomic mass unit? Why is it necessary to introduce such a unit?

Classify the following ionic compounds as soluble or insoluble: (a) CuS, (b) \(\mathrm{Ca}(\mathrm{OH})_{2}\), (c) \(\mathrm{Zn}\left(\mathrm{NO}_{3}\right)_{2}\).

The diagrams here show three compounds \(A B_{2}\) (a), \(A C_{2}\) (b), and \(A D_{2}\) (c) dissolved in water. Which is the strongest electrolyte and which is the weakest? (For simplicity, water molecules are not shown.)

What is the difference between a nonelectrolyte and an electrolyte? Between a weak electrolyte and a strong electrolyte?

Predict the precipitate produced by mixing an \(\mathrm{Al}\left(\mathrm{NO}_{3}\right)_{3}\) solution with a NaOH solution. Write the net ionic equation for the reaction.

Which of the diagrams here accurately describes the reaction between \(\mathrm{Ca}\left(\mathrm{NO}_{3}\right)_{2}(a q)\) and \(\mathrm{Na}_{2} \mathrm{CO}_{3}(a q)\)? For simplicity, only the \(\mathrm{Ca}^{2+}\) (yellow) and \(\mathrm{CO}_{3}^{2-}\) (blue) ions are shown.

Describe hydration. What properties of water enable its molecules to interact with ions in solution?

Classify each of the following species as a Brønsted acid or base: (a) \(\mathrm{SO}_{4}^{2-}\), (b) HI.

Which of the following diagrams best represents a weak acid? Which represents a very weak acid? Which represents a strong acid? The proton exists in water as the hydronium ion. All acids are monoprotic. (For simplicity, water molecules are not shown.)

What is the difference between the following symbols in chemical equations: \(\rightarrow\) and \(\rightleftharpoons\) ?

Write a molecular equation, an ionic equation. and a net ionic equation for the reaction between aqueous solutions of phosphoric acid and sodium hydroxide.

Which of the following combination reactions is not a redox reaction? (a) \(2 \mathrm{Mg}(s)+\mathrm{O}_{2}(s) \rightarrow 2 \mathrm{MgO}(\mathrm{s})\) (b) \(\mathrm{H}_{2}(g)+\mathrm{F}_{2}(g) \rightarrow 2 \mathrm{HF}(g)\) (c) \(\mathrm{NH}_{3}(g)+\mathrm{HCl}(g) \rightarrow \mathrm{NH}_{4} \mathrm{Cl}(\mathrm{s})\) (d) \(2 \mathrm{Na}(\mathrm{s})+\mathrm{S}(s) \rightarrow \mathrm{Na}_{2} \mathrm{~S}(\mathrm{~s})\)

Water is an extremely weak electrolyte and therefore cannot conduct electricity. Why are we often cautioned not to operate electrical appliances when our hands are wet?

Assign oxidation numbers to all the elements in the following compound and ion: (a) \(\mathrm{PF}_{3}\), (b) \(\mathrm{MnO}_{4}^{-}\).

What is the final concentration of a 0.6 M NaCl solution if its volume is doubled and the number of moles of solute is tripled?

Sodium sulfate \(\left(\mathrm{Na}_{2} \mathrm{SO}_{4}\right)\) is a strong electrolyte. What species are present in \(\mathrm{Na}_{2} \mathrm{SO}_{4}(a q)\) ?

Identify the following redox reactions by type: (a) \(\mathrm{Fe}+\mathrm{H}_{2} \mathrm{SO}_{4} \rightarrow \mathrm{FeSO}_{4}+\mathrm{H}_{2}\) (b) \(\mathrm{~S}+3 \mathrm{~F}_{2} \rightarrow \mathrm{SF}_{6}\) (c) \(2 \mathrm{CuCl} \rightarrow \mathrm{Cu}+\mathrm{CuCl}_{2}\) (d) \(2 \mathrm{Ag}+\mathrm{PtCl}_{2} \rightarrow 2 \mathrm{AgCl}+\mathrm{Pt}\)

Calculate the mass of AgBr formed if a solution containing 6.00 g of KBr is treated with an excess of \(\mathrm{AgNO}_{3}\).

The aqueous solutions of three compounds are shown in the diagram. Identify each compound as a nonelectrolyte, a weak electrolyte, and a strong electrolyte.

What is the molarity of an 85.0-mL ethanol (\(\mathrm{C}_{2} \mathrm{H}_{5} \mathrm{OH}\)) solution containing 1.77 g of ethanol?

Which of the following diagrams best represents the hydration of NaCl when dissolved in water? The \(\mathrm{Cl}^{-}\) ion is larger in size than the \(\mathrm{Na}^{+}\) ion.

What volume (in milliliters) of a 0.315 M NaOH solution contains 6.22 g of NaOH?

Identify each of the following substances as a strong electrolyte, weak electrolyte, or nonelectrolyte: (a) \(\mathrm{H}_{2} \mathrm{O}\), (b) KCl, (c) \(\mathrm{HNO}_{3}\), (d) \(\mathrm{CH}_{3} \mathrm{COOH}\), (e) \(\mathrm{C}_{12} \mathrm{H}_{22} \mathrm{O}_{11}\).

How would you prepare \(2.00\times10^2\mathrm{\ mL}\) of a \(0.866 \mathrm{M}\) NaOH solution, starting with a 5.07 \(M\) stock solution?

Identify each of the following substances as a strong electrolyte, weak electrolyte, or nonelectrolyte: (a) \(\mathrm{Ba}\left(\mathrm{NO}_{3}\right)_{2}\), (b) Ne, (c) \(\mathrm{NH}_{3}\), (d) NaOH.

A sample of \(0.3220 \mathrm{~g}\) of an ionic compound containing the bromide ion \(\left(\mathrm{Br}^{-}\right)\) is dissolved in water and treated with an excess of \(\mathrm{AgNO}_{3}\). If the mass of the \(\mathrm{AgBr}\) precipitate that forms is \(0.6964 \mathrm{~g}\), what is the percent by mass of Br in the original compound?

The passage of electricity through an electrolyte solution is caused by the movement of (a) electrons only, (b) cations only, (c) anions only, (d) both cations and anions.

How many grams of KHP are needed to neutralize 18.64 mL of a 0.1004 M NaOH solution?

Predict and explain which of the following systems are electrically conducting: (a) solid NaCl, (b) molten NaCl, (c) an aqueous solution of NaCl.

Practice Exercise How many milliliters of a \(1.28 \ \mathrm{M} \mathrm{H}_{2} \mathrm{SO}_{4}\) solution are needed to neutralize \(60.2 \mathrm{~mL}\) of a \(0.427 \ \mathrm{M} \ \mathrm{KOH}\) solution?

You are given a water-soluble compound X. Describe how you would determine whether it is an electrolyte or a nonelectrolyte. If it is an electrolyte, how would you determine whether it is strong or weak?

How many milliliters of a 0.206 M HI solution are needed to reduce 22.5 mL of a 0.374 M \(\mathrm{KMnO}_{4}\) solution according to the following equation: \(10 \mathrm{HI}+2 \mathrm{KMnO}_{4}+3 \mathrm{H}_{2} \mathrm{SO}_{4} \rightarrow 5 \mathrm{I}_{2}+2 \mathrm{MnSO}_{4}+\mathrm{K}_{2} \mathrm{SO}_{4}+8 \mathrm{H}_{2} \mathrm{O}\)

Explain why a solution of HCl in benzene does not conduct electricity but in water it does.

What is the difference between an ionic equation and a molecular equation?

Two aqueous solutions of \(\mathrm{AgNO}_{3}\) and NaCl are mixed. Which of the following diagrams best represents the mixture? For simplicity, water molecules are not shown. (Color codes are: \(\mathrm{Ag}^{+}\) = gray, \(\mathrm{Cl}^{-}\) = orange, \(\mathrm{Na}^{+}\) = green. \(\mathrm{NO}_{3}^{-}\) = blue.)

What is the advantage of writing net ionic equations?

Two aqueous solutions of KOH and MgCl2 are mixed. Which of the following diagrams best represents the mixture? For simplicity, water molecules are not shown. (Color codes are: \(K^+\) = purple, \(OH^–\) = red, \(Mg^{2+}\) = green, \(Cl^–\) = orange.)

Characterize the following compounds as soluble or insoluble in water: (a) \(\mathrm{Ca}_{3}\left(\mathrm{PO}_{4}\right)_{2}\), (b) \(\mathrm{Mn}(\mathrm{OH})_{2}\), (c) \(\mathrm{AgClO}_{3}\), (d) \(\mathrm{K}_2\mathrm{S}\).

Characterize the following compounds as soluble or insoluble in water: (a) \(\mathrm{CaCO}_{3}\), (b) \(\mathrm{ZnSO}_{4}\), (c) \(\mathrm{Hg}\left(\mathrm{NO}_{3}\right)_{2}\), (d) \(\mathrm{HgSO}_{4}\), (e) \(\mathrm{NH}_{4} \mathrm{ClO}_{4}\).

Write ionic and net ionic equations for the following reactions: (a) \(\mathrm{AgNO}_{3}(a q)+\mathrm{Na}_{2} \mathrm{SO}_{4}(a q) \rightarrow\) (b) \(\mathrm{BaCl}_{2}(a q)+\mathrm{ZnSO}_{4}(a q) \rightarrow\) (c) \(\left(\mathrm{NH}_{4}\right)_{2} \mathrm{CO}_{3}(a q)+\mathrm{CaCl}_{2}(a q) \rightarrow\)

Which of the following processes will likely result in a precipitation reaction? (a) Mixing a \(\mathrm{NaNO}_{3}\) solution with a \(\mathrm{CuSO}_{4}\) solution. (b) Mixing a \(\mathrm{BaCl}_{2}\) solution with a \(\mathrm{~K}_{2} \mathrm{SO}_{4}\) solution. Write a net ionic equation for the precipitation reaction.

Write ionic and net ionic equations for the following reactions: (a) \(\mathrm{Na}_{2} \mathrm{~S}(a q)+\mathrm{ZnCl}_{2}(a q) \rightarrow\) (b) \(\mathrm{K}_{3} \mathrm{PO}_{4}(a q)+3 \mathrm{Sr}\left(\mathrm{NO}_{3}\right)_{2}(a q) \rightarrow\) (c) \(\mathrm{Mg}\left(\mathrm{NO}_{3}\right)_{2}(a q)+2 \mathrm{NaOH}(a q) \rightarrow\)

With reference to Table 4.2, suggest one method by which you might separate (a) \(\mathrm{K}^{+}\) from \(\mathrm{Ag}^{+}\), (b) \(\mathrm{Ba}^{2+}\) from \(\mathrm{Pb}^{2+}\), (c) \(\mathrm{NH}_{4}^{+}\) from \(\mathrm{Ca}^{2+}\), (d) \(\mathrm{Ba}^{2+}\) from \(\mathrm{Cu}^{2+}\). All cations are assumed to be in aqueous solution, and the common anion is the nitrate ion.

List the general properties of acids and bases.

Give Arrhenius’s and Brønsted’s definitions of an acid and a base. Why are Brønsted’s definitions more useful in describing acid–base properties?

Describe how you would prepare the following compounds: (a) \(M g(O H)_{2}\), (b) Agl, (c) \(B a_{3}\left(P O_{4}\right)_{2}\).

Someone spilled concentrated sulfuric acid on the floor of a chemistry laboratory. To neutralize the acid, would it be preferrable to pour concentrated sodium hydroxide solution or spray solid sodium bicarbonate over the acid? Explain your choice and the chemical basis for the action.

Describe in each case how you would separate the cations or anions in an aqueous solution of: (a) \(\mathrm{NaNO}_{3}\) and \(\mathrm{Ba}\left(\mathrm{NO}_{3}\right)_{2}\), (b) \(\mathrm{Mg}\left(\mathrm{NO}_{3}\right)_{2}\) and \(\mathrm{KNO}_{3}\), (c) KBr and \(\mathrm{KNO}_{3}\), (d) \(\mathrm{K}_{3} \mathrm{PO}_{4}\) and \(\mathrm{KNO}_{3}\), (e) \(\mathrm{Na}_{2} \mathrm{CO}_{3}\) and \(\mathrm{NaNO}_{3}\).

The following are common household compounds: table salt (NaCl), table sugar (sucrose), vinegar (contains acetic acid), baking soda (\(\mathrm{NaHCO}_{3}\)), washing soda (\(\mathrm{Na}_{2} \mathrm{CO}_{3} \cdot 10 \mathrm{H}_{2} \mathrm{O}\)), boric acid (\(\mathrm{H}_{3} \mathrm{BO}_{3}\), used in eyewash), epsom salt (\(\mathrm{MgSO}_{4} \cdot 7 \mathrm{H}_{2} \mathrm{O}\)), sodium hydroxide (used in drain openers), ammonia, milk of magnesia [\(\mathrm{Mg}(\mathrm{OH})_{2}\)], and calcium carbonate. Based on what you have learned in this chapter, describe test(s) that would enable you to identify each of these compounds.

Sulfites (compounds containing the \(\mathrm{SO}_{3}^{2-}\) ions) are used as preservatives in dried fruit and vegetables and in wine making. In an experiment to test the presence of sulfite in fruit, a student first soaked several dried apricots in water overnight and then filtered the solution to remove all solid particles. She then treated the solution with hydrogen peroxide (\(\mathrm{H}_{2} \mathrm{O}_{2}\)) to oxidize the sulfite ions to sulfate ions. Finally, the sulfate ions were precipitated by treating the solution with a few drops of a barium chloride (\(\mathrm{BaCl}_{2}\)) solution. Write a balanced equation for each of the preceding steps.

A 0.9157-g mixture of \(\mathrm{CaBr}_{2}\) and NaBr is dissolved in water, and \(\mathrm{AgNO}_{3}\) is added to the solution to form AgBr precipitate. If the mass of the precipitate is 1.6930 g, what is the percent by mass of NaBr in the original mixture?

Hydrogen halides (HF, HCl, HBr, HI) are highly reactive compounds that have many industrial and laboratory uses. (a) In the laboratory, HF and HCl can be generated by reacting \(\mathrm{CaF}_{2}\) and NaCl with concentrated sulfuric acid. Write appropriate equations for the reactions. (Hint: These are not redox reactions.) (b) Why is it that HBr and HI cannot be prepared similarly, that is, by reacting NaBr and Nal with concentrated sulfuric acid? (Hint: \(\mathrm{H}_{2} \mathrm{SO}_{4}\) is a stronger oxidizing agent than both \(B r_{2}\) and \(I_{2}\).) (c) HBr can be prepared by reacting phosphorus tri- bromide (\(\mathrm{PBr}_{3}\)) with water. Write an equation for this reaction.

A 325-mL sample of solution contains 25.3 g of \(\mathrm{CaCl}_{2}\). (a) Calculate the molar concentration of \(\mathrm{Cl}^{-}\) in this solution, (b) How many grams of \(\mathrm{Cl}^{-}\) are in 0.100 L of this solution?

Phosphoric acid (\(\mathrm{H}_{3} \mathrm{PO}_{4}\)) is an important industrial chemical used in fertilizers, in detergents, and in the food industry. It is produced by two different methods. In the electric furnace method, elemental phosphorus (\(\mathrm{P}_{4}\)) is burned in air to form \(\mathrm{P}_{4} \mathrm{O}_{10}\). which is then reacted with water to give \(\mathrm{H}_{3} \mathrm{PO}_{4}\). In the wet process, the mineral phosphate rock fluorapatite [\(\mathrm{Ca}_{5}\left(\mathrm{PO}_{4}\right)_{3} \mathrm{~F}\)] is reacted with sulfuric acid to give \(\mathrm{H}_{3} \mathrm{PO}_{4}\) (and HF and \(\mathrm{CaSO}_{4}\)). Write equations for these processes and classify each step as precipitation, acid–base, or redox reaction.

Ammonium nitrate \(\left(\mathrm{NH}_4 \mathrm{NO}_3\right)\) is one of the most important nitrogen-containing fertilizers. Its purity can be analyzed by titrating a solution of \(\mathrm{NH}_4 \mathrm{NO}_3\) with a standard \(\mathrm{NaOH}\) solution. In one experiment a \(0.2041 \mathrm{~g}\) sample of industrially prepared \(\mathrm{NH}_4 \mathrm{NO}_3\) required \(24.42 \mathrm{~mL}\) of \(0.1023 \mathrm{M}\) \(\mathrm{NaOH}\) for neutralization. (a) Write a net ionic equation for the reaction. (b) What is the percent purity of the sample?

Is the following reaction a redox reaction? Explain. \(3 \mathrm{O}_{2}(\mathrm{~g}) \rightarrow 2 \mathrm{O}_{3}(\mathrm{~g})\)

What is the oxidation number of O in HFO?

Use molecular models like those in Figures 4.7 and 4.8 to represent the following acid–base reactions: (a) \(\mathrm{OH}^{-}+\mathrm{H}_{3} \mathrm{O}^{+} \rightarrow 2 \mathrm{H}_{2} \mathrm{O}\) (b) \(\mathrm{NH}_{4}^{+}+\mathrm{NH}_{2}^{-} \rightarrow 2 \mathrm{NH}_{3}\) Identify the Brønsted acid and base in each case.

The alcohol content in a 10.0-g sample of blood from a driver required 4.23 mL of 0.07654 M \(\mathrm{K}_{2} \mathrm{Cr}_{2} \mathrm{O}_{7}\) for titration. Should the police prosecute the individual for drunken driving? (Hint: See Chemistry in Action essay on p. 144.)

On standing, a concentrated nitric acid gradually turns yellow in color. Explain. (Hint: Nitric acid slowly decomposes. Nitrogen dioxide is a colored gas.)

Describe the laboratory preparation for the following gases: (a) hydrogen, (b) oxygen, (c) carbon dioxide, and (d) nitrogen. Indicate the physical states of the reactants and products in each case. [Hint:Nitrogen can be obtained by heating ammonium nitrite \(\left(\mathrm{NH}_4 \mathrm{NO}_2\right)\).]

Referring to Figure 4.18, explain why one must first dissolve the solid completely before making up the solution to the correct volume.

Can the following decomposition reaction be characterized as an acid–base reaction? Explain. \(\mathrm{NH}_{4} \mathrm{Cl}(\mathrm{s}) \rightarrow \mathrm{NH}_{3}(\mathrm{~g})+\mathrm{HCl}(\mathrm{g})\)

Give a chemical explanation for each of the following: (a) When calcium metal is added to a sulfuric acid solution, hydrogen gas is generated. After a few minutes, the reaction slows down and eventually stops even though none of the reactants is used up. Explain. (b) In the activity series, aluminum is above hydrogen, yet the metal appears to be unreactive toward steam and hydrochloric acid. Why? (c) Sodium and potassium lie above copper in the activity series. Explain why \(\mathrm{Cu}^{2+}\) ions in \(\mathrm{a} \mathrm{CuSO}_4\) solution are not converted to metallic copper upon the addition of these metals. (d) A metal M reacts slowly with steam. There is no visible change when it is placed in a pale green iron(II) sulfate solution. Where should we place M in the activity series? (e) Before aluminum metal was obtained by electrolysis, it was produced by reducing its chloride \(\left(\mathrm{AlCl}_3\right)\) with an active metal. What metals would you use to produce aluminum in that way?

The recommended procedure for preparing a very dilute solution is not to weigh out a very small mass or measure a very small volume of a stock solution. Instead, it is done by a series of dilutions. A sample of \(0.8214 \mathrm{~g}\) of \(\mathrm{KMnO}_{4}\) was dissolved in water and made up to the volume in a \(500-\mathrm{mL}\) volumetric flask. A 2.000-mL sample of this solution was transferred to a \(1000-\mathrm{mL}\) volumetric flask and diluted to the mark with water. Next, \(10.00 \mathrm{~mL}\) of the diluted solution were transferred to a $250-\mathrm{mL}\) flask and diluted to the mark with water. (a) Calculate the concentration (in molarity) of the final solution. (b) Calculate the mass of \(\mathrm{KMnO}_{4}\) needed to directly prepare the final solution.

The following “cycle of copper” experiment is performed in some general chemistry laboratories. The series of reactions starts with copper and ends with metallic copper. The steps are as follows: (1) A piece of copper wire of known mass is allowed to react with concentrated nitric acid [the products are copper(II) nitrate, nitrogen dioxide, and water]. (2) The copper(II) nitrate is treated with a sodium hydroxide solution to form copper(II) hydroxide precipitate. (3) On heating, copper(II) hydroxide decomposes to yield copper(II) oxide. (4) The copper(II) oxide is reacted with concentrated sulfuric acid to yield copper(II) sulfate. (5) Copper(II) sulfate is treated with an excess of zinc metal to form metallic copper. (6) The remaining zinc metal is removed by treatment with hydrochloric acid, and metallic copper is filtered, dried, and weighed. (a) Write a balanced equation for each step and classify the reactions. (b) Assuming that a student started with 65.6 g of copper, calculate the theoretical yield at each step. (c) Considering the nature of the steps, comment on why it is possible to recover most of the copper used at the start.

A quantity of 25.0 mL of a solution containing both \(\mathrm{Fe}^{2+}\) and \(\mathrm{Fe}^{3+}\) ions is titrated with 23.0 mL of 0.0200 M \(\mathrm{KMnO}_{4}\) (in dilute sulfuric acid). As a result, all of the \(\mathrm{Fe}^{2+}\) ions are oxidized to \(\mathrm{Fe}^{3+}\) ions. Next, the solution is treated with Zn metal to convert all of the \(\mathrm{Fe}^{3+}\) ions to \(\mathrm{Fe}^{2+}\) ions. Finally, the solution containing only the \(\mathrm{Fe}^{2+}\) ions requires 40.0 mL of the same \(\mathrm{KMnO}_{4}\) solution for oxidation to \(\mathrm{Fe}^{3+}\). Calculate the molar concentrations of \(\mathrm{Fe}^{2+}\) and \(\mathrm{Fe}^{3+}\) in the original solution. The net ionic equation is \(\mathrm{MnO}_{4}^{-}+5 \mathrm{Fe}^{2+}+8 \mathrm{H}^{+} \rightarrow \mathrm{Mn}^{2+}+5 \mathrm{Fe}^{3+} 4 \mathrm{H}_{2} \mathrm{O}\)

Use the periodic table framework shown to show the names and positions of two metals that can (a) displace hydrogen from cold water, (b) displace hydrogen from steam, and (c) displace hydrogen from acid. Also show two metals that can react neither with water nor acid.

Referring to the Chemistry in Action essay on page 156, answer the following questions: (a) Identify the precipitation, acid–base, and redox processes. (b) Instead of calcium oxide, why don’t we simply add sodium hydroxide to seawater to precipitate magnesium hydroxide? (c) Sometimes a mineral called dolomite (a mixture of \(\mathrm{CaCO}_{3}\) and \(\mathrm{MgCO}_{3}\)) is substituted for limestone to bring about the precipitation of magnesium hydroxide. What is the advantage of using dolomite?

\(\mathrm{A} 22.02-\mathrm{mL}\) solution containing \(1.615 \mathrm{~g} \mathrm{Mg}\left(\mathrm{NO}_{3}\right)_{2}\) is mixed with a \(28.64-\mathrm{mL}\) solution containing \(1.073 \mathrm{~g}\) \(\mathrm{NaOH}\). Calculate the concentrations of the ions remaining in solution after the reaction is complete. Assume volumes are additive.

Chemical tests of four metals A, B, C, and D show the following results. (a) Only B and C react with 0.5 M HCl to give \(H_{2}\) gas. (b) When B is added to a solution containing the ions of the other metals, metallic A, C, and D are formed. (c) A reacts with 6 M \(\mathrm{HNO}_{3}\) but D does not. Arrange the metals in the increasing order as reducing agents. Suggest four metals that fit these descriptions.

The antibiotic gramicidin A can transport \(\mathrm{Na}^{+}\) ions into a certain cell at the rate of \(5.0 \times 10^{7}\) \(\mathrm{Na}^{+}\) ions \(s^{-1}\). Calculate the time in seconds to transport enough \(\mathrm{Na}^{+}\) ions to increase its concentration by \(8.0 \times 10^{-3}\) M in a cell whose intracellular volume is \(2.0 \times 10^{-10}\) mL.

Shown here are two aqueous solutions containing various ions. The volume of each solution is 600 mL. (a) Write a net ionic equation for the reaction after the solutions are mixed. (b) Calculate the mass of the precipitates formed and the concentrations of the ions in the mixed solution. Treat each sphere as 0.0500 mol.

Many proteins contain metal ions for structural and/or redox functions. Which of the following metals fit into one or both categories: Ca, Cu, Fe, Mg, Mn, Ni, Zn?

The fastest way to introduce therapeutic agents into the bloodstream is by direct delivery into a vein (intravenous therapy, or IV therapy). A clinical researcher wishes to establish an initial concentration of \(6 \times 10^{-4}\) mmol/L in the bloodstream of an adult male participating in a trial study of a new drug. The drug serum is prepared in the hospital’s pharmacy at a concentration of \(1.2 \times 10^{-3}\) mol/L. How much of the serum should be introduced intravenously in order to achieve the desired initial blood concentration of the drug?

Public water supplies are often “fluoridated” by the addition of compounds such as NaF, \(\mathrm{H}_{2} \mathrm{SiF}_{6}\), and \(\mathrm{Na}_{2} \mathrm{SiF}_{6}\). It is well established that fluoride helps prevent tooth decay; however, care must be taken not to exceed safe levels of fluoride, which can stain or etch tooth enamel (dental fluorosis). A safe and effective concentration of fluoride in drinking water is generally considered to be around 1 mg/L. How much fluoride would a person consume by drinking fluoridated water in 1 year? What would be the equivalent mass as sodium fluoride?

Potassium superoxide (\(\mathrm{KO}_{2}\)), a useful source of oxygen employed in breathing equipment, reacts with water to form potassium hydroxide, hydrogen peroxide, and oxygen. Furthermore, potassium superoxide also reacts with carbon dioxide to form potassium carbonate and oxygen. (a) Write equations for these two reactions and comment on the effectiveness of potassium superoxide in this application. (b) Focusing only on the reaction between \(\mathrm{KO}_{2}\) and \(\mathrm{CO}_{2}\), estimate the amount of \(\mathrm{KO}_{2}\) needed to sustain a worker in a polluted environment for 30 min. See Problem 1.69 for useful information.

Because acid–base and precipitation reactions discussed in this chapter all involve ionic species, their progress can be monitored by measuring the electrical conductance of the solution. Match the following reactions with the diagrams shown here. The electrical conductance is shown in arbitrary units. (1) A 1.0 M KOH solution is added to 1.0 L of 1.0 M \(\mathrm{CH}_{3} \mathrm{COOH}\). (2) A 1.0 M NaOH solution is added to 1.0 L of 1.0 M HCl. (3) A 1.0 M \(\mathrm{BaCl}_{2}\) solution is added to 1.0 L of 1.0 M \(\mathrm{K}_{2} \mathrm{SO}_{4}\). (4) A 1.0 M NaCl solution is added to 1.0 L of 1.0 M \(\mathrm{AgNO}_{3}\). (5) A 1.0 M \(\mathrm{CH}_{3} \mathrm{COOH}\) solution is added to 1.0 Lof 1.0 M \(\mathrm{NH}_{3}\).

Distilled water must be used in the gravimetric analysis of chlorides. Why?

If \(30.0 \mathrm{~mL}\) of \(0.150 \mathrm{M} \mathrm{CaCl}_{2}\) is added to \(15.0 \mathrm{~mL}\) of \(0.100 \mathrm{M} \mathrm{AgNO}_{3}\), what is the mass in grams of \(\mathrm{AgCl}\) precipitate?

A sample of \(0.6760 \mathrm{~g}\) of an unknown compound containing barium ions \(\left(\mathrm{Ba}^{2+}\right)\) is dissolved in water and treated with an excess of \(\mathrm{Na}_{2} \mathrm{SO}_{4}\). If the mass of the \(\mathrm{BaSO}_{4}\) precipitate formed is \(0.4105 \mathrm{~g}\), what is the percent by mass of Ba in the original unknown compound?

How many grams of NaCl are required to precipitate most of the \(Ag^+\) ions from \(2.50\times 10^2\) mL of 0.0113 M \(AgNO_3\) solution? Write the net ionic equation for the reaction.

The concentration of sulfate in water can be determined by adding a solution of barium chloride to precipitate the sulfate ion. Write the net ionic equation for this reaction. Treating a 145-mL sample of water with excess \(\mathrm{BaCl}_{2}(a q)\) precipitated 0.330 g of \(\mathrm{BaSO}_{4}\). Determine the concentration of sulfate in the original water sample.

Describe the basic steps involved in an acid–base titration. Why is this technique of great practical value?

How does an acid–base indicator work?

A student carried out two titrations using a NaOH solution of unknown concentration in the buret. In one titration she weighed out 0.2458 g of KHP (see p. 152) and transferred it to an Erlenmeyer flask. She then added 20.00 mL of distilled water to dissolve the acid. In the other titration she weighed out 0.2507 g of KHP but added 40.00 mL of distilled water to dissolve the acid. Assuming no experimental error, would she obtain the same result for the concentration of the NaOH solution?

Would the volume of a 0.10 M NaOH solution needed to titrate 25.0 mL of a 0.10 M \(\mathrm{HNO}_{2}\) (a weak acid) solution be different from that needed to titrate 25.0 mL of a 0.10 M HCl (a strong acid) solution?

A quantity of 18.68 mL of a KOH solution is needed to neutralize 0.4218 g of KHP. What is the concentration (in molarity) of the KOH solution?

Calculate the concentration (in molarity ) of a NaOH solution if 25.0 mL of the solution are needed to neutralize 17.4 mL of a 0.312 M HCl solution.

Calculate the volume in \(\mathrm{mL}\) of a \(1.420 \ \mathrm{MaOH}\) solution required to titrate the following solutions: (a) \(25.00 \mathrm{~mL}\) of a \(2.430 \mathrm{M} \ \mathrm{HCl}\) solution (b) \(25.00 \mathrm{~mL}\) of a \(4.500 \mathrm{M} \ \mathrm{H}_{2} \mathrm{SO}_{4}\) solution (c) \(25.00 \mathrm{~mL}\) of a \(1.500 \mathrm{M} \ \mathrm{H}_{3} \mathrm{PO}_{4}\) solution

What volume of a 0.500 M HCl solution is needed to neutralize each of the following: (a) 10.0 mL of a 0.300 M NaOH solution (b) 10.0 mL of a 0.200 M \(B a(O H)_{2}\) solution

What are the similarities and differences between acid–base titrations and redox titrations?

Explain why potassium permanganate \(\left(\mathrm{KMnO}_{4}\right)\) and potassium dichromate \(\left(\mathrm{K}_{2} \mathrm{Cr}_{2} \mathrm{O}_{7}\right)\) can serve as internal indicators in redox titrations.

Iron(II) can be oxidized by an acidic \(\mathrm{K}_{2} \mathrm{Cr}_{2} \mathrm{O}_{7}\) solution according to the net ionic equation: \(\mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}+6 \mathrm{Fe}^{2+}+14 \mathrm{H}^{+} \stackrel{\longrightarrow}{2 \mathrm{Cr}^{3+}}+6 \mathrm{Fe}^{3+}+7 \mathrm{H}_{2} \mathrm{O}\) If it takes \(26.0 \mathrm{~mL}\) of \(0.0250 \mathrm{M} \mathrm{K}_{2} \mathrm{Cr}_{2} \mathrm{O}_{7}\) to titrate \(25.0 \mathrm{~mL}\) of a solution containing \(\mathrm{Fe}^{2+}\), what is the molar concentration of \(\mathrm{Fe}^{2+}\)?

The \(\mathrm{SO}_{2}\) present in air is mainly responsible for the acid rain phenomenon. Its concentration can be determined by titrating against a standard permanganate solution as follows: \(5 \mathrm{SO}_{2}+2 \mathrm{MnO}_{4}^{-}+2 \mathrm{H}_{2} \mathrm{O} \rightarrow 5 \mathrm{SO}_{4}^{2-}+2 \mathrm{Mn}^{2+}+4 \mathrm{H}^{+}\) Calculate the number of grams of \(\mathrm{SO}_{2}\) in a sample of air if 7.37 mL of 0.00800 M \(\mathrm{KMnO}_{4}\) solution are required for the titration.

A sample of iron ore (containing only \(\mathrm{Fe}^{2+}\) ions) weighing 0.2792 g was dissolved in dilute acid solution, and all the Fe(II) was converted to Fe(III) ions. The solution required 23.30 mL of 0.0194 M \(\mathrm{K}_{2} \mathrm{Cr}_{2} \mathrm{O}_{7}\) for titration. Calculate the percent by mass of iron in the ore. (Hint: See Problem 4.95 for the balanced equation.)

The concentration of a hydrogen peroxide solution can be conveniently determined by titration against a standardized potassium permanganate solution in an acidic medium according to the following equation: \(2 \mathrm{MnO}_{4}^{-}+5 \mathrm{H}_{2} \mathrm{O}_{2}+6 \mathrm{H}^{+} \longrightarrow_{5 \mathrm{O}_{2}}+2 \mathrm{Mn}^{2+}+8 \mathrm{H}_{2} \mathrm{O}\) If \(36.44 \mathrm{~mL}\) of a \(0.01652 \mathrm{M} \ \mathrm{KMnO}_{4}\) solution are required to oxidize \(25.00 \mathrm{~mL}\) of a \(\mathrm{H}_{2} \mathrm{O}_{2}\) solution, calculate the molarity of the \(\mathrm{H}_{2} \mathrm{O}_{2}\) solution.

Oxalic acid (\(\mathrm{H}_{2} \mathrm{C}_{2} \mathrm{O}_{4}\)) is present in many plants and vegetables. If 24.0 mL of 0.0100 M \(\mathrm{KMnO}_{4}\) solution is needed to titrate 1.00 g of a sample of \(\mathrm{H}_{2} \mathrm{C}_{2} \mathrm{O}_{4}\) to the equivalence point, what is the percent by mass of \(\mathrm{H}_{2} \mathrm{C}_{2} \mathrm{O}_{4}\) in the sample? The net ionic equation is \(2 \mathrm{MnO}_{4}^{-}+16 \mathrm{H}^{+}+5 \mathrm{C}_{2} \mathrm{O}_{4}^{2-} \rightarrow 2 \mathrm{Mn}^{2+}+10 \mathrm{CO}_{2}+8 \mathrm{H}_{2} \mathrm{O}\)

A 15.0-mL sample of an oxalic acid solution requires 25.2 mL of 0.149 M NaOH for neutralization. Calculate the volume of a 0.122 M \(\mathrm{KMnO}_{4}\) solution needed to react with a second 15.0-mL sample of the oxalic acid solution. (Hint: Oxalic acid is a diprotic acid. See Problem 4.99 for redox equation.)

Iodate ion, \(\mathrm{IO}_{3}^{-}\) oxidizes \(\mathrm{SO}_{3}^{2-}\) in acidic solution. The half-reaction for the oxidation is \(\mathrm{SO}_{3}^{2-}+\mathrm{H}_{2} \mathrm{O} \rightarrow \mathrm{SO}_{4}^{2-}+2 \mathrm{H}^{+}+2 e^{-}\) A 100.0-mL sample of solution containing 1.390 g of \(\mathrm{KIO}_{3}\) reacts with 32.5 mL of 0.500 M \(\mathrm{Na}_{2} \mathrm{SO}_{3}\). What is the final oxidation state of the iodine after the reaction has occurred?

Calcium oxalate \(\left(\mathrm{CaC}_{2} \mathrm{O}_{4}\right)\), the main component of kidney stones, is insoluble in water. For this reason it can be used to determine the amount of \(\mathrm{Ca}^{2+}\) ions in fluids such as blood. The calcium oxalate isolated from blood is dissolved in acid and titrated against a standardized \(\mathrm{KMnO}_{4}\) solution, as shown in Problem 4.99. In one test it is found that the calcium oxalate isolated from a 10.0-mL sample of blood requires 24.2 mL of \(9.56 \times 10^{-4}\) M \(\mathrm{KMnO}_{4}\) for titration. Calculate the number of milligrams of calcium per milliliter of blood.

Classify the following reactions according to the types discussed in the chapter: (a) \(\mathrm{Cl}_{2}+2 \mathrm{OH}^{-} \longrightarrow \mathrm{Cl}^{-}+\mathrm{ClO}^{-}+\mathrm{H}_{2} \mathrm{O}\) (b) \(\mathrm{Ca}^{2+}+\mathrm{CO}_{3}^{2-} \longrightarrow \mathrm{CaCO}_{3}\) (c) \(\mathrm{NH}_{3}+\mathrm{H}^{+} \longrightarrow \mathrm{NH}_{4}^{+}\) (d) \(2 \mathrm{CCl}_{4}+\mathrm{CrO}_{4}^{2-} \longrightarrow$ $2 \mathrm{COCl}_{2}+\mathrm{CrO}_{2} \mathrm{Cl}_{2}+2 \mathrm{Cl}^{-}\) (e) \(\mathrm{Ca}+\mathrm{F}_{2} \longrightarrow \mathrm{CaF}_{2}\) (f) \(2 \mathrm{Li}+\mathrm{H}_{2} \longrightarrow 2 \mathrm{LiH}\) (g) \(\mathrm{Ba}\left(\mathrm{NO}_{3}\right)_{2}+\mathrm{Na}_{2} \mathrm{SO}_{4} \longrightarrow 2 \mathrm{NaNO}_{3}+\mathrm{BaSO}_{4}\) (h) \(\mathrm{CuO}+\mathrm{H}_{2} \longrightarrow \mathrm{Cu}+\mathrm{H}_{2} \mathrm{O}\) (i) \(\mathrm{Zn}+2 \mathrm{HCl} \longrightarrow \mathrm{ZnCl}_{2}+\mathrm{H}_{2}\) (j) \(2 \mathrm{FeCl}_{2}+\mathrm{Cl}_{2} \longrightarrow 2 \mathrm{FeCl}_{3}\) (k) \(\mathrm{LiOH}+\mathrm{HNO}_{3} \longrightarrow \mathrm{LiNO}_{3}+\mathrm{H}_{2} \mathrm{O}\)

Oxygen \(\left(\mathrm{O}_{2}\right)\) and carbon dioxide \(\left(\mathrm{CO}_{2}\right)\) are colorless and odorless gases. Suggest two chemical tests that would enable you to distinguish between these two gases.

Which of the following aqueous solutions would you expect to be the best conductor of electricity at 25°C? Explain your answer. (a) 0.20 M NaCl (b) 0.60 M \(\mathrm{CH}_{3} \mathrm{COOH}\) (c) 0.25 M HCl (d) 0.20 M \(\mathrm{Mg}\left(\mathrm{NO}_{3}\right)_{2}\)

A \(5.00 \times 10^{2}\)-mL sample of 2.00 M HCl solution is treated with 4.47 g of magnesium. Calculate the concentration of the acid solution after all the metal has reacted. Assume that the volume remains unchanged.

Shown here are two aqueous solutions containing various ions. The volume of each solution is 200 mL. (a) Calculate the mass of the precipitate (in g) after the solutions are mixed. (b) What are the concentrations (in M) of the ions in the final solution? Treat each sphere as 0.100 mol. Assume the volumes are additive.

Shown here are two aqueous solutions containing various ions. The volume of each solution is 200 mL. (a) Calculate the mass of the precipitate (in g) after the solutions are mixed. (b) What are the concentrations (in M) of the ions in the final solution? Treat each sphere as 0.100 mol. Assume the volumes are additive.

Calculate the volume of a 0.156 M \(\mathrm{CuSO}_{4}\) solution that would react with 7.89 g of zinc.

Sodium carbonate (\(\mathrm{Na}_{2} \mathrm{CO}_{3}\)) is available in very pure form and can be used to standardize acid solutions. What is the molarity of a HCl solution if 28.3 mL of the solution are required to react with 0.256 g of \(\mathrm{Na}_{2} \mathrm{CO}_{3}\)?

A 3.664-g sample of a monoprotic acid was dissolved in water. It took 20.27 mL of a 0.1578 M NaOH solution to neutralize the acid. Calculate the molar mass of the acid.

Acetic acid (\(\mathrm{CH}_{3} \mathrm{COOH}\)) is an important ingredient of vinegar. A sample of 50.0 mL of a commercial vinegar is titrated against a 1.00 M NaOH solution. What is the concentration (in M) of acetic acid present in the vinegar if 5.75 mL of the base are needed for the titration?

A 15.00-mL solution of potassium nitrate (\(\mathrm{KNO}_{3}\)) was diluted to 125.0 mL, and 25.00 mL of this solution were then diluted to \(1.000 \times 10^{3}\) mL. The concentration of the final solution is 0.00383 M. Calculate the concentration of the original solution.

When 2.50 g of a zinc strip were placed in a \(\mathrm{AgNO}_{3}\) solution, silver metal formed on the surface of the strip. After some time had passed, the strip was removed from the solution, dried, and weighed. If the mass of the strip was 3.37 g. calculate the mass of Ag and Zn metals present.

The molecular formula of malonic acid is \(\mathrm{C}_{3} \mathrm{H}_{4} \mathrm{O}_{4}\). If a solution containing 0.762 g of the acid requires 12.44 mL of 1.174 M NaOH for neutralization, how many ionizable H atoms are present in the molecule?

A quantitative definition of solubility is the maximum number of grams of a solute that will dissolve in a given volume of water at a particular temperature. Describe an experiment that would enable you to determine the solubility of a soluble compound.

A 60.0-mL 0.513 M glucose (\(\mathrm{C}_{6} \mathrm{H}_{12} \mathrm{O}_{6}\)) solution is mixed with 120.0 mL of 2.33 M glucose solution. What is the concentration of the final solution? Assume the volumes are additive.

An ionic compound X is only slightly soluble in water. What test would you employ to show that the compound does indeed dissolve in water to a certain extent?

A student is given an unknown that is either iron(II) sulfate or iron (III) sulfate. Suggest a chemical procedure for determining its identity. (Both iron compounds are water soluble.)

You are given a colorless liquid. Describe three chemical tests you would perform on the liquid to show that it is water.

Give an example of a monoprotic acid, a diprotic acid, and a triprotic acid.

What are the characteristics of an acid–base neutralization reaction?

What factors qualify a compound as a salt? Specify which of the following compounds are salts: \(\mathrm{CH}_{4}\), NaF, NaOH, CaO, \(\mathrm{BaSO}_{4}\), \(\mathrm{HNO}_{3}\), \(\mathrm{NH}_{3}\), KBr?

Identify the following as a weak or strong acid or base: (a) \(\mathrm{NH}_{3}\), (b) \(\mathrm{H}_{3} \mathrm{PO}_{4}\), (c) LiOH, (d) HCOOH (formic acid), (e) \(\mathrm{H}_{2} \mathrm{SO}_{4}\), (f) HF, (g) \(\mathrm{Ba}(\mathrm{OH})_{2}\).

Identify each of the following species as a Brønsted acid, base, or both: (a) HI, (b) \(\mathrm{CH}_{3} \mathrm{COO}^{-}\), (c) \(\mathrm{H}_{2} \mathrm{PO}_{4}^{-}\) (d) \(\mathrm{HSO}_{4}^{-}\)

Identify each of the following species as a Brønsted acid, base, or both: \(\mathrm{PO}_{4}^{3-}\), (b) \(\mathrm{ClO}_{2}^{-}\), (c) \(\mathrm{NH}_{4}^{+}\), (d) \(\mathrm{HCO}_{3}^{-}\).

Balance the following equations and write the corresponding ionic and net ionic equations (if appropriate): (a) \(\mathrm{HBr}(a q)+\mathrm{NH}_{3}(a q) \rightarrow\) (b) \(\mathrm{Ba}(\mathrm{OH})_{2}(a q)+\mathrm{H}_{3} \mathrm{PO}_{4}(a q) \rightarrow\) (c) \(\mathrm{HClO}_{4}(a q)+\mathrm{Mg}(\mathrm{OH})_{2}(s) \rightarrow\)

Calculate the concentration of the acid (or base) remaining in solution when 10.7 mL of 0.211 M \(\mathrm{HNO}_{3}\), are added to 16.3 mL of 0.258 M NaOH.

(a) Describe a preparation for magnesium hydroxide \(\left[\mathrm{Mg}(\mathrm{OH})_2\right]\) and predict its solubility. (b) Milk of magnesia contains mostly \(\mathrm{Mg}(\mathrm{OH})_2\) and is effective in treating acid (mostly hydrochloric acid) indigestion. Calculate the volume of a \(0.035 \mathrm{M} \mathrm{HCl}\) solution (a typical acid concentration in an upset stomach) needed to react with two spoonfuls (approximately \(10 \mathrm{~mL}\)) of milk of magnesia [at \(0.080 \mathrm{~g}\) \(\left.\mathrm{Mg}(\mathrm{OH})_2 / \mathrm{mL}\right]\).

A 1.00-g sample of a metal X (that is known to form \(X^{2+}\) ions) was added to 0.100 L of 0.500 M \(\mathrm{H}_{2} \mathrm{SO}_{4}\). After all the metal had reacted, the remaining acid required 0.0334 L of 0.500 M NaOH solution for neutralization. Calculate the molar mass of the metal and identify the element.

Calculate the mass of the precipitate formed when 2.27 L of 0.0820 M \(\mathrm{Ba}(\mathrm{OH})_{2}\) are mixed with 3.06 L of 0.0664 M \(\mathrm{Na}_{2} \mathrm{SO}_{4}\).

Carbon dioxide in air can be removed by an aqueous metal hydroxide solution such as LiOH and \(B a(O H)_{2}\). (a) Write equations for the reactions. (Carbon dioxide reacts with water to form carbonic acid.) (b) Calculate the mass of \(\mathrm{CO}_{2}\) that can be removed by \(5.00 \times 10^{2}\) mL of a 0.800 M LiOH and a 0.800 M \(B a(O H)_{2}\) solution. (c) Which solution would you choose for use in a space capsule and which for use in a submarine?

Using the apparatus shown in Figure 4.1, a student found that a sulfuric acid solution caused the lightbulb to glow brightly. However, after the addition of a certain amount of a barium hydroxide [\(B a(O H)_{2}\)] solution, the light began to dim even though \(B a(O H)_{2}\) is also a strong electrolyte. Explain.

The molar mass of a certain metal carbonate, \(\mathrm{MCO}_{3}\), can be determined by adding an excess of HCl acid to react with all the carbonate and then “back titrating” the remaining acid with a NaOH solution. (a) Write an equation for these reactions, (b) In a certain experiment, 20.00 mL of 0.0800 M HCl were added to a 0.1022-g sample of \(\mathrm{MCO}_{3}\). The excess HCl required 5.64 mL of 0.1000 M NaOH for neutralization. Calculate the molar mass of the carbonate and identify M.

A 5.012-g sample of an iron chloride hydrate was dried in an oven. The mass of the anhydrous compound was 3.195 g. The compound was then dissolved in water and reacted with an excess of \(\mathrm{AgNO}_{3}\). The AgCl precipitate formed weighed 7.225 g. What is the formula of the original compound?

You are given a soluble compound of unknown molecular formula. (a) Describe three tests that would show that the compound is an acid. (b) Once you have established that the compound is an acid, describe how you would determine its molar mass using a NaOH solution of known concentration. (Assume the acid is monoprotic.) (c) How would you find out whether the acid is weak or strong? You are provided with a sample of NaCl and an apparatus like that shown in Figure 4.1 for comparison.

You are given two colorless solutions, one containing NaCl and the other sucrose (\(\mathrm{C}_{12} \mathrm{H}_{22} \mathrm{O}_{11}\)). Suggest a chemical and a physical test that would allow you to distinguish between these two solutions.

The concentration of lead ions (\(P b^{2+}\)) in a sample of polluted water that also contains nitrate ions (\(\mathrm{NO}_{3}^{-}\)) is determined by adding solid sodium sulfate (\(\mathrm{Na}_{2} \mathrm{SO}_{4}\)) to exactly 500 mL of the water. (a) Write the molecular and net ionic equations for the reaction. (b) Calculate the molar concentration of \(P b^{2+}\) if 0.00450 g of \(\mathrm{Na}_{2} \mathrm{SO}_{4}\) was needed for the complete precipitation of \(P b^{2+}\) ions as \(\mathrm{PbSO}_{4}\).

Hydrochloric acid is not an oxidizing agent in the sense that sulfuric acid and nitric acid are. Explain why the chloride ion is not a strong oxidizing agent like \(\mathrm{SO}_{4}^{2-}\) and \(\mathrm{NO}_{3}^{-}\)

Explain how you would prepare potassium iodide (KI) by means of (a) an acid–base reaction and (b) a reaction between an acid and a carbonate compound.

Sodium reacts with water to yield hydrogen gas. Why is this reaction not used in the laboratory preparation of hydrogen?

Chlorine forms a number of oxides with the following oxidation numbers: +1, +3, +4, +6. and +7. Write a formula for each of these compounds.

A useful application of oxalic acid is the removal of rust (\(\mathrm{Fe}_{2} \mathrm{O}_{3}\)) from, say, bathtub rings according to the reaction \(\mathrm{Fe}_{2} \mathrm{O}_{3}(\mathrm{~S})+6 \mathrm{H}_{2} \mathrm{C}_{2} \mathrm{O}_{4}(\mathrm{aq}) \rightarrow 2 \mathrm{Fe}\left(\mathrm{C}_{2} \mathrm{O}_{4}\right)_{3}^{3-}(\mathrm{aq})+3 \mathrm{H}_{2} \mathrm{O}+6 \mathrm{H}^{+}(\mathrm{aq})\) Calculate the number of grams of rust that can be removed by \(5.00 \times 10^{2}\) mL of a 0.100 M solution of oxalic acid.

Acetylsalicylic acid (\(\mathrm{C}_{9} \mathrm{H}_{8} \mathrm{O}_{4}\)) is a monoprotic acid commonly known as “aspirin.” A typical aspirin tablet, however, contains only a small amount of the acid. In an experiment to determine its composition, an aspirin tablet was crushed and dissolved in water. It took 12.25 mL of 0.1466 M NaOH to neutralize the solution. Calculate the number of grains of aspirin in the tablet. (One grain = 0.0648 g.)

A NaoH is initially mixed with an acid solution shown in (a). Which of the diagrams shown in (b) - (d) corresponds to one of the following acids: HCl, \(\mathrm{H}_{2} \mathrm{SO}_{4}\), \(\mathrm{H}_{3} \mathrm{PO}_{4}\)? Color codes: Blue spheres (\(\mathrm{OH}^{-}\) ions); red spheres (acid molecules); green spheres (anions of the acids). Assume all the acid-base neutralization reactions go to completion.

A 0.8870-g sample of a mixture of NaCl and KCl is dissolved in water, and the solution is then treated with an excess of \(\mathrm{AgNO}_{3}\) to yield 1.913 g of AgCl. Calculate the percent by mass of each compound in the mixture.

Based on oxidation number consideration, explain why carbon monoxide (CO) is flammable but carbon dioxide (\(\mathrm{CO}_{2}\)) is not.

Which of the diagrams shown here corresponds to the reaction between AgOH(s) and \(\mathrm{HNO}_3(a q)\) ? Write a balanced equation for the reaction. The green spheres represent the \(\mathrm{Ag}^{+}\) and the red spheres represent the \(\mathrm{NO}_3^{-}\) ions.

Muriatic acid, a commercial-grade hydrochloric acid used for cleaning masonry surfaces, is typically around 10 percent HCl by mass and has a density of 1.2 g/\(\mathrm{cm}^{3}\). A 0.5-in layer of boiler scale has accumulated on a 6.0-ft section of hot water pipe with an internal diameter of 2.0 in (see the Chemistry in Action on p. 126). What is the minimum volume of muriatic acid in gallons that would be needed to remove the boiler scale?