Chapter 5

Following are some laboratory methods occasionally used for the preparation of small quantities of chemicals. Write a balanced equation for each. (a) preparation of \(\mathrm{H}_{2} \mathrm{S}(\mathrm{g}): \mathrm{HCl}(\mathrm{aq})\) is heated with \(\mathrm{FeS}(\mathrm{s})\) (b) preparation of \(\mathrm{Cl}_{2}(\mathrm{g}): \mathrm{HCl}(\mathrm{aq})\) is heated with \(\mathrm{MnO}_{2}(\mathrm{s}) ; \mathrm{MnCl}_{2}(\mathrm{aq})\) and \(\mathrm{H}_{2} \mathrm{O}(1)\) are other products (c) preparation of \(\mathrm{N}_{2}: \mathrm{Br}_{2}\) and \(\mathrm{NH}_{3}\) react in aqueous solution; \(\mathrm{NH}_{4} \mathrm{Br}\) is another product (d) preparation of chlorous acid: an aqueous suspension of solid barium chlorite is treated with dilute \(\mathrm{H}_{2} \mathrm{SO}_{4}(\mathrm{aq})\)

When concentrated \(\mathrm{CaCl}_{2}(\mathrm{aq})\) is added to \(\mathrm{Na}_{2} \mathrm{HPO}_{4}(\mathrm{aq}),\) a white precipitate forms that is \(38.7 \%\) Ca by mass. Write a net ionic equation representing the probable reaction that occurs.

Sodium hydroxide used to make standard \(\mathrm{NaOH}(\mathrm{aq})\) solutions for acid-base titrations is invariably contaminated with some sodium carbonate. (a) Explain why, except in the most precise work, the presence of this sodium carbonate generally does not seriously affect the results obtained, for example, when \(\mathrm{NaOH}(\mathrm{aq})\) is used to titrate HCl(aq). (b) Conversely, show that if \(\mathrm{Na}_{2} \mathrm{CO}_{3}\) comprises more than \(1 \%\) to \(2 \%\) of the solute in NaOH(aq), the titration results are affected.

A \(110.520 \mathrm{g}\) sample of mineral water is analyzed for its magnesium content. The \(\mathrm{Mg}^{2+}\) in the sample is first precipitated as \(\mathrm{MgNH}_{4} \mathrm{PO}_{4},\) and this precipitate is then converted to \(\mathrm{Mg}_{2} \mathrm{P}_{2} \mathrm{O}_{7},\) which is found to weigh 0.0549 g. Express the quantity of magnesium in the sample in parts per million (that is, in grams of \(\mathrm{Mg}\) per million grams of \(\mathrm{H}_{2} \mathrm{O}\) ).

What volume of \(0.248 \mathrm{M} \mathrm{CaCl}_{2}\) must be added to \(335 \mathrm{mL}\) of \(0.186 \mathrm{M} \mathrm{KCl}\) to produce a solution with a concentration of \(0.250 \mathrm{M} \mathrm{Cl}^{-2}\) Assume that the solution volumes are additive.

An unknown white solid consists of two compounds, each containing a different cation. As suggested in the illustration, the unknown is partially soluble in water. The solution is treated with \(\mathrm{NaOH}(\mathrm{aq})\) and yields a white precipitate. The part of the original solid that is insoluble in water dissolves in \(\mathrm{HCl}(\mathrm{aq})\) with the evolution of a gas. The resulting solution is then treated with \(\left(\mathrm{NH}_{4}\right)_{2} \mathrm{SO}_{4}(\mathrm{aq})\) and yields a white precipitate. (a) Is it possible that any of the cations \(M g^{2+}, C u^{2+}\) \(\mathrm{Ba}^{2+}, \mathrm{Na}^{+},\) or \(\mathrm{NH}_{4}^{+}\) were present in the original unknown? Explain your reasoning. (b) What compounds could be in the unknown mixture (that is, what anions might be present)?

Balance these equations for reactions in acidic solution. (a) \(\mathrm{IBr}+\mathrm{BrO}_{3}^{-}+\mathrm{H}^{+} \longrightarrow \mathrm{IO}_{3}^{-}+\mathrm{Br}^{-}+\mathrm{H}_{2} \mathrm{O}\) (b) \(\mathrm{C}_{2} \mathrm{H}_{5} \mathrm{NO}_{3}+\mathrm{Sn} \longrightarrow\) \(\mathrm{NH}_{2} \mathrm{OH}+\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{OH}+\mathrm{Sn}^{2+}\) (c) \(\mathrm{As}_{2} \mathrm{S}_{3}+\mathrm{NO}_{3}^{-} \longrightarrow \mathrm{H}_{3} \mathrm{AsO}_{4}+\mathrm{S}+\mathrm{NO}\) (d) \(\mathrm{H}_{5} \mathrm{IO}_{6}+\mathrm{I}_{2} \longrightarrow \mathrm{IO}_{3}^{-}+\mathrm{H}^{+}+\mathrm{H}_{2} \mathrm{O}\) (e) \(\mathrm{S}_{2} \mathrm{F}_{2}+\mathrm{H}_{2} \mathrm{O} \longrightarrow \mathrm{S}_{8}+\mathrm{H}_{2} \mathrm{S}_{4} \mathrm{O}_{6}+\mathrm{HF}\)

A method of producing phosphine, \(\mathrm{PH}_{3}\), from elemental phosphorus, \(P_{4}\), involves heating the \(P_{4}\) with \(\mathrm{H}_{2} \mathrm{O} .\) An additional product is phosphoric acid, \(\mathrm{H}_{3} \mathrm{PO}_{4}\) Write a balanced equation for this reaction.

Iron (Fe) is obtained from rock that is extracted from open pit mines and then crushed. The process used to obtain the pure metal from the crushed rock produces solid waste, called tailings, which are stored in disposal areas near the mines. The tailings pose a serious environmental risk because they contain sulfides, such as pyrite ( \(\mathrm{FeS}_{2}\) ), which oxidize in air to produce metal ions and \(\mathrm{H}^{+}\) ions that can enter into surface water or ground water. The oxidation of \(\mathrm{FeS}_{2}\) to \(\mathrm{Fe}^{3+}\) is described by the unbalanced chemical equation below. \(\mathrm{FeS}_{2}(\mathrm{s})+\mathrm{O}_{2}(\mathrm{g})+\mathrm{H}_{2} \mathrm{O}(\mathrm{l}) \longrightarrow\) \(\quad \mathrm{Fe}^{3+}(\mathrm{aq})+\mathrm{SO}_{4}^{2-}(\mathrm{aq})+\mathrm{H}^{+}(\mathrm{aq}) \quad(\text { not balanced })\) Thus, the oxidation of pyrite produces \(\mathrm{Fe}^{3+}\) and \(\mathrm{H}^{+}\) ions that can leach into surface or ground water. The leaching of \(\mathrm{H}^{+}\) ions causes the water to become very acidic. To prevent acidification of nearby ground or surface water, limestone \(\left(\mathrm{CaCO}_{3}\right)\) is added to the tailings to neutralize the \(\mathrm{H}^{+}\) ions: \(\mathrm{CaCO}_{3}(\mathrm{s})+2 \mathrm{H}^{+}(\mathrm{aq}) \underset{\mathrm{Ca}^{2+}}{\longrightarrow}(\mathrm{aq})+\mathrm{H}_{2} \mathrm{O}(\mathrm{l})+\mathrm{CO}_{2}(\mathrm{g})\) (a) Balance the equation above for the reaction of \(\mathrm{FeS}_{2}\) and \(\mathrm{O}_{2}\). [ Hint: Start with the half-equations \(\mathrm{FeS}_{2}(\mathrm{s}) \rightarrow\) \(\left.\mathrm{Fe}^{3+}(\mathrm{aq})+\mathrm{SO}_{4}^{2-}(\mathrm{aq}) \text { and } \mathrm{O}_{2}(\mathrm{g}) \rightarrow \mathrm{H}_{2} \mathrm{O}(1) .\right]\) (b) What is the minimum amount of \(\mathrm{CaCO}_{3}(\mathrm{s})\) required, per kilogram of tailings, to prevent contamination if the tailings contain \(3 \%\) S by mass? Assume that all the sulfur in the tailings is in the form \(\mathrm{FeS}_{2}\).

A sample of battery acid is to be analyzed for its sulfuric acid content. A \(1.00 \mathrm{mL}\) sample weighs \(1.239 \mathrm{g}\). This \(1.00 \mathrm{mL}\) sample is diluted to \(250.0 \mathrm{mL}\), and \(10.00 \mathrm{mL}\) of this diluted acid requires \(32.44 \mathrm{mL}\) of \(0.00498 \mathrm{M} \mathrm{Ba}(\mathrm{OH})_{2}\) for its titration. What is the mass percent of \(\mathrm{H}_{2} \mathrm{SO}_{4}\) in the battery acid? (Assume that complete ionization and neutralization of the \(\mathrm{H}_{2} \mathrm{SO}_{4}\) occurs.)

A piece of marble (assume it is pure \(\mathrm{CaCO}_{3}\) ) reacts with \(2.00 \mathrm{L}\) of \(2.52 \mathrm{M} \mathrm{HCl}\). After dissolution of the marble, a \(10.00 \mathrm{mL}\) sample of the resulting solution is withdrawn, added to some water, and titrated with 24.87 mL of 0.9987 M NaOH. What must have been the mass of the piece of marble? Comment on the precision of this method; that is, how many significant figures are justified in the result?

The active component in one type of calcium dietary supplement is calcium carbonate. A \(1.2450 \mathrm{g}\) tablet of the supplement is added to \(50.00 \mathrm{mL}\) of \(0.5000 \mathrm{M} \mathrm{HCl}\) and allowed to react. After completion of the reaction, the excess HCl(aq) requires \(40.20 \mathrm{mL}\) of \(0.2184 \mathrm{M}\) NaOH for its titration to the equivalence point. What is the calcium content of the tablet, expressed in milligrams of \(\mathrm{Ca}^{2+} ?\)

A \(0.4324 \mathrm{g}\) sample of a potassium hydroxide-lithium hydroxide mixture requires \(28.28 \mathrm{mL}\) of \(0.3520 \mathrm{M} \mathrm{HCl}\) for its titration to the equivalence point. What is the mass percent lithium hydroxide in this mixture?

Chile saltpeter is a natural source of \(\mathrm{NaNO}_{3}\); it also contains \(\mathrm{NaIO}_{3} .\) The \(\mathrm{NaIO}_{3}\) can be used as a source of iodine. Iodine is produced from sodium iodate in a two-step process occurring under acidic conditions: \(\begin{aligned} \mathrm{IO}_{3}^{-}(\mathrm{aq})+\mathrm{HSO}_{3}^{-}(\mathrm{aq}) & \longrightarrow \mathrm{I}^{-}(\mathrm{aq}) +\mathrm{SO}_{4}^{2-}(\mathrm{aq}) \end{aligned} \quad\) ( not balanced) \(\mathrm{I}^{-}(\mathrm{aq})+\mathrm{IO}_{3}^{-}(\mathrm{aq}) \longrightarrow\) \(\mathrm{I}_{2}(\mathrm{s})+\mathrm{H}_{2} \mathrm{O}(\mathrm{l}) \quad(\text { not balanced })\) In the illustration, a 5.00 L sample of a \(\mathrm{NaIO}_{3}(\mathrm{aq})\) solution containing \(5.80 \mathrm{g} \mathrm{NaIO}_{3} / \mathrm{L}\) is treated with the stoichiometric quantity of \(\mathrm{NaHSO}_{3}\) (no excess of either reactant). Then, a further quantity of the initial \(\mathrm{NaIO}_{3}(\mathrm{aq})\) is added to the reaction mixture to bring about the second reaction. (a) How many grams of NaHSO \(_{3}\) are required in the first step? (b) What additional volume of the starting solution must be added in the second step?

The active ingredients in a particular antacid tablet are aluminum hydroxide, \(\mathrm{Al}(\mathrm{OH})_{3},\) and magnesium hydroxide, \(\mathrm{Mg}(\mathrm{OH})_{2} . \quad \mathrm{A} 5.00 \times 10^{2} \mathrm{mg}\) sample of the active ingredients was dissolved in \(50.0 \mathrm{mL}\) of \(0.500 \mathrm{M} \mathrm{HCl} .\) The resulting solution, which was still acidic, required \(16.5 \mathrm{mL}\) of \(0.377 \mathrm{M} \mathrm{NaOH}\) for neutralization. What are the mass percentages of \(\mathrm{Al}(\mathrm{OH})_{3}\) and \(\mathrm{Mg}(\mathrm{OH})_{2}\) in the sample?

A compound contains only Fe and O. A \(0.2729 \mathrm{g}\) sample of the compound was dissolved in \(50 \mathrm{mL}\) of concentrated acid solution, reducing all the iron to \(\mathrm{Fe}^{2+}\) ions. The resulting solution was diluted to \(100 \mathrm{mL}\) and then titrated with a \(0.01621 \mathrm{M} \mathrm{KMnO}_{4}\) solution. The unbalanced chemical equation for reaction between \(\mathrm{Fe}^{2+}\) and \(\mathrm{MnO}_{4}^{-}\) is given below. \(\begin{aligned} \mathrm{MnO}_{4}^{-}(\mathrm{aq})+& \mathrm{Fe}^{2+}(\mathrm{aq}) \longrightarrow \mathrm{Mn}^{2+}(\mathrm{aq})+\mathrm{Fe}^{3+}(\mathrm{aq}) \quad(\text { not balanced }) \end{aligned}\) The titration required \(42.17 \mathrm{mL}\) of the \(\mathrm{KMnO}_{4}\) solution to reach the pink endpoint. What is the empirical formula of the compound?

Warfarin, \(\mathrm{C}_{19} \mathrm{H}_{16} \mathrm{O}_{4},\) is the active ingredient used in some anticoagulant medications. The amount of warfarin in a particular sample was determined as follows. A 13.96 g sample was first treated with an alkaline I_ solution to convert \(\mathrm{C}_{19} \mathrm{H}_{16} \mathrm{O}_{4}\) to \(\mathrm{CHI}_{3}\). This treatment gives one mole of \(\mathrm{CHI}_{3}\) for every mole of \(\mathrm{C}_{19} \mathrm{H}_{16} \mathrm{O}_{4}\) that was initially present in the sample. The iodine in \(\mathrm{CHI}_{3}\) is then precipitated as \(\mathrm{AgI}(\mathrm{s})\) by treatment with excess \(\mathrm{AgNO}_{3}(\mathrm{aq}):\) $$\begin{aligned} \mathrm{CHI}_{3}(\mathrm{aq})+3 \mathrm{AgNO}_{3}(\mathrm{aq})+\mathrm{H}_{2} \mathrm{O}(\mathrm{l}) \longrightarrow & \longrightarrow 3 \mathrm{AgI}(\mathrm{s})+3 \mathrm{HNO}_{3}(\mathrm{aq}) &+\mathrm{CO}(\mathrm{g}) \end{aligned}$$ If \(0.1386 \mathrm{g}\) solid \(\mathrm{AgI}\) were obtained, then what is the percentage by mass of warfarin in the sample analyzed?

Phosphorus is essential for plant growth, but an excess of phosphorus can be catastrophic in aqueous ecosystems. Too much phosphorus can cause algae to grow at an explosive rate and this robs the rest of the ecosystem of oxygen. Effluent from sewage treatment plants must be treated before it can be released into lakes or streams because the effluent contains significant amounts of \(\mathrm{H}_{2} \mathrm{PO}_{4}^{-}\) and \(\mathrm{HPO}_{4}^{2-}\). (Detergents are a major contributor to phosphorus levels in domestic sewage because many detergents contain \(\mathrm{Na}_{2} \mathrm{HPO}_{4}\) ) A simple way to remove \(\mathrm{H}_{2} \mathrm{PO}_{4}^{-}\) and \(\mathrm{HPO}_{4}^{2-}\) from the effluent is to treat it with lime, \(\mathrm{CaO}\) which produces \(\mathrm{Ca}^{2+}\) and \(\mathrm{OH}^{-}\) ions in water. The \(\mathrm{OH}^{-}\) ions convert \(\mathrm{H}_{2} \mathrm{PO}_{4}^{-}\) and \(\mathrm{HPO}_{4}^{2-}\) ions into \(\mathrm{PO}_{4}^{3-}\) ions and, finally, \(\mathrm{Ca}^{2+}, \mathrm{OH}^{-}\), and \(\mathrm{PO}_{4}^{3-}\) ions combine to form a precipitate of \(\mathrm{Ca}_{5}\left(\mathrm{PO}_{4}\right)_{3} \mathrm{OH}(\mathrm{s})\) (a) Write balanced chemical equations for the four reactions described above. [Hint: The reactants are \(\mathrm{CaO}\) and \(\mathrm{H}_{2} \mathrm{O} ; \mathrm{H}_{2} \mathrm{PO}_{4}^{-}\) and \(\left.\mathrm{OH}^{-} ; \mathrm{HPO}_{4}^{2-} \text { and } \mathrm{OH}^{-} ; \mathrm{Ca}^{2+}, \mathrm{PO}_{4}^{3-}, \text { and } \mathrm{OH}^{-} .\right]\) (b) How many kilograms of lime are required to remove the phosphorus from a \(1.00 \times 10^{4}\) L holding tank filled with contaminated water, if the water contains \(10.0 \mathrm{mg}\) of phosphorus per liter?

Blood alcohol content (BAC) is often reported in weight-volume percent (w/v\%). For example, a BAC of \(0.10 \%\) corresponds to \(0.10 \mathrm{g} \mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{OH}\) per 100 mL of blood. Estimates of BAC can be obtained from breath samples by using a number of commercially available instruments, including the Breathalyzer for which a patent was issued to R. F. Borkenstein in 1958\. The chemistry behind the Breathalyzer is described by the oxidation- reduction reaction below, which occurs in acidic solution: \(\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{OH}(\mathrm{g})+\mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}(\mathrm{aq}) \longrightarrow\) \(\mathrm{CH}_{3} \mathrm{COOH}(\mathrm{aq})+\mathrm{Cr}^{3+}(\mathrm{aq}) \quad(\text { not balanced })\)A Breathalyzer instrument contains two ampules, each of which contains \(0.75 \mathrm{mg} \mathrm{K}_{2} \mathrm{Cr}_{2} \mathrm{O}_{7}\) dissolved in \(3 \mathrm{mL}\) of \(9 \mathrm{mol} / \mathrm{L} \mathrm{H}_{2} \mathrm{SO}_{4}(\mathrm{aq}) .\) One of the ampules is used as reference. When a person exhales into the tube of the Breathalyzer, the breath is directed into one of the ampules, and ethyl alcohol in the breath converts \(\mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}\) into \(\mathrm{Cr}^{3+} .\) The instrument compares the colors of the solutions in the two ampules to determine the breath alcohol content (BrAC), and then converts this into an estimate of BAC. The conversion of BrAC into BAC rests on the assumption that 2100 mL of air exhaled from the lungs contains the same amount of alcohol as \(1 \mathrm{mL}\) of blood. With the theory and assumptions described in this problem, calculate the molarity of \(\mathrm{K}_{2} \mathrm{Cr}_{2} \mathrm{O}_{7}\) in the ampules before and after a breath test in which a person with a BAC of \(0.05 \%\) exhales 0.500 Lof his breath into a Breathalyzer instrument.

In your own words, define or explain the terms or symbols \((\mathrm{a}) \rightleftharpoons(\mathrm{b})[] ;(\mathrm{c})\) spectator ion; (d) weak acid.

Briefly describe (a) half-equation method of balancing redox equations; (b) disproportionation reaction; (c) titration; (d) standardization of a solution.

Explain the important distinctions between (a) a strong electrolyte and strong acid; (b) an oxidizing agent and reducing agent; (c) precipitation reactions and neutralization reactions; (d) half-reaction and overall reaction.

The number of moles of hydroxide ion in 0.300 L of \(0.0050 \mathrm{M} \mathrm{Ba}(\mathrm{OH})_{2}\) is (a) \(0.0015 ;(\mathrm{b}) 0.0030 ;(\mathrm{c}) 0.0050\) (d) 0.010.

The highest \(\left[\mathrm{H}^{+}\right]\) will be found in an aqueous solution that is (a) \(0.10 \mathrm{M} \mathrm{HCl} ;\) (b) \(0.10 \mathrm{M} \mathrm{NH}_{3} ;\) (c) \(0.15 \mathrm{M}\) \(\mathrm{CH}_{3} \mathrm{COOH} ;(\mathrm{d}) 0.10 \mathrm{M} \mathrm{H}_{2} \mathrm{SO}_{4}\).

To precipitate \(\mathrm{Zn}^{2+}\) from \(\mathrm{Zn}\left(\mathrm{NO}_{3}\right)_{2}(\mathrm{aq}),\) add (a) \(\mathrm{NH}_{4} \mathrm{Cl} ;\) (b) \(\mathrm{MgBr}_{2} ;\) (c) \(\mathrm{K}_{2} \mathrm{CO}_{3} ;\) (d) \(\left(\mathrm{NH}_{4}\right)_{2} \mathrm{SO}_{4}\).

When treated with dilute \(\mathrm{HCl}(\mathrm{aq}),\) the solid that reacts to produce a gas is (a) \(\mathrm{BaSO}_{3} ;\) (b) \(\mathrm{ZnO};\) (c) \(\mathrm{NaBr} ;\) (d) \(\mathrm{Na}_{2} \mathrm{SO}_{4}\).

What is the net ionic equation for the reaction that occurs when an aqueous solution of \(\mathrm{KI}\) is added to an aqueous solution of \(\mathrm{Pb}\left(\mathrm{NO}_{3}\right)_{2} ?\)

When aqueous sodium carbonate, \(\mathrm{Na}_{2} \mathrm{CO}_{3}\), is treated with dilute hydrochloric acid, HCl, the products are sodium chloride, water, and carbon dioxide gas. What is the net ionic equation for this reaction?

Describe the synthesis of each of the following ionic compounds, starting from solutions of sodium and nitrate salts. Then write the net ionic equation for each synthesis. (a) \(\mathrm{Zn}_{3}\left(\mathrm{PO}_{4}\right)_{2};\) (b) \(\mathrm{Cu}(\mathrm{OH})_{2};\) (c) \(\mathrm{NiCO}_{3}.\)

Consider the following redox reaction: $$\begin{array}{r}4 \mathrm{NO}(\mathrm{g})+3 \mathrm{O}_{2}(\mathrm{g})+2 \mathrm{H}_{2} \mathrm{O}(\mathrm{l}) \longrightarrow 4 \mathrm{NO}_{3}^{-}(\mathrm{aq})+4 \mathrm{H}^{+}(\mathrm{aq})\end{array} $$ (a) Which species is oxidized? (b) Which species is reduced? (c) Which species is the oxidizing agent? (d) Which species is the reducing agent? (e) Which species gains electrons? (f) Which species loses electrons?

In the equation \(\begin{aligned} ? \mathrm{Fe}^{2+}(\mathrm{aq})+\mathrm{O}_{2}(\mathrm{g})+4 \mathrm{H}^{+}(\mathrm{aq}) & \longrightarrow ? \mathrm{Fe}^{3+}(\mathrm{aq})+2 \mathrm{H}_{2} \mathrm{O}(1) \end{aligned}\) the missing coefficients (a) are each \(2 ;\) (b) are each 4; (c) can have any values as long as they are the same; (d) must be determined by experiment.

In the half-reaction in which \(\mathrm{NpO}_{2}^{+}\) is converted to \(\mathrm{Np}^{4+},\) the number of electrons appearing in the half-equation is (a) \(1 ;(b) 2 ;(c) 3 ;\) (d) 4.

Classify each of the following statements as true or false. (a) Barium chloride, \(\mathrm{BaCl}_{2^{\prime}}\) is a weak electrolyte in aqueous solution. (b) In the reaction \(\mathrm{H}^{-}(\mathrm{aq})+\mathrm{H}_{2} \mathrm{O}(1) \rightarrow \mathrm{H}_{2}(\mathrm{g})+\) \(\mathrm{OH}^{-}(\mathrm{aq}),\) water acts as both an acid and an oxidizing agent. (c) A precipitate forms when aqueous sodium carbonate, \(\mathrm{Na}_{2} \mathrm{CO}_{3}(\mathrm{aq}),\) is treated with excess aqueous hydrochloric acid, HCl(aq). (d) Hydrofluoric acid, \(\overline{\mathrm{HF}}\), is a strong acid in water. (e) Compared with a 0.010 M solution of \(\mathrm{NaNO}_{3}\), a \(0.010 \mathrm{M}\) solution of \(\mathrm{Mg}\left(\mathrm{NO}_{3}\right)_{2}\) is a better conductor of electricity.

Which of the following reactions are oxidationreduction reactions? (a) \(\mathrm{H}_{2} \mathrm{CO}_{3}(\mathrm{aq}) \longrightarrow \mathrm{H}_{2} \mathrm{O}(1)+\mathrm{CO}_{2}(\mathrm{g})\) (b) \(2 \mathrm{Li}(\mathrm{s})+2 \mathrm{H}_{2} \mathrm{O}(1) \longrightarrow 2 \mathrm{LiOH}(\mathrm{aq})+\mathrm{H}_{2}(\mathrm{g})\) (c) \(4 \mathrm{Ag}(\mathrm{s})+\mathrm{PtCl}_{4}(\mathrm{aq}) \longrightarrow 4 \mathrm{AgCl}(\mathrm{s})+\mathrm{Pt}(\mathrm{s})\) (d) \(2 \mathrm{HClO}_{4}(\mathrm{aq})+\mathrm{Ca}(\mathrm{OH})_{2}(\mathrm{aq}) \longrightarrow\) \(2 \mathrm{H}_{2} \mathrm{O}(1)+\mathrm{Ca}\left(\mathrm{ClO}_{4}\right)_{2}(\mathrm{aq})\)