SGCh20 - Chapter Twenty METALLURGY AND THE CHEMISTRY OF...

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Unformatted text preview: Chapter Twenty METALLURGY AND THE CHEMISTRY OF METALS • • • • • Metallurgical Processes The Band Theory of Conductivity The Alkali Metals The Alkaline Earth Metals Aluminum METALLURGICAL PROCESSES STUDY OBJECTIVES 1. 2. 3. List the procedures used to separate minerals from gangue. Describe briefly the means used to produce and purify a metal. Describe the steps in the production of iron. Preliminary Treatment. Metallurgy is the science and technology of extracting metals from their ores and preparing these metals for use. Most metals are found in nature in the combined state as minerals. A mineral is a naturally occurring substance with a characteristic range of composition. An ore is a mineral deposit of high enough concentration to make an economical recoverery of a desired metal. Ores generally must be cleaned before the metal is extracted. The waste materials, called gangue, are removed by flotation, ferromagnetic separation, or distillation of mercury amalgams. Reduction of Metals. The process of producing a free metal is always one of reduction. Reduction of metal oxides is usually less complex than reduction of sulfides. Roasting is a preliminary operation used to convert a metal sulfide or carbonate into an oxide. Two types of reductions are ordinarily employed: 1. 2. Chemical reduction is the use of a reducing agent to prepare the elemental form of a metal from a compound. Calcium, magnesium, aluminum, hydrogen, and carbon are often used in chemical reductions. Electrolytic reduction is necessary to produce the very electropositive metals such as sodium, magnesium and aluminum. Iron. Iron is prepared by a chemical reduction process. The raw materials for making iron are (1) iron ore, either hematite Fe2 O3 or magnetite Fe 3 O4 , (2) coke, and (3) limestone. These three materials are mixed and fed into a huge blast furnace (Figure 20.3 text). A strong blast of air preheated to 1500°C is blown in at the bottom, where oxygen in the air reacts with the coke. This reaction supplies heat for the furnace, and carbon monoxide which is the reducing agent in the reaction of the iron oxides. 2C(g) + O2 (g) → 2CO(g) Fe2 O3 (s) + 3CO(g) → 2Fe(l) + 3CO2 (g) 4 03 Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website 4 04 / Metallurgy and the Chemistry of Metals A temperature gradient is set up in the furnace. The temperature at the bottom is 1500°C, and it decreases with height; falling to about 250°C at the top. Much of the carbon dioxide produced is reduced by reaction with excess coke as follows: CO2 (g) + C(s) → 2CO(g) thereby forming more of the principal reducing agent. The gas that escapes at the top of the furnace is mostly nitrogen and carbon monoxide. The molten iron runs to the bottom where it is withdrawn periodically. Limestone serves as a flux in the removal of impurities such as silica (SiO2 ) and alumina (Al2 O3 ). The limestone decomposes at temperatures above 900°C, forming calcium oxide and carbon dioxide. CaCO3 (s) → CaO(s) + CO2 (g) Calcium oxide unites with silica and alumina to form a glassy, molten substance called slag which is composed mainly of CaSiO3 and some calcium aluminate. Slag is less dense than iron, and so it collects as a pool on top of the metal. It is drawn off, leaving the molten iron. Iron prepared in this manner contains many impurities and is called pig iron, or cast iron. Cast iron is made into steel by further treatment in a basic oxygen furnace. Purification of Metals. Once a metal is prepared it may need to be further purified. Distillation, electrolysis, and zone refining are three common procedures. Metals with low boiling points, such as mercury (357°C), cadmium (767°C), and zinc (907°C), are referred to as volatile. They can be purified by fractional distillation. On heating, these metals vaporize, leaving behind any nonvolatile impurities. Condensation of the vapor yields the purified metal. Metals can also be purified by electrolysis. The metals that plate out on a cathode can be controlled by the voltage at which the electrolysis is carried out. In the purification of copper, Cu2+ is reduced much more readily than iron or zinc ions, which are common impurities. At just the right voltage Cu is plated out and the metal impurities remain in solution. Zone refining takes advantage of the fact that when liquids begin to freeze, the impurities tend to remain in the liquid phase. See Figure 20.8 of the textbook. Extremely pure metals can be obtained by repeating this process a number of times. _______________________________________________________________________________ EXAMPLE 20.1 Chemical Reduction Look up the oxidation potentials of calcium, magnesium, aluminum, and hydrogen and suggest five metals listed in Table 19.1 that calcium, magnesium, aluminum can reduce, but that hydrogen cannot reduce. •Method of Solution Ca, Mg, and Al are strong reducing agents. They have high oxidation potentials which means they have a strong tendency to lose electrons and thereby reduce other metal cations to the pure metal.. o Ca → Ca 2+ + 2e – Mg → Mg 2+ + 2e – Al → Al 3+ + 3e – Eox = 2.87 o Eox = 2.37 V o Eox = 1.66 V H2 is not as strong a reducing agent as the others. o H2 → 2H+ + 2e – Eox = 0 V Ca, Mg, and Al will reduce ions of most of the active metals such as Zn2+ , Cr 3+ , Co 2+ , Ni 2+ , and Sn 2+ . H 2 cannot reduce these metals, but is used to reduce more noble metals such as copper. _______________________________________________________________________________ Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website Metallurgy and the Chemistry of Metals / 4 05 _______________________________________________________________________________ EXAMPLE 20.2 Steel Production How are the impurities present in iron ore, such as silica and sulfur, removed during steel production? •Method of Solution Acidic impurities such as silica and sulfides combine with CaO, a base, to form a molten slag. SiO 2 + CaO → CaSiO 3 FeS + CaO → CaS + FeO The slag has a lower density than molten iron. Therefore it collects as a pool on top of the metal. This slag is drawn off and used as a component in making cement. _______________________________________________________________________________ EXERCISES 1. 2. 3. 4. 5. Distinguish between an ore and a mineral. Why must electrolytic reduction be used in the production of some metals such as lithium and sodium? What is zone refining used for? How is limestone used in iron production? Write an equation for the chemical reduction of Cr2 O3 with Al metal yielding chromium. THE BAND THEORY OF CONDUCTIVITY STUDY OBJECTIVES 1. 2. Compare the energies of the conduction bands and valence bands in metals, insulators, and semiconductors. Use the band theory to describe conduction in metals. The Conduction Band. The band theory is the result of the application of molecular orbital theory to metals. In metals, the atoms lie in a three-dimensional array and take part in bonding that spreads over the entire crystal. Take, for example, an alkali metal atom that carries a single valence electron in an s orbital. Recall that the atomic orbitals of two atoms will overlap when the atoms are close together. This results in the formation of two molecular orbitals: a bonding orbital and an antibonding orbital. The total number of molecular orbitals produced always equals the number of atomic orbitals that overlap to produce them. The overlap of four atomic orbitals from four atoms, for instance, will form four molecular orbitals: two bonding and two antibonding orbitals. A crystal made up of N atoms with overlapping atomic orbitals will have a total of N molecular orbitals. These bonding and antibonding orbitals are so closely spaced in terms of energy that they are called a "band." The formation of a band is shown in Figure 20.1. The band containing the valence shell electrons (3s for Na) is called the v alence band . For alkali metals there are N valence electrons and N orbitals in the band. These N orbitals have a capacity for 2N electrons. Any band that is either vacant or partially filled is called a c onduction band . Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website 4 06 / Metallurgy and the Chemistry of Metals Antibonding orbital E a. Bonding orbital b. c. Figure 20.1 The formation of an energy band by the successive overlap of atomic orbitals. a. When the orbitals of two atoms overlap, one bonding orbital and one antibonding orbital of significantly different energies are formed. b. When the orbitals of four atoms overlap, four molecular orbitals are formed. c. When the orbitals of N atoms overlap where N is a large number, N orbitals are formed. With so many orbitals, the orbital energies differ only infinitesimally from each other and form a virtually continuous band. Electrical Conduction. The band theory explains conduction in the alkali metals in the following way. When a voltage is applied across a piece of sodium metal, conduction occurs. The current is the result of electrons in the 3s band being free to jump from atom to atom. The free movement of electrons is possible for two reasons. In alkali metals the conduction band and the valence band are the same. The orbitals within the band are so similar in energy that an electron does not need to gain appreciable energy to reach the conduction band. Also, as discussed above, the conduction band is only half filled to capacity which means that an electron is free to move through the entire metal. In an i nsulator , such as glass or plastic, the valence band is filled. Thus the next vacant higher-energy band becomes the conduction band. An energy gap exists between the valence band and the conduction band, as shown in Figure 20.10 of the text. This large separation prevents electrons in insulators from entering the conduction band. A s emiconductor , such as Si and Ge, has a filled valence band and an empty conduction band, but in contrast to an insulator, a relatively small gap exists between these bands. A relatively small amount of thermal energy will promote an electron into the conduction band. Thus as temperature increases, the conductivity of semiconductors increases correspondingly. _______________________________________________________________________________ EXAMPLE 20.3 Band Theory Show for sodium that the valence band is half-filled. •Method of Solution Sodium atoms have filled 1s, 2s, and 2p orbitals; therefore the corresponding bands in the solid are filled. The valence band for Na is the 3s. In the metal the 3s orbitals of N sodium atoms overlap to form a total of N molecular orbitals that make up the valence band. The capacity of this band is 2N electrons. Since each Na atom has one 3s electron, the N Na atoms have N electrons in the valence band. This means that the valence band is half-filled. The 3s band of Na is the conduction band. •Comment The highest occupied band is called the valence band. This unfilled valence band is called the conduction band of sodium metal. _______________________________________________________________________________ EXERCISES 6. 7. Back Define the terms valence band and conduction band. Define the terms conductor, insulator, and semiconductor. Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website Metallurgy and the Chemistry of Metals / 4 07 THE ALKALI METALS STUDY OBJECTIVES 1. 2. Describe the chemical properties of the alkali metal elements. Describe the sources and preparation of sodium and potassium metals. Properties of the Alkali Metals. As a group, the alkali metals are extremely reactive and never occur naturally in the elemental form. They are the least electronegative group of elements and exist as +1 cations combined with halides, sulfate, carbonate, and silicate ions. A number of properties of these elements are listed in Table 20.4 of the text. Compared with other metals, they are very soft and have low melting points and low densities. Lithium, sodium and potassium will float on water. The low density is a result of their large atomic volume and the fact that they all possess a body-centered crystal structure which has a low packing efficiency. Sources and Preparations. The preparation of Group 1A metals requires large amounts of energy because their positive ions are difficult to reduce. Lithium, being the most active metal, is prepared by electrolysis of molten LiCl. Sodium and potassium occur in a wide variety of minerals, and are the sixth and seventh most abundant elements in Earth's crust. Sodium compounds are so abundant and wide-spread that it is difficult to find matter free of this element. Sodium chloride makes up about two-thirds of the solid matter dissolved in sea water. The minerals carnallite (KMgCl 3 ·6H2 O), and sylvite (KCl) are found in ancient lake and seabeds and serve as commercial sources of potassium and its compounds. Both metals, sodium and potassium, were first prepared in 1807 by the English chemist Humphry Davy. He used the electrolysis of the corresponding moist hydroxides. As their discoverer, Davy was allowed to name the elements giving them the above names. However, you will notice that neither has an element symbol consistent with its name. At the time both potash and soda were recognized as carbonates of these unisolated metal elements. Potassium obtained its name because it was the metallic element in potash, potassium carbonate. The metallic element of soda (sodium carbonate) Davy called sodium. Sodium and Potassium. Metallic sodium is prepared commercially by the electrolysis of molten sodium chloride. Most of the sodium made in the United States is produced in the Downs cell (Section 19.8) by the electrolysis of a mixture of sodium chloride and calcium chloride. This electrolyte mixture melts at 505°C, whereas pure NaCl melts at 801°C. The lower temperature reduces the cost of production. A Downs cell consists of a carbon anode and an iron cathode. The chloride ions are oxidized at the anode, and the sodium ions are reduced at the cathode. The half-reactions are: anode cathode 2Cl – → C l 2 (g) + 2e– 2Na+ + 2e – → 2Na(l) The liquid sodium is drawn off and kept from contact with chlorine and oxygen. Sodium can also be obtained by the electrolysis of molten sodium hydroxide. Potassium is made by reaction of potassium chloride with sodium vapor in the absence of air at 900°C. KCl(l) + Na(g) → K(g) + NaCl(l) This reaction should occur to only a small extent. However, removal of potassium vapor as it is formed drives the reaction to completion. Reactions of Alkali Metals. The alkali metals are extremely reactive. The reactivity increases with atomic mass. Sodium forms two oxides on reaction with oxygen: sodium oxide (Na2 O) and sodium peroxide (Na2 O2 ). 2Na(s) + 1 /2 O2 (g) → Na2 O(s) 2Na(s) + O2 (g) → Na2 O2 (s) Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website 4 08 / Metallurgy and the Chemistry of Metals The reaction of potassium with oxygen forms the peroxide (K2 O2 ), but in addition when K burns in air it forms potassium superoxide K(s) + O 2 (g) → KO2 (s) Potassium superoxide is a source of O2 in breathing equipment. KO2 reacts with water: 2KO2 (s) + 2H 2 O(l) → 2KOH(aq) + O2 (g) + H 2 O2 (aq) Sodium chloride, sodium carbonate, sodium bicarbonate, sodium hydroxide, sodium nitrate, potassium hydroxide, and potassium nitrate are important compounds of these elements. EXERCISES 8. 9. List four physical properties of the alkali metals. List three chemical properties of the alkali metals. THE ALKALINE EARTH METALS STUDY OBJECTIVES 1. 2. Write chemical equations to show the preparations of magnesium and calcium. Describe the chemical properties and give the major uses of these elements and their compounds. Properties of the Alkaline Earth Metals. The alkaline earth metals are somewhat less reactive than the alkali metals. Except for beryllium the alkaline earth metals have similar chemical properties. The oxidation number of the Group 2A metals in almost all compounds is +2. The electronegativity of these elements is low enough that they form predominately ionic compounds with nonmetals. Table 20.5 in the text lists some common properties of these metals. As with the other groups of elements in the periodic table, metallic character increases as you read down the group. The properties of magnesium and calcium are reviewed below. The formulas of the important ores of magnesium and calcium are given in Table 20.1. Table 20.1 Ores Containing Magnesium and Calcium ________________________________________________________________ Element Rank Ore Formula (Earth's crust) ________________________________________________________________ Magnesium 6th brucite Mg(OH)2 dolomiteMgCO 3 ·CaCO 3 epsomite MgSO 4 ·7H2 O Calcium5th limestone CaCO3 gypsum CaSO4 ·2H2 O fluorite CaF 2 ________________________________________________________________ Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website Metallurgy and the Chemistry of Metals / 4 09 Preparation of Magnesium and Calcium. Seawater is an important source of magnesium. It contains 0.13 percent Mg 2+ by weight. Mg(OH) 2 being an insoluble compound is precipitated from seawater by adding Ca(OH)2 . Mg2+ (aq) + Ca2+ (aq) + 2OH– (aq) → Mg(OH)2 (s) + Ca 2+ (aq) Magnesium hydroxide is then converted to magnesium chloride by neutralization with hydrochloric acid. Metallic magnesium is obtained by electrolysis of molten magnesium chloride. electrolysis Mg2+ + 2 Cl – → Mg(l) + Cl2 (g) Calcium is obtained from limestone. Heating limestone causes it to decompose into calcium oxide and carbon dioxide. Calcium metal is prepared by a thermal reduction process, rather than electrolysis. Calcium oxide is reduced with aluminum at high temperature (1200°C). CaCO3 (s) → CaO(s) + CO2 (g) 6CaO(s) + 2Al(l) → 3Ca(l) + Ca3 Al2 O6 (s) Properties of Magnesium and Calcium. Magnesium is less reactive than calcium. Magnesium does not react with cold water, but it reacts slowly with steam. Calcium, on the other hand, reacts with cold water. Mg(s) + H 2 O(g) → MgO(s) + H 2 (g) Ca(s) + 2H 2 O(l) → Ca(OH)2 (aq) + H2 (g) Both MgO and CaO react with water to give hydroxides. MgO(s) + H 2 O(l) → Mg(OH)2 (s) (reacts with steam) CaO(s) + H2 O(l) → Ca(OH)2 (s) (reacts with cold water) When magnesium burns in air, considerable nitride is formed along with the oxide. 2Mg(s) + O 2 (g) → 2MgO(s) 3Mg(s) + N 2 (g) → Mg 3 N2 (s) Uses of the Alkaline Earth Metals. Beryllium has excellent alloying qualities and is used to make alloys that are corrosion resistant. Beryllium is used as a "window" in X-ray tubes. The major uses of magnesium are also in alloys. Because of its low atomic mass, it is a good lightweight structural metal. It is also used in batteries, and for cathodic protection of buried metal pipelines and storage tanks. Metallic calcium finds use mainly in alloys. Calcium salts are used as dehydrating agents; anhydrous calcium chloride, for example, has a strong affinity for water. Quicklime (CaO) is used is steel production, and in the removal of SO2 from coal-fired power plants. Slaked lime (Ca(OH)2 ) is used in water treatment. Strontium nitrate and carbonate are used in fireworks and highway flares to provide their brilliant red color. Barium metal is the most active of the alkaline earth metals and so it has very few uses. _______________________________________________________________________________ EXAMPLE 20.4 Properties of Alkaline Earth Metals Identify the element among the alkaline earth metals (excluding radium) that will have the following properties. a. Most reactive oxide towards water. b. Lowest electronegativity c. Smallest atomic radius d. Lowest first ionization energy. Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website 4 10 / Metallurgy and the Chemistry of Metals •Method of Solution Applying peroidic trends in these properties. a. Ba because reactivity increases as you read down in the group. b. Ba because electronegativity decreases as you read down in the group. c. Be because atomic radius increases as you read down in the group. d. Ba because ionization energy decreases as you read down in the group. _______________________________________________________________________________ EXAMPLE 20.5 Electrolytic Reduction of Calcium Starting with Ca2+ ions in limestone, write equations to show how you would obtain pure calcium, using electrolysis, rather than thermal reduction. •Method of Solution Electrolysis of molten metal chlorides is often used to prepare pure metals. First limestone, chalk, or sea shells, all of which contain CaCO 3 , must be decomposed at 900°C. CaCO3 (s) → CaO(s) + CO2 (g) The CaO(s) is slaked to yield calcium hydroxide. CaO(s) + H2 O(l) → Ca(OH)2 (s) This base is then neutralized with hydrochloric acid to give the desired salt CaCl2 . Ca(OH)2 (aq) + 2HCl(aq) → CaCl2 (aq) + 2H2 O(l) After drying, the CaCl2 can be melted and then electrolyzed. e lectrolysis CaCl 2 (l) → Ca(l) + Cl2 (g) _______________________________________________________________________________ EXERCISES 10. Give the sources of magnesium and calcium. 11. Write chemical formulas for limestone, quicklime, and slaked lime. Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website Metallurgy and the Chemistry of Metals / 4 11 ALUMINUM STUDY OBJECTIVES 1. 2. 3. Write chemical equations to show the reactions of aluminum with acids, bases, oxygen, and other metal oxides. Write chemical equations that show the amphoteric nature of Al(OH)3 . List several uses of aluminum and its compounds. Source and Preparation. Aluminum is the third most abundant element in the earth's crust, making up 7.5 percent by mass. It is too active chemically to occur free in nature, but is found in compounds in over 200 different minerals. The most important ore of Al is bauxite, which contains hydrated aluminum oxide (Al2 O3 ·2H2 O), along with silica and hydrated iron oxide (Fe2 O3 ·2H2 O). Another important mineral is cryolite (Na3 AlF 6 ), which is used in the metallurgy of aluminum. The Hall process for the production of aluminum metal is described in Figure 20.20 of the text. Aluminum is obtained from bauxite ore (Al 2 O3 ·2H2 O). The first step in its preparation is to separate pure aluminum oxide from the silica and iron impurities. First, bauxite is pulverized and digested with sodium hydroxide solution. This converts the silica into soluble silicates, and aluminum oxide is converted to the – aluminate ion AlO 2 (aq) which remains in solution. However this digestion treatment has no effect on the iron which remains as insoluble Fe 2 O3 (s). It is removed by filtration. Aluminum hydroxide is then precipitated by acidification (with carbonic acid) to about pH 6. The precipitate is heated strongly to produce pure Al2 O3 . The chemical changes involving aluminum are: – Al2 O3 (s) + 2OH – (aq) → 2AlO2 (aq) + H2 O(l) – AlO2 (aq) + H+(aq) + H2 O(l) → Al(OH)3 (s) 2Al(OH)3 (s) → Al 2 O3 (s) + 3H 2 O(g) The aluminum ions in Al2 O3 can be reduced to metallic aluminum efficiently only by electrolysis. The melting point of Al2 O3 is 2050°C, which makes electrolysis of pure molten Al2 O3 extremely expensive owing to the need to maintain the high temperature. In the Hall process, Al2 O3 is dissolved in molten cryolite (Na3 AlF 6 ) which melts at 1000°C. The use of cryolite makes it possible to lower the temperature of electrolysis by 1050°C! The mixture is electrolyzed to produce aluminum and oxygen. cathode 2Al 3+ + 6e – → 2Al(l) 3 anode 3O2– → 2 O 2 + 6e – ________________________________________ overall 2Al3+ + 3O2– → 2Al(l) + 3 /2 O2 Properties. Pure aluminum is a silvery-white metal with low density (2.7 g/cm3 ) and high tensile strength. It is malleable and can be rolled into thin foils. Its electrical conductivity is about 65 percent that of copper. Aluminum is an amphoteric element, reacting with both acids and bases. 2Al(s) + 6HCl(aq) → AlCl3 (aq) + 3H2 (g) 2Al(s) + 2NaOH(aq) + 2H2 O(l) → 2NaAlO2 (aq) + 3H2 (g) Aluminum has a strong affinity for oxygen. Indeed the metal is usually covered by an oxide film. 4Al(s) + 3O 2 (g) → 2Al 2 O3 (s) Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website 4 12 / Metallurgy and the Chemistry of Metals This layer of Al 2 O3 forms a compact, adherent, protective, surface coating on the metal and is responsible for preventing further oxidation and corrosion. Because of this surface oxide layer, Al is practically insoluble in weak or dilute acids. o The large enthalpy of formation of aluminum oxide (∆ Hf = –1670 kJ/mol) makes the metal an excellent reducing agent. Thus a variety of metals can be produced in a series of similar reactions involving Al powder with the corresponding metal oxides. So much heat is liberated in these reactions, called aluminothermic reactions, that the metal is usually obtained in the molten state. The thermite reaction, used in welding steel and iron, is one example. 2Al(s) + Fe 2 O3 (s) → Al 2 O3 (l) + 2Fe(l) Aluminum hydroxide is an amphoteric hydroxide, dissolving in both acid and base. Al(OH)3 (s) + 3H +(aq) → Al 3+ (aq) + 3H2 O(l) – Al(OH)3 (s) + OH – (aq) → Al(OH)4 (aq) Uses. Aluminum metal is used in high-voltage transmission lines because it is cheaper and lighter than copper. Its chief use is in aircraft construction. Aluminum is employed also as a solid propellant in rockets. This is another example of the great affinity that aluminum has for oxygen. Ammonium perchlorate, NH4 ClO 4 , is the oxidizer in the reaction. The formation of aluminum hydroxide precipitate is used in water treatment plants. The process requires large amounts of aluminum sulfate. Al2 (SO4 )3 (aq) + 3Ca(OH)2 (aq) → 2Al(OH)3 (s) + 3CaSO 4 (aq) Aluminum hydroxide is a gelatinous substance. As it settles in treatment pools it coprecipitates suspended matter such as bacteria and colloidal sized particles. This process clarifies drinking water. Alums are compounds which have the general formula: M+M3+ (SO4 )2 ·12H2 O + where M+ = K+, Na +, NH 4 , and M 3+ = Al 3+ , Cr 3+ , Fe 3+ . They are used in the dying industry where the formation of the gelatinous aluminum hydroxide plays in "fixing" the dye to the cloth. _______________________________________________________________________________ EXAMPLE 20.6 Amphoterism Explain the meaning of the term amphoterism. Write equations for the amphoteric behavior of aluminum. •Method of Solution An amphoteric substance can react either as an acid or as a base. Aluminum reacts with hydrochloric acid and with strong bases as follows: 2Al(s) + 6HCl(aq) → AlCl3 (aq) + 3H2 (g) 2Al(s) + 2NaOH(aq) + 2H2 O(l) → 2NaAlO2 (aq) + 3H2 (g) _______________________________________________________________________________ EXAMPLE 20.7 Oxidation of Aluminum Aluminum is a good reducing agent, and its reduction potential is quite negative (E° = –1.66 V) which means that aluminum should react with water and liberate hydrogen. But we know that airplanes do not dissolve in thunderstorms. Explain. Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website Metallurgy and the Chemistry of Metals / 4 13 •Method of Solution Aluminum readily forms the oxide Al2 O3 when exposed to air. The oxide forms a tenacious surface film that protects the aluminum metal from further corrosion due to water. Therefore it is the presence of this oxide layer that makes aluminum practically insoluble even in dilute acids. _______________________________________________________________________________ EXERCISES 12. What is the role of cryolite, Na 3 AlF 6 , in aluminum production? 13. Write chemical equations that show the amphoteric nature of Al(OH)3 . 14. The enthalpy of formation of Al 2 O3 is large and exothermic. Why doesn't aluminum metal just completely oxidize away. __________________________________________________________________________ PRACTICE TEST 1. Distinguish between: a. Leaching and flotation b. Roasting and reduction 2. Can carbon be used to reduce Al2 O3 to Al(s)? Hint: Is the following reaction spontaneous? 2Al2 O3 (s) + 2C(s) → 4Al(s) + 3CO2 (g) 3. When a solution of NaOH is added dropwise to a test tube containing a solution of Al3+ , a white gelatinous precipitate is formed. Upon continued addition of NaOH, the precipitate disappears. Write the chemical equations to explain this observation. 4. In the production of sodium, the metal is prevented from coming into contact with Cl2 and O2 . Write chemical equations for the reactions of these elements with sodium. 5. What is the agent that actually reduces Fe in the blast furnace? Write a chemical equation to show the reaction. 6. Distinguish between an alloy and an amalgam. 7. According to the band theory, why is copper a conductor? 8. Considering that electrons in metals have random thermal motion like molecules of gas, explain why electrical conductivity of conductors decreases with increasing temperature? 9. Why does the electrical conductivity of semiconductors increase with increasing temperature? 10. What are the sources of aluminum and potassium? 11. What is quicklime? How is it made? 12. Aluminum metal was first prepared in 1825 by the action of potassium metal on aluminum chloride. Write a balance equation for this reaction. 13. What is the chemical form of aluminum in bauxite? 14. How is aluminum separated from Fe2 O3 ? 15. What is an alum? Which alum is used in baking powder? Hint: Read a label. What is its function in baking powder? 16. What is the role of aluminum sulfate in water purification? 17. What is the pH of a 0.10 M aluminum chloride solution if it is hydrolyzed to the extent of 10 percent at room temperature. [Al(H2 O)6 ]3+ → H+ + [Al(H 2 O)5 OH]2+ Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website 4 14 / Metallurgy and the Chemistry of Metals ANSWERS Exercises 1. A mineral is a naturally occurring substance with a characteristic range of chemical composition. An ore is a mineral, or a mixture of minerals, from which a particular metal can be profitably extracted. 2. The electropositive metals can be reduced by a more electropositive metal such as Li, but there is no metal electropositive enough to reduce Li+ to Li. Electrolytic reduction is the only way to prepare Li. It is possible to reduce other metals with Li. Rather than prepare Li by electrolysis and then use it to reduce aluminum, magnesium, or sodium, it is more convenient simply to prepare these metals in one step by electrolytic reduction. 3. Zone refining is a technique for purifying metals. Figure 20.8 in the text explains how it works. 4. Limestone is used to remove impurities from iron during production. Limestone decomposes to lime (CaO). CaO reacts with SiO 2 and Al2 O3 impurities from the iron ore forming calcium silicate and calcium aluminate. The mixture of calcium silicate and calcium aluminate is known as slag. 5. Cr2 O3 (s) + 2Al(s) → 2Cr(l) + Al2 O3 (s) 6. A band is a large number of molecular orbitals that are closely spaced in energy. The valence band is a set of closely spaced MOs that are filled with electrons. The conduction band is a set of closely spaced empty orbitals. An electron can travel freely through the metal since the conduction band is void of electrons. 7. A conductor is capable of conducting an electrical current. In a conductor there is essentially no energy gap between the valence band and the conduction band. Insulators are materials that do not conduct. In an insulator the energy gap between the valence band and the conduction band is much greater than the gap in a conductor. Semiconductors are normally not conductors, but will conduct electricity at higher temperatures, or when combined with small amounts of certain elements. 8. Good conductors of heat and electricity. Low density. Soft enough to cut with a knife. Low melting points. 9. Extremely reactive. React with water to form hydroxides. React with oxygen to form a varity of oxides, peroxides and superoxides. 10. Magnesiun is precipitated from seawater. Calcium is obtained from limestone. 11. CaCO3 , CaO, and Ca(OH)2 respectively. 12. Molten cryolite is used as a solvent for alumina, Al2 O3 , in aluminum production. Cryolite melts at 1000°C as compared to 2050°C for alumina and therefore it lowers the energy consumption. 13. Reaction with acid: Al(OH)3 (s) + 3H +(aq) → Al 3+ +(aq) + 3H2 O(l) – Reaction with base: Al(OH)3 (s) + OH – (aq) → Al(OH)4 (aq) 14. A fresh surface of aluminum does react readily with oxygen and forms a surface coating of Al2 O3 . This layer is very impenetrable and does not allow oxygen and water through to continue a reaction with the rest of the aluminum. Practice Test 1. a. Leaching is the selective dissolution of a metal from an ore. Flotation is a technique used to separate mineral particles from waste clays and silicates called gangue. b. Roasting involves heating an ore in the presence of air. The idea is to convert metal carbonates and sulfides to oxides, which can be more conveniently reduced to yield pure metals. Reduction is the gain of electrons by a cation to yield the pure metal. 2. No. Use Appendix 3 of the text to calculate ∆ G°. The sign of ∆ G° indicates whether the reaction will go or not. 3. Al3+ + 3OH – → Al(OH)3 (s) – Al(OH)3 (s) + OH – → Al(OH)4 (aq) Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website Metallurgy and the Chemistry of Metals / 4 15 4. 2Na + Cl2 → 2NaCl 2Na + O 2 → Na2 O2 1 2Na + 2 O 2 → Na2 O 5. Carbon monoxide. Fe2 O3 + 3CO → 2Fe + 3CO2 6. An alloy is a mixture of two or more metals, or of metals and nonmetal elements that have metallic properties. An amalgam is an alloy containing mercury. 7. Copper atoms have a half-filled 4s orbital. This means that copper metal has a half-filled valence band. A partially filled valence band is a conduction band. 8. The increasingly random motion of electrons due to a temperature rise opposes motion in one unified direction. 9. As temperature increases more electrons acquire enough energy to "jump" the energy gap between the valence band and the conduction band. 10. The source of Al is bauxite. Potassium is obtained from sylvite. 11. Quicklime is CaO. It is made the thermal decomposition of limestone. 12. 3K + AlCl3 → Al + 3KCl 13. Al2 O3 ·2H2 O 14. The bauxite is pulverized and digested with sodium hydroxide solution. This converts the aluminum oxide – to the aluminate ion AlO2 (aq) which remains in solution. However treatment with base has no effect on the iron oxide which remains as insoluble Fe 2 O3 (s), and is removed by filtration. Aluminum hydroxide is then precipitated by acidification to about pH 6. + 15. Alums are sulfates with the general formula M+M3+ (SO4 )2 ·12H2 O, where M+ = K+, Na +, NH 4 , and M3+ = Al 3+ , Cr 3+ , Fe 3+ . Sodium aluminum sulfate is used in baking powder as an anticaking agent. 16. Al2 (SO4 )3 is used to clarify water through its reaction with Ca(OH)2 . The precipitation of Al(OH)3 traps dirt and dust particles. 17. pH = 2 ____________________________________________________________________________ Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website ...
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This note was uploaded on 09/15/2009 for the course CHEM 102 taught by Professor Bastos during the Spring '08 term at Adelphi.

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