Chapter 8

Eutrophication, the rapid growth of algae and the death of fish, may be caused by the presence of an excess of phosphates in water. Treatment plants that process sewage may add \(\mathrm{Ca}(\mathrm{OH})_{2}\) (slaked lime) to water to remove phosphates before returning the water to the environment. Although the phosphates may be present in several forms, we can use HPO \(_{4}^{-}\) as a representative phosphate in the net ionic equation: \(5 \mathrm{Ca}^{2+}(a q)+3 \mathrm{HPO}_{4}^{-}(a q)+4 \mathrm{OH}^{-}(a q) \rightarrow\) $$ \mathrm{Ca}_{5} \mathrm{OH}\left(\mathrm{PO}_{4}\right)_{3}(s)+3 \mathrm{H}_{2} \mathrm{O}(\ell) $$ If phosphates (as HPO \(_{4}^{-}\) ) are present at a level of \(15.7 \mathrm{mg} / \mathrm{L}\) in wastewater, how much \(\mathrm{Ca}(\mathrm{OH})_{2}\) would need to be added to \(1.00 \times 10^{5} \mathrm{L}\) of water to precipitate \(95 \%\) of the phosphate ion present?

Iron(II) can be precipitated from a slightly basic aqueous solution by bubbling oxygen through the solution, which converts \(\mathrm{Fe}^{2+}\) to insoluble \(\mathrm{Fe}^{3+}\) \(4 \mathrm{Fe}(\mathrm{OH})^{+}(a q)+4 \mathrm{OH}^{-}(a q)+\mathrm{O}_{2}(g)+2 \mathrm{H}_{2} \mathrm{O}(\ell) \rightarrow\) $$ 4 \mathrm{Fe}(\mathrm{OH})_{3}(s) $$ How many grams of \(\mathrm{O}_{2}\) are consumed to precipitate all of the iron in \(75 \mathrm{mL}\) of \(0.090 M \mathrm{Fe}^{2+} ?\)

Given the following equation, how many grams of \(\mathrm{Pb} \mathrm{CO}_{3}\) will dissolve when 1.00 L of \(1.00 M \mathrm{H}^{+}\) is added to \(5.00 \mathrm{g}\) of \(\mathrm{PbCO}_{3} ?\) $$ \mathrm{PbCO}_{3}(s)+2 \mathrm{H}^{+}(a q) \rightarrow \mathrm{Pb}^{2+}(a q)+\mathrm{H}_{2} \mathrm{O}(\ell)+\mathrm{CO}_{2}(g) $$

Rhubarb leaves contain \(0.520 \mathrm{g}\) of oxalic \(\operatorname{acid}\left(\mathrm{H}_{2} \mathrm{C}_{2} \mathrm{O}_{4}\right)\) per \(100.0 \mathrm{g}\) of leaves. The oxalic acid can react with calcium ion to make insoluble calcium oxalate, a major constituent of kidney stones. To study this reaction in the laboratory, 375 mL of a \(0.866 M\) solution of oxalic acid is treated with excess \(0.133 M\) calcium hydroxide solution. a. How much calcium oxalate is formed in this reaction? "b. What volume of calcium hydroxide solution is required if it must be present in \(20 \%\) excess to ensure complete reaction?

How are the gains or losses of electrons related to changes in oxidation numbers?

What is the sum of the oxidation numbers of the atoms in a molecule?

What is the sum of the oxidation numbers of all the atoms in each of the following polyatomic ions? (a) \(\mathrm{OH}^{-}\) (b) \(\mathrm{NH}_{4}^{+} ;\) (c) \(\mathrm{SO}_{4}^{2-} ;\) (d) \(\mathrm{PO}_{4}^{3-}\)

Gold does not dissolve in concentrated \(\mathrm{H}_{2} \mathrm{SO}_{4}\) but readily dissolves in \(\mathrm{H}_{2} \mathrm{SeO}_{4}\) (selenic acid). Which acid is the stronger oxidizing agent?

Silver dissolves in sulfuric acid to form silver sulfate and \(\mathrm{H}_{2},\) but gold does not dissolve in sulfuric acid to form gold sulfate. Which of the two metals is the better reducing agent?

Rank the following oxoanions in order of decreasing oxidation number of chlorine: (a) \(\mathrm{ClO}^{-} ;\) (b) \(\mathrm{ClO}_{2}^{-}\) (c) \(\mathrm{ClO}_{3}^{-} ;\) (d) \(\mathrm{ClO}_{4}^{-}\)

Why is the oxidizing agent in a redox reaction reduced and the reducing agent oxidized?

What is the oxidation number of chlorine in each of the following oxoacids? (a) hypochlorous acid (HC1O); (b) chloric acid \(\left(\mathrm{HClO}_{3}\right) ;\) (c) perchloric acid \(\left(\mathrm{HClO}_{4}\right)\)

What is the oxidation number of nitrogen in each of the following species? (a) elemental nitrogen \(\left(\mathrm{N}_{2}\right) ;\) (b) hydrazine \(\left(\mathrm{N}_{2} \mathrm{H}_{4}\right) ;(\mathrm{c})\) ammonium ion \(\left(\mathrm{NH}_{4}^{+}\right)\)

What is the change (if any) in the oxidation state of carbon in this reaction? $$ \mathrm{C}_{12} \mathrm{H}_{22} \mathrm{O}_{11}(s) \rightarrow 12 \mathrm{C}(s)+11 \mathrm{H}_{2} \mathrm{O}(\ell) $$

In one stage of the nitrogen cycle, Nitrosomonas bacteria convert ammonia and oxygen into nitrite ions. a. What is the change in oxidation state of nitrogen during the reaction? b. Write a balanced net ionic equation for the reaction in acidic groundwater.

Iron is oxidized in a number of chemical weathering processes. How many moles of \(\mathrm{O}_{2}\) are consumed when one mole of magnetite (Fe \(_{3} \mathrm{O}_{4}\) ) is converted into hematite \(\left(\mathrm{Fe}_{2} \mathrm{O}_{3}\right) ?\)

The mineral rhodochrosite [manganese(II) carbonate, \(\left.\mathrm{MnCO}_{3}\right]\) is a commercially important source of manganese. How many moles of \(\mathrm{O}_{2}\) are consumed when one mole \(\mathrm{MnCO}_{3}\) is converted into \(\mathrm{MnO}_{2}\) and \(\mathrm{CO}_{2} ?\)

The following chemical reactions have helped to shape Earth's crust. Determine the oxidation numbers of all the elements in the reactants and products, and identify which elements are oxidized and which are reduced: a. \(3 \mathrm{SiO}_{2}(s)+2 \mathrm{Fe}_{3} \mathrm{O}_{4}(s) \rightarrow 3 \mathrm{Fe}_{2} \mathrm{SiO}_{4}(s)+\mathrm{O}_{2}(g)\) b. \(\operatorname{si} \mathrm{O}_{2}(s)+2 \mathrm{Fe}(s)+\mathrm{O}_{2}(g) \rightarrow \mathrm{Fe}_{2} \mathrm{SiO}_{4}(s)\) c. \(4 \mathrm{FeO}(s)+\mathrm{O}_{2}(g)+6 \mathrm{H}_{2} \mathrm{O}(\ell) \rightarrow 4 \mathrm{Fe}(\mathrm{OH})_{3}(s)\)

Determine the oxidation numbers of each of the elements in the following reactions, and identify which of them, if any, are oxidized or reduced: a. \(\mathrm{SiO}_{2}(s)+2 \mathrm{H}_{2} \mathrm{O}(\ell) \rightarrow \mathrm{H}_{4} \mathrm{SiO}_{4}(a q)\) b. \(2 \mathrm{MnCO}_{3}(s)+\mathrm{O}_{2}(g) \rightarrow 2 \mathrm{MnO}_{2}(s)+2 \mathrm{CO}_{2}(g)\) c. \(3 \mathrm{NO}_{2}(g)+\mathrm{H}_{2} \mathrm{O}(\ell) \rightarrow 2 \mathrm{NO}_{3}^{-}(a q)+\mathrm{NO}(g)+2 \mathrm{H}^{+}(a q)\)

How many moles of \(\mathrm{O}_{2}\) are consumed in the conversion of one mole of \(\mathrm{FeCO}_{3}\) to each of the following compounds? Assume \(\mathrm{CO}_{2}\) is also produced. (a) \(\mathrm{Fe}_{2} \mathrm{O}_{3} ;\) (b) \(\mathrm{Fe}_{3} \mathrm{O}_{4}\)

Uranium is found in Earth's crust as UO, and in an assortment of compounds containing UO \(_{2}^{n+}\) cations. How many moles of electrons are transferred in the conversion of one mole of \(\mathrm{UO}_{2}\) to each of the following species? In which of the conversions is uranium oxidized? (a) \(\mathrm{UO}_{2}\left(\mathrm{CO}_{3}\right)_{3}^{4-}(a q) ;\) (b) \(\mathrm{UO}_{2}\left(\mathrm{HPO}_{4}\right)_{2}^{2-}(a q)\)

A method for determining the quantity of dissolved oxygen in natural waters requires a series of redox reactions. Balance the following chemical equations in that series under the conditions indicated: a. \(\mathrm{Mn}^{2+}(a q)+\mathrm{O}_{2}(g) \rightarrow \mathrm{MnO}_{2}(s)\) (basic solution) b. \(\mathrm{MnO}_{2}(s)+\mathrm{I}^{-}(a q) \rightarrow \mathrm{Mn}^{2+}(a q)+\mathrm{I}_{2}(s)\) (acidic solution) c. \(\mathrm{I}_{2}(s)+\mathrm{S}_{2} \mathrm{O}_{3}^{2-}(a q) \rightarrow\) \(\mathrm{I}^{-}(a q)+\mathrm{S}_{4} \mathrm{O}_{6}^{2-}(a q) \quad\) (neutral solution)

Silver can be extracted from rocks using cyanide ion. Complete and balance the following reaction for this process: \(\mathrm{Ag}(s)+\mathrm{CN}^{-}(a q)+\mathrm{O}_{2}(g) \rightarrow\) \(\mathrm{Ag}(\mathrm{CN})_{2}-(a q) \quad\) (basic solution)

Permanganate ion \(\left(\mathrm{MnO}_{4}^{-}\right)\) is used in water purification to remove oxidizable substances. Complete and balance the following reactions for the removal of sulfide, cyanide, and sulfite. Assume that reaction conditions are basic: a. \(\mathrm{MnO}_{4}^{-}(a q)+\mathrm{S}^{2-}(a q) \rightarrow \mathrm{MnS}(s)+\mathrm{S}_{8}(s)\) b. \(\mathrm{MnO}_{4}^{-}(a q)+\mathrm{CN}^{-}(a q) \rightarrow \mathrm{MnO}_{2}(s)+\mathrm{CNO}^{-}(a q)\) c. \(\operatorname{Mn} \mathrm{O}_{4}^{-}(a q)+\mathrm{SO}_{3}^{2-}(a q) \rightarrow \mathrm{MnO}_{2}(s)+\mathrm{SO}_{4}^{2-}(a q)\)

The water-soluble gas \(\mathrm{ClO}_{2}\) is known as an oxidative biocide. It destroys bacteria by oxidizing their cell walls and viruses by attacking their viral envelopes. \(\mathrm{ClO}_{2}\) may be prepared for use as a decontaminating agent from several different starting materials in slightly acidic solutions. Complete and balance the following chemical reactions for the synthesis of \(\mathrm{ClO}_{2}\) a. \(\mathrm{ClO}_{3}^{-}(a q)+\mathrm{SO}_{2}(g) \rightarrow \mathrm{ClO}_{2}(g)+\mathrm{SO}_{4}^{2-}(a q)\) b. \(\mathrm{ClO}_{3}^{-}(a q)+\mathrm{Cl}^{-}(a q) \rightarrow \mathrm{ClO}_{2}(g)+\mathrm{Cl}_{2}(g)\) c. \(\mathrm{ClO}_{3}^{-}(a q)+\mathrm{Cl}_{2}(g) \rightarrow \mathrm{ClO}_{2}(g)+\mathrm{O}_{2}(g)\)

How many milliliters of \(0.100 M\) NaOH are required to neutralize the following solutions? a. \(10.0 \mathrm{mL}\) of \(0.0500 M \mathrm{HCl}\) b. \(25.0 \mathrm{mL}\) of \(0.126 M \mathrm{HNO}_{3}\) c. \(50.0 \mathrm{mL}\) of \(0.215 M \mathrm{H}_{2} \mathrm{SO}_{4}\)

How many milliliters of \(0.100 M\) HNO \(_{3}\) are needed to neutralize the following solutions? a. \(45.0 \mathrm{mL}\) of \(0.667 M \mathrm{KOH}\) b. \(58.5 \mathrm{mL}\) of \(0.0100 M \mathrm{Al}(\mathrm{OH})_{3}\) c. \(34.7 \mathrm{mL}\) of \(0.775 M \mathrm{NaOH}\)

The solubility of slaked lime, \(\mathrm{Ca}(\mathrm{OH})_{2},\) in water at \(20^{\circ} \mathrm{C}\) is \(0.185 \mathrm{g} / 100.0 \mathrm{mL} .\) What volume of \(0.00100 M \mathrm{HCl}\) is needed to neutralize \(10.0 \mathrm{mL}\) of a saturated \(\mathrm{Ca}(\mathrm{OH})_{2}\) solution?

The solubility of magnesium hydroxide, \(\mathrm{Mg}(\mathrm{OH})_{2},\) in water is \(9.0 \times 10^{-4} \mathrm{g} / 100.0 \mathrm{mL} .\) What volume of \(0.00100 M \mathrm{HNO}_{3}\) is required to neutralize \(1.00 \mathrm{L}\) of a saturated Mg(OH), solution?

Explain how a mixture of anion and cation exchangers can be used to deionize water.

Describe the process by which the ion exchanger in a home water softener is regenerated for further use.

What would be the advantage in using a \(\mathrm{K}^{+}\) resin rather than a \(\mathrm{Na}^{+}\) resin to soften water?

To determine the concentration of \(\mathrm{SO}_{4}^{2-}\) ion in a sample of groundwater, \(100.0 \mathrm{mL}\) of the sample is titrated with \(0.0250 M \mathrm{Ba}\left(\mathrm{NO}_{3}\right)_{2},\) forming insoluble \(\mathrm{BaSO}_{4} .\) If \(3.19 \mathrm{mL}\) of the \(\mathrm{Ba}\left(\mathrm{NO}_{3}\right)_{2}\) solution is required to reach the end point of the titration, what is the molarity of the \(\mathrm{SO}_{4}^{2-} ?\)

Ethylene glycol is the common name for the liquid used to keep the coolant in automobile cooling systems from freezing. It is \(38.7 \%\) carbon, \(9.7 \%\) hydrogen, and \(51.6 \%\) oxygen by mass. Its molar mass is \(62.07 \mathrm{g} / \mathrm{mol}\) and its density is \(1.106 \mathrm{g} / \mathrm{mL}\) at \(20^{\circ} \mathrm{C}\) a. What is the empirical formula of ethylene glycol? b. What is the molecular formula of ethylene glycol? c. In a solution prepared by mixing equal volumes of water and ethylene glycol, which ingredient is the solute and which is the solvent?

According to the label on a bottle of concentrated hydrochloric acid, the contents are \(36.0 \%\) HCl by mass and have a density of \(1.18 \mathrm{g} / \mathrm{mL}\) a. What is the molarity of this concentrated HCl? b. What volume of it would you need to prepare \(0.250 \mathrm{L}\) of \(2.00 \mathrm{MHCl} ?\) c. What mass of sodium hydrogen carbonate would be needed to neutralize the spill if a bottle containing 1.75 L of this concentrated HCl dropped on a lab floor and broke open?

Chlorine was first prepared in 1774 by heating a mixture of \(\mathrm{NaCl}\) and \(\mathrm{MnO}_{2}\) in sulfuric acid: $$\begin{aligned} \mathrm{NaCl}(a q)+\mathrm{H}_{2} \mathrm{SO}_{4}(a q)+\mathrm{MnO}_{2}(s) & \rightarrow \\ & \mathrm{Na}_{2} \mathrm{SO}_{4}(a q)+\mathrm{MnCl}_{2}(a q)+\mathrm{H}_{2} \mathrm{O}(\ell)+\mathrm{Cl}_{2}(g) \end{aligned}$$a. Assign oxidation numbers to the elements in each compound, and balance the redox reaction in acid solution. b. Write a net ionic equation describing the reaction for the formation of chlorine. c. If chlorine gas is inhaled, it causes pulmonary edema (fluid in the lungs) because it reacts with water in the alveolar sacs of the lungs to produce the strong acid \(\mathrm{HCl}\) and the weaker acid HOC1. Balance the equation for the conversion of \(\mathrm{Cl}_{2}\) to \(\mathrm{HCl}\) and \(\mathrm{HOCl}\).

When a solution of dithionite ions \(\left(\mathrm{S}_{2} \mathrm{O}_{4}^{2-}\right)\) is added to a solution of chromate ions \(\left(\mathrm{CrO}_{4}^{2-}\right),\) the products of the ensuing chemical reaction that occurs under basic conditions include soluble sulfite ions and solid chromium(III) hydroxide. This reaction is used to remove \(\mathrm{Cr}^{6+}\) from wastewater generated by factories that make chrome- plated metals. a. Write the net ionic equation for this redox reaction. b. Which element is oxidized and which is reduced? c. Identify the oxidizing and reducing agents in the reaction. d. How many grams of sodium dithionite would be needed to remove the \(\mathrm{Cr}^{6+}\) in \(100.0 \mathrm{L}\) of wastewater that contains \(0.00148 M\) chromate ion?

A prototype battery based on iron compounds with large, positive oxidation numbers was developed in \(1999 .\) In the following reactions, assign oxidation numbers to the clements in each compound and balance the redox reactions in basic solution: a. \(\mathrm{FeO}_{4}^{2-}(a q)+\mathrm{H}_{2} \mathrm{O}(\ell) \rightarrow \mathrm{FeOOH}(s)+\mathrm{O}_{2}(g)+\mathrm{OH}^{-}(a q)\) b. \(\mathrm{FeO}_{4}^{2-}(a q)+\mathrm{H}_{2} \mathrm{O}(\ell) \rightarrow \mathrm{Fe}_{2} \mathrm{O}_{3}(s)+\mathrm{O}_{2}(g)+\mathrm{OH}^{-}(a q)\)

Silver tarnish is the result of silver metal reacting with sulfur compounds, such as \(\mathrm{H}_{2} \mathrm{S},\) and \(\mathrm{O}_{2}\) in the air. The tarnish on silverware \(\left(\mathrm{Ag}_{2} \mathrm{S}\right)\) can be removed by soaking the silverware in a slightly basic solution of \(\mathrm{NaHCO}_{3}\) (baking soda) in a basin lined with aluminum foil. a. Write a balanced chemical equation for the tarnish formation reaction. b. Write a balanced net ionic equation for the tarnish removal process in which Ag_S S reacts with A1 metal, forming \(\mathrm{Al}(\mathrm{OH})_{3}(s), \mathrm{Ag}\) metal, and \(\mathrm{HS}^{-}\) ions.

Give the formulas of the acids formed in the following chemical reactions of chlorine oxides. a. \(\mathrm{ClO}+\mathrm{H}_{2} \mathrm{O} \rightarrow ?+?\) b. \(\mathrm{Cl}_{2} \mathrm{O}+\mathrm{H}_{2} \mathrm{O} \rightarrow \mathrm{HCl}+?\) c. \(\mathrm{Cl}_{2} \mathrm{O}_{6}+\mathrm{H}_{2} \mathrm{O} \rightarrow ?+?\)

Many nonmetal oxides react with water to form acidic solutions. Give the formulas of the acids produced in the following reactions: a. \(P_{4} O_{10}+6 H_{2} O \rightarrow ?\) b. \(\mathrm{SeO}_{2}+\mathrm{H}_{2} \mathrm{O} \rightarrow ?\) c. \(\mathrm{B}_{2} \mathrm{O}_{3}+3 \mathrm{H}_{2} \mathrm{O} \rightarrow ?\)

Write net ionic equations for the reactions that occur when a. a sample of acetic acid is titrated with a solution of \(\mathrm{KOH}\) b. a solution of sodium carbonate is mixed with a solution of calcium chloride. c. calcium oxide dissolves in water.

Sodium fluoride is added to drinking water in many municipalities to protect teeth against cavities. The target of the fluoridation is hydroxyapatite, \(\mathrm{Ca}_{10}\left(\mathrm{PO}_{4}\right)_{6}(\mathrm{OH})_{2},\) a compound in tooth enamel. There is concern, however, that fluoride ions in water may contribute to skeletal fluorosis, an arthritis-like disease. a. Write a net ionic equation for the reaction between hydroxyapatite and sodium fluoride that produces fluorapatite, \(\mathrm{Ca}_{10}\left(\mathrm{PO}_{4}\right)_{6} \mathrm{F}_{2}\) b. The U.S. EPA currently restricts the concentration of \(\mathrm{F}^{-}\) in drinking water to \(4 \mathrm{mg} / \mathrm{L}\). Express this concentration of \(F^{-}\) in molarity. c. One study of skeletal fluorosis suggests that drinking water with a fluoride concentration of \(4 \mathrm{mg} / \mathrm{L}\) for 20 years raises the fluoride content in bone to \(6 \mathrm{mg} / \mathrm{g}\), a level at which a patient may experience stiff joints and other symptoms. How much fluoride (in milligrams) is present in a 100 mg sample of bone with this fluoride concentration?

Near Las Vegas, NV, improper disposal of perchlorates used to manufacture rocket fuel contaminated a stream flowing into Lake Mead, the largest artificial lake in the United States and a major supply of drinking and irrigation water for the American Southwest. The U.S. EPA has proposed an advisory range for perchlorate concentrations in drinking water of 4 to \(18 \mu \mathrm{g} / \mathrm{L} .\) The perchlorate concentration in the stream averages \(700.0 \mu \mathrm{g} / \mathrm{L},\) and the stream flows at an average rate of 161 million gallons per day \((1 \text { gal }=3.785\) L) a. What are the formulas of sodium perchlorate and ammonium perchlorate? b. How many kilograms of perchlorate flow from the Las Vegas stream into Lake Mead each day? c. What volume of perchlorate-free lake water would have to mix with the stream water each day to dilute the stream's perchlorate concentration from 700.0 to \(4 \mu \mathrm{g} / \mathrm{L} ?\) d. since \(2003,\) the states of Maryland, Massachusetts, and New Mexico have limited perchlorate concentrations in drinking water to \(0.1 \mu \mathrm{g} / \mathrm{L} .\) Five replicate samples were analyzed for perchlorates by laboratories in each state, and the following data ( \(\mu \mathrm{g} / \mathrm{L}\) ) were collected: $$\begin{array}{ccc} \mathrm{MD} & \mathrm{MA} & \mathrm{NM} \\ 1.1 & 0.90 & 1.2 \\ \hline 1.1 & 0.95 & 1.2 \\ \hline 1.4 & 0.92 & 1.3 \\ \hline 1.3 & 0.90 & 1.4 \\ \hline 0.9 & 0.93 & 1.1 \\ \hline \end{array}$$Which of the labs produced the most precise analytical results?

Some people who prefer natural foods make their own apple cider vinegar. They start with freshly squeezed apple juice that contains about \(6 \%\) natural sugars. These sugars, which all have nearly the same empirical formula, \(\mathrm{CH}_{2} \mathrm{O},\) are fermented with yeast in a chemical reaction that produces equal numbers of moles of ethanol \(\left(\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{OH}\right)\) and carbon dioxide. The product of fermentation, called hard cider, undergoes an acid fermentation step in which ethanol and dissolved oxygen gas react together to form acetic acid (CH \(_{3} \mathrm{COOH}\) ) and water. This acetic acid is the principal solute in vinegar. a. Write a balanced chemical equation for the fermentation of natural sugars to ethanol and carbon dioxide. You may use in the equation the empirical formula given in the preceding paragraph. b. Write a balanced chemical equation for the acid fermentation of ethanol to acetic acid. c. What are the oxidation states of carbon in the reactants and products of the two fermentation reactions? d. If a sample of apple juice contains \(1.00 \times 10^{2} \mathrm{g}\) of natural sugar, what is the maximum quantity of acetic acid that could be produced by the two fermentation reactions?

The stalactites and stalagmites in most caves are made of calcium carbonate (see Figure 8.10 ). In the Lower Kane Cave in Wyoming, however, they are made of gypsum (calcium sulfate). The presence of \(\mathrm{CaSO}_{4}\) is explained by the following sequence of reactions: $$ \begin{array}{c} \mathrm{H}_{2} \mathrm{S}(a q)+2 \mathrm{O}_{2}(g) \rightarrow \mathrm{H}_{2} \mathrm{SO}_{4}(a q) \\ \mathrm{H}_{2} \mathrm{SO}_{4}(a q)+\mathrm{CaCO}_{3}(s) \rightarrow \mathrm{CaSO}_{4}(s)+\mathrm{H}_{2} \mathrm{O}(\ell)+\mathrm{CO}_{2}(g) \end{array} $$ a. Which (if either) of these reactions is a redox reaction? b. Write a net ionic equation for the reaction of \(\mathrm{H}_{2} \mathrm{SO}_{4}\) with \(\mathrm{CaCO}_{3}\) c. How would the net ionic equation be different if the reaction were written as follows? $$ \mathrm{H}_{2} \mathrm{SO}_{4}(a q)+\mathrm{CaCO}_{3}(s) \rightarrow \mathrm{CaSO}_{4}(s)+\mathrm{H}_{2} \mathrm{CO}_{3}(a q) $$

Which of the following reactions of calcium compounds is or are redox reactions? a. \(\mathrm{CaCO}_{3}(s) \rightarrow \mathrm{CaO}(s)+\mathrm{CO}_{2}(g)\) b. \(\mathrm{CaO}(s)+\mathrm{SO}_{2}(g) \rightarrow \mathrm{CaSO}_{3}(s)\) c. \(\mathrm{CaCl}_{2}(s) \rightarrow \mathrm{Ca}(s)+\mathrm{Cl}_{2}(g)\) d. \(3 \mathrm{Ca}(s)+\mathrm{N}_{2}(g) \rightarrow \mathrm{Ca}_{3} \mathrm{N}_{2}(s)\)

HF is prepared by reacting \(\mathrm{CaF}_{2}\) with \(\mathrm{H}_{2} \mathrm{SO}_{4}:\) $$ \mathrm{CaF}_{2}(s)+\mathrm{H}_{2} \mathrm{SO}_{4}(\ell) \rightarrow 2 \mathrm{HF}(g)+\mathrm{CaSO}_{4}(s) $$ HF can be electrolyzed, in turn, when dissolved in molten KF to produce fluorine gas: $$ 2 \mathrm{HF}(\ell) \rightarrow \mathrm{F}_{2}(g)+\mathrm{H}_{2}(g) $$ Fluorine is extremely reactive, so it is typically sold as a \(5 \%\) mixture by volume in an inert gas such as helium. How much \(\mathrm{CaF}_{2}\) is required to produce \(500.0 \mathrm{L}\) of \(5 \% \mathrm{F}_{2}\) in helium? Assume the density of \(\mathrm{F}_{2}\) gas is \(1.70 \mathrm{g} / \mathrm{L}\)