Solution: As stated in the chapter, potassium superoxide

Define mineral, ore, and metallurgy.

List three metals that are usually found in an uncombined state in nature and three metals that are always found in a combined state in nature.

Write chemical formulas for the following minerals: (a) calcite, (b) dolomite, (c) fluorite, (d) halite, (e) corundum, (f) magnetite, (g) beryl, (h) galena, (i) epsomite, (j) anhydrite.

Name the following minerals: (a) MgCO3, (b) Na3AlF6, (c) Al2O3, (d) Ag2S, (e) HgS, (f) ZnS, (g) SrSO4, (h) PbCO3, (i) MnO2, (j) TiO2.

Describe the main steps involved in the preparation of an ore

What does roasting mean in metallurgy? Why is roasting a major source of air pollution and acid rain?

Describe with examples the chemical and electrolytic reduction processes used in the production of metals

Describe the main steps used to purify metals.

Describe the extraction of iron in a blast furnace.

Briefly discuss the steelmaking process.

In the Mond process for the purification of nickel, CO is passed over metallic nickel to give Ni(CO)4: Ni(s) 1 4CO(g) Ni(CO)4(g) Given that the standard free energies of formation of CO(g) and Ni(CO)4(g) are 2137.3 kJ/mol and 2587.4 kJ/mol, respectively, calculate the equilibrium constant of the reaction at 80C. (Assume DGf to be independent of temperature.)

Copper is purified by electrolysis (see Figure 21.6). A 5.00-kg anode is used in a cell where the current is 37.8 A. How long (in hours) must the current run to dis solve this anode and electroplate it onto the cathode?

Consider the electrolytic procedure for purifying copper described in Figure 21.6. Suppose that a sample of copper contains the following impurities: Fe, Ag, Zn, Au, Co, Pt, and Pb. Which of the metals will be oxidized and dissolved in solution and which will be unaffected and simply form the sludge that accumulates at the bottom of the cell?

How would you obtain zinc from sphalerite (ZnS)?

Starting with rutile (TiO2), explain how you would obtain pure titanium metal. (Hint: First convert TiO2 to TiCl4. Next, reduce TiCl4 with Mg. Look up physical properties of TiCl4, Mg, and MgCl2 in a chemistry handbook.)

A certain mine produces 2.0 3 108 kg of copper from chalcopyrite (CuFeS2) each year. The ore contains only 0.80 percent Cu by mass. (a) If the density of the ore is 2.8 g/cm3 , calculate the volume (in cm3 ) of ore removed each year. (b) Calculate the mass (in kg) of SO2 produced by roasting (assume chalcopyrite to be the only source of sulfur).

Which of the following compounds would require electrolysis to yield the free metals? Ag2S, CaCl2, NaCl, Fe2O3, Al2O3, TiCl4.

Although iron is only about two-thirds as abundant as aluminum in Earths crust, mass for mass it costs only about one-quarter as much to produce. Why?

Define the following terms: conductor, insulator, semiconducting elements, donor impurities, acceptor impurities, n-type semiconductors, p-type semiconductors.

Briefly discuss the nature of bonding in metals, insulators, and semiconducting elements.

Describe the general characteristics of n-type and p-type semiconductors.

State whether silicon would form n-type or p-type semiconductors with the following elements: Ga, Sb, Al, As.

How is sodium prepared commercially?

Why is potassium usually not prepared electrolytically from one of its salts?

Describe the uses of the following compounds: NaCl, Na2CO3, NaOH, KOH, KO2.

Under what conditions do sodium and potassium form Na2 and K2 ions?

Complete and balance the following equations: (a) K(s) 1 H2O(l) (b) NaH(s) 1 H2O(l) (c) Na(s) 1 O2(g) (d) K(s) 1 O2(g)

Write a balanced equation for each of the following reactions: (a) sodium reacts with water; (b) an aqueous solution of NaOH reacts with CO2; (c) solid Na2CO3 reacts with a HCl solution; (d) solid NaHCO3 reacts with a HCl solution; (e) solid NaHCO3 is heated; (f ) solid Na2CO3 is heated.

Sodium hydride (NaH) can be used as a drying agent for many organic solvents. Explain how it works.

Calculate the volume of CO2 at 10.0C and 746 mmHg pressure obtained by treating 25.0 g of Na2CO3 with an excess of hydrochloric acid.

List the common ores of magnesium and calcium.

How are the metals magnesium and calcium obtained commercially?

From the thermodynamic data in Appendix 3, calculate the DH values for the following decompositions: (a) MgCO3(s) MgO(s) 1 CO2(g) (b) CaCO3(s) CaO(s) 1 CO2(g) Which of the two compounds is more easily decomposed by heat?

Starting with magnesium and concentrated nitric acid, describe how you would prepare magnesium oxide. [Hint: First convert Mg to Mg(NO3)2. Next, MgO can be obtained by heating Mg(NO3)2.]

Describe two ways of preparing magnesium chloride.

The second ionization energy of magnesium is only about twice as great as the first, but the third ionization energy is 10 times as great. Why does it take so much more energy to remove the third electron?

List the sulfates of the Group 2A metals in order of increasing solubility in water. Explain the trend. (Hint: You need to consult a chemistry handbook.)

Helium contains the same number of electrons in its outer shell as do the alkaline earth metals. Explain why helium is inert whereas the Group 2A metals are not.

When exposed to air, calcium first forms calcium oxide, which is then converted to calcium hydroxide, and finally to calcium carbonate. Write a balanced equation for each step.

Write chemical formulas for (a) quicklime, (b) slaked lime, (c) limewater.

Describe the Hall process for preparing aluminum.

What action renders aluminum inert?

Before Hall invented his electrolytic process, aluminum was produced by the reduction of its chloride with an active metal. Which metals would you use for the production of aluminum in that way?

With the Hall process, how many hours will it take to deposit 664 g of Al at a current of 32.6 A?

Aluminum forms the complex ions AlCl4 2 and AlF6 32. Describe the shapes of these ions. AlCl6 32 does not form. Why? (Hint: Consider the relative sizes of Al31, F2, and Cl2 ions.)

The overall reaction for the electrolytic production of aluminum by means of the Hall process may be represented as Al2O3 (s) 1 3C(s) 2Al(l) 1 3CO(g) At 1000C, the standard free-energy change for this process is 594 kJ/mol. (a) Calculate the minimum voltage required to produce 1 mole of aluminum at this temperature. (b) If the actual voltage applied is exactly three times the ideal value, calculate the energy required to produce 1.00 kg of the metal.

In basic solution, aluminum metal is a strong reducing agent and is oxidized to AlO2 2. Give balanced equations for the reaction of Al in basic solution with the following: (a) NaNO3, to give ammonia; (b) water, to give hydrogen; (c) Na2SnO3, to give metallic tin.

Write a balanced equation for the thermal decomposition of aluminum nitrate to form aluminum oxide, nitrogen dioxide, and oxygen gas

Describe some of the properties of aluminum that make it one of the most versatile metals known.

The pressure of gaseous Al2Cl6 increases more rapidly with temperature than predicted by the ideal gas equation even though Al2Cl6 behaves like an ideal gas. Explain.

Starting with aluminum, describe with balanced equations how you would prepare (a) Al2Cl6, (b) Al2O3, (c) Al2(SO4)3, (d) NH4Al(SO4)2 ? 12H2O

Explain the change in bonding when Al2Cl6 dissociates to form AlCl3 in the gas phase.

In steelmaking, nonmetallic impurities such as P, S, and Si are removed as the corresponding oxides. The inside of the furnace is usually lined with CaCO3 and MgCO3, which decompose at high temperatures to yield CaO and MgO. How do CaO and MgO help in the removal of the nonmetallic oxides?

When 1.164 g of a certain metal sulfide was roasted in air, 0.972 g of the metal oxide was formed. If the oxidation number of the metal is 12, calculate the molar mass of the metal.

An early view of metallic bonding assumed that bonding in metals consisted of localized, shared electron-pair bonds between metal atoms. What evidence would help you to argue against this viewpoint?

Referring to Figure 21.6, would you expect H2O and H1 to be reduced at the cathode and H2O oxidized at the anode?

A 0.450-g sample of steel contains manganese as an impurity. The sample is dissolved in acidic solution and the manganese is oxidized to the permanganate ion MnO4 2. The MnO4 2 ion is reduced to Mn21 by reacting with 50.0 mL of 0.0800 M FeSO4 solution. The excess Fe21 ions are then oxidized to Fe31 by 22.4 mL of 0.0100 M K2Cr2O7. Calculate the percent by mass of manganese in the sample

Given that DGf(Fe2O3) 5 2741.0 kJ/mol and that DGf(Al2O3) 5 21576.4 kJ/mol, calculate DG for the following reactions at 25C: (a) 2Fe2O3(s) 4Fe(s) 1 3O2(g) (b) 2Al2O3(s) 4Al(s) 1 3O2(g)

Use compounds of aluminum as an example to explain what is meant by amphoterism.

When an inert atmosphere is needed for a metallurgical process, nitrogen is frequently used. However, in the reduction of TiCl4 by magnesium, helium is used. Explain why nitrogen is not suitable for this process.

It has been shown that Na2 species form in the vapor phase. Describe the formation of the disodium mole cule in terms of a molecular orbital energylevel diagram. Would you expect the alkaline earth metals to exhibit a similar property?

Explain each of the following statements: (a) An aqueous solution of AlCl3 is acidic. (b) Al(OH)3 is soluble in NaOH solutions but not in NH3 solution.

Write balanced equations for the following reactions: (a) the heating of aluminum carbonate; (b) the reaction between AlCl3 and K; (c) the reaction between solutions of Na2CO3 and Ca(OH)2.

Write a balanced equation for the reaction between calcium oxide and dilute HCl solution.

What is wrong with the following procedure for obtaining magnesium? MgCO3 MgO(s) 1 CO2(g) MgO(s) 1 CO(g) Mg(s) 1 CO2(g)

Explain why most metals have a flickering appearance

Predict the chemical properties of francium, the last member of Group 1A.

Describe a medicinal or health-related application for each of the following compounds: NaF, Li2CO3, Mg(OH)2, CaCO3, BaSO4, Al(OH)2NaCO3. (You would need to do a Web search for some of these compounds.)

The following are two reaction schemes involving magnesium. Scheme I: When magnesium burns in oxygen, a white solid (A) is formed. A dissolves in 1 M HCl to give a colorless solution (B). Upon addition of Na2CO3 to B, a white precipitate is formed (C). On heating, C decomposes to D and a colorless gas is generated (E). When E is passed through limewater [an aqueous suspension of Ca(OH)2], a white precipitate appears (F). Scheme II: Magnesium reacts with 1 M H2SO4 to produce a colorless solution (G). Treating G with an excess of NaOH produces a white precipitate (H). H dissolves in 1 M HNO3 to form a colorless solution. When the solution is slowly evaporated, a white solid (I) appears. On heating I, a brown gas is given off. Identify AI and write equations representing the reactions involved.

Lithium and magnesium exhibit a diagonal relationship in some chemical properties. How does lithium resemble magnesium in its reaction with oxygen and nitrogen? Consult a handbook of chemistry and compare the solubilities of carbonates, fluorides, and phosphates of these metals.

To prevent the formation of oxides, peroxides, and superoxides, alkali metals are sometimes stored in an inert atmosphere. Which of the following gases should not be used for lithium? Why? Ne, Ar, N2, Kr.

Which of the following metals is not found in the free state in nature: Ag, Cu, Zn, Au, Pt?

After heating, a metal surface (such as that of a cooking pan or skillet) develops a color pattern like an oil slick on water. Explain.

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 H2 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 HNO3 but D does not. Arrange the metals in the increasing order as reducing agents. Suggest four metals that fit these descriptions.

The electrical conductance of copper metal decreases with temperature, but that of a CuSO4 solution increases with temperature. Explain.

As stated in the chapter, potassium superoxide (KO2) is a useful source of oxygen employed in breathing equipment. Calculate the pressure at which oxygen gas stored at 20C would have the same density as the oxygen gas provided by KO2. The density of KO2 at 20C is 2.15 g/cm3 .

A sample of 10.00 g of sodium reacts with oxygen to form 13.83 g of sodium oxide (Na2O) and sodium peroxide (Na2O2). Calculate the percent composition of the mixture.