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22-1 Lecture PowerPoint Chemistry The Molecular Nature of Matter and Change Sixth Edition Martin S. Silberberg Copyright  The McGraw-Hill Companies, Inc.

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Presentation on theme: "22-1 Lecture PowerPoint Chemistry The Molecular Nature of Matter and Change Sixth Edition Martin S. Silberberg Copyright  The McGraw-Hill Companies, Inc."— Presentation transcript:

1 22-1 Lecture PowerPoint Chemistry The Molecular Nature of Matter and Change Sixth Edition Martin S. Silberberg Copyright  The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

2 22-2 Chapter 22 The Elements in Nature and Industry

3 22-3 The Elements in Nature and Industry 22.1How the Elements Occur in Nature 22.2The Cycling of Elements Through the Environment 22.3Metallurgy: Extracting a Metal from Its Ore 22.4Tapping the Crust: Isolation and Uses of Selected Elements 22.5Chemical Manufacturing: Two Case Studies

4 22-4 Figure 22.1The layered internal structure of Earth. The processes by which the Earth formed led to its differentiation into different layers.

5 22-5 Elemental Abundance The abundance of an element is the amount of that element in a particular region of the natural world. Elements are not equally abundant in all regions – abundances differ due to the differences in physical and chemical behavior of the elements. - The core of the Earth is rich in dense Group 8B(8) to 8B(10) metals. - The Earth’s crust has the largest share of nonmetals, metalloids, and light active metals.

6 22-6 Table 22.1 Cosmic and Terrestrial Abundances (Mass %) of Selected Elements.

7 22-7 Compositional Phases of the Earth As the Earth cooled to form its major layers, gravity and convection caused materials of different densities to separate, giving several phases. Fe was the major component of the core or iron phase. The outer silicate phase, containing oxygen combined with Si, Al, Mg, and some Fe, separated into the mantle and crust. The sulfide phase, with intermediate density, consisted mostly of iron sulfide mixed with parts of the other phases.

8 22-8 Figure 22.2Geochemical differentiation of the elements.

9 22-9 Distribution of the Elements The distribution of the elements in the Earth’s layers was controlled by their chemical affinity for one of the three phases. Elements with low or high electronegativity tended to congregate in the silicate phase as ionic compounds. These included active metals and nonmetals. Metals with intermediate EN dissolved in the iron phase. Lower-melting transition metals, and many metals and metalloids in Groups 11 to 16, became concentrated in the sulfide phase.

10 22-10 Impact of Life on Crustal Abundances Photosynthesis resulted in an increase in the O 2 levels of the atmosphere. –Oxidation became the major source of free energy in the crust and biosphere. The [K + ] of the oceans is much lower than the [Na + ] –since growing plants have absorbed dissolved K +. Subterranean deposits of organic carbon formed from the decomposition of ancient plants under high pressure. Fossilized skeletal remains of early marine organisms have formed vast deposits of C, O and Ca.

11 22-11 Figure 22.3Ancient effect of an O 2 -rich atmosphere. Banded-iron formations containing Fe 2 O 3 Fossil of early multicellular organism

12 22-12 Table 22.2 Abundances of Selected Elements in the Crust, Its Regions, and the Human Body as Representative of the Biosphere (Mass %)

13 22-13 Sources of the Elements O 2, N 2, and the noble gases (except He) are obtained from the atmosphere, where they occur as the free elements. A few elements occur naturally in their uncombined (native) state. These include S, carbon in coal, and unreactive metals. Most elements occur in ores, natural compounds or mixtures from which an element must be extracted.

14 22-14 Figure 22.4Sources of the elements.

15 22-15 Environmental Cycles An environmental cycle is a natural process in which elements are continuously cycled in various forms between different regions of the Earth’s crust. The most important of these cycles from the perspective of living organisms are the carbon, nitrogen, and phosphorus cycles. Elements are cycled through physical, biological, and chemical pathways.

16 22-16 Sources of Carbon Carbon appears in all three portions of the Earth’s crust. Elemental C occurs as diamond and graphite in the lithosphere. C is a component of carbonate minerals. C forms the basis of all organic substances. CO 2 (g) is found in the atmosphere and dissolved in rivers, lakes and seas.

17 22-17 The Carbon Cycle Fixation is the process of converting a gaseous substance into a condensed form. Through photosynthesis, plants use sunlight to fix atmospheric CO 2 into carbohydrates. All living organisms release CO 2 by respiration and by decomposition. Animals eat the plants, incorporating the carbon into their systems. The portions of the carbon cycle are linked by the atmosphere – one link between oceans and air, and the other between land and air.

18 22-18 Figure 22.5 The carbon cycle.

19 22-19 The Nitrogen Cycle The nitrogen cycle involves a direct interaction between land and sea. Atmospheric N 2 must be fixed to enter the land and sea. –Atmospheric fixation occurs when lightning provides the energy for the reaction between N 2 (g) and O 2 (g). –Industrial fixation results from the production of NH 3, which is used to make fertilizers. –Biological fixation involves the conversion of atmospheric N 2 to NO 3 - by blue-green algae and nitrogen-fixing bacteria. Denitrifying bacteria reduce NO 3 - back to N 2 in various stages, allowing it to cycle back to the atmosphere.

20 22-20 Figure 22.6The nitrogen cycle.

21 22-21 The Phosphorus Cycle P cycles between land and sea in the form of phosphate (PO 4 3- ) and the various hydrogen phosphate ions. There are three interlocking subcycles: –The inorganic cycle involves slow weathering of phosphate- containing rocks, which causes PO 4 3- to leach into the rivers and seas. –The land-based biological cycle involves incorporation of PO 4 3- into organisms and its release through excretion and decomposition. –The water-based biological cycle consists of the cycling of phosphate oxoanions by aquatic organisms.

22 22-22 Figure 22.7The phosphorus cycle.

23 22-23 Figure 22.8Industrial uses of phosphorus.

24 22-24 Metallurgy Metallurgy is the technology of metals, and is concerned with their extraction and utilization. Pyrometallurgy uses heat to obtain the metal. Electrometallurgy employs an electrochemical step. Hydrometallurgy relies on the metal’s aqueous solution chemistry Extraction of a pure element from its ore involves a series of steps, beginning with mining the ore. The mineral contains the element while the gangue is the portion of the ore with no commercial value.

25 22-25 Table 22.3 Common Mineral Sources of Some Elements ElementMineral, Formula AlGibbsite (in bauxite), Al(OH) 3 BaBarite, BaSO 4 BeBeryl, Be 3 Al 2 Si 6 O 18 CaLimestone, CaCO 3 FeHematite, Fe 2 O 3 HgCinnabar, HgS NaHalite, NaCl PbGalena PbS SnCassiterite, SnO 2 ZnSphalerite, ZnS

26 22-26 Figure 22.9Steps in metallurgy. Mining Pretreating Magnetic attraction Cyclone separation Flotation Leaching Converting (mineral to compound) Pyrometallurgy (roasting, etc.) Hydrometallurgy Converting (compound to metal) Chemical redox (smelting, etc.) Electrochemical redox Refining Electrorefining Distillation Zone refining Alloying

27 22-27 Pretreating the Ore Magnetic attraction can be used to separate magnetic minerals from the gangue. Density separation exploits large differences in densities. –Panning for gold relied on the density differences between gold and sand. Flotation uses an oil-detergent mixture to form a slurry, which is mixed rapidly with air to give an oily, mineral- rich froth that floats. Leaching involves the formation of a complex ion that is water soluble.

28 22-28 Figure 22.10The cyclone separator. The cyclone separator blows high-pressure air through the pulverized mixture to separate the particles. Panning for gold also relies on density differences.

29 22-29 Figure 22.11The flotation process.

30 22-30 Conversion of Minerals by Roasting Minerals are often converted to their oxides because oxides can be reduced easily. Δ CaCO 3 (s) CaO(s) + CO 2 (g) Metal sulfides are converted to oxides by roasting in air: Δ 2ZnS(s) + 3O 2 (g) 2ZnO(s) + 2SO 2 (g)

31 22-31 Conversion of Minerals by Chemical Redox Reduction with carbon is common. Heating an oxide with a reducing agent to obtain the metal is called smelting. ZnO(s) + C(s) → Zn(g) + CO(g) Oxides of less active metals are reduced with hydrogen instead of carbon. WO 3 (s) + H 2 (g) → W(s) + 3H 2 O(g) A more active metal may be used as the reducing agent. Cr 2 O 3 (s) + 2Al(s) → 2Cr(l) + Al 2 O 3 (s)  H° << 0

32 22-32 Figure 22.12The thermite reaction. Al powder is used to reduce a metal oxide in a spectacular exothermic reaction, which produces the molten metal.

33 22-33 Electrochemical Redox in Mineral Conversion The mineral is converted to the element in an electrolytic cell. Often the pure mineral is used to prevent unwanted side reactions. The cation is reduced to the metal at the cathode, and the anion is oxidized to the nonmetal at the anode: BeCl 2 (l) → Be(s) + Cl 2 (g) Specially designed cells separate the products to prevent recombination.

34 22-34 Figure 22.13The redox step in converting a mineral to the element.

35 22-35 Refining and Alloying the Element Refining (purifying) the element is carried out by one of the following three methods: –Electrorefining employs an electrolytic cell, in which the impure metal acts as the anode and the pure metal as the cathode. –Distillation is used to refine metals with relatively low boiling points. –Zone refining exploits freezing point depression. Many metals are in alloys rather than their pure forms. Alloying involves combining the metals in the correct proportions to produce a material of the desired physical properties.

36 22-36 Figure 22.14Zone refining of silicon. The impure metal or metalloid is passed slowly through a heating coil in an inert atmosphere. A small zone of the impure solid melts, and as the next zone melts, the dissolved impurities from the first zone lowers the freezing point. The purer solid of the first zone refreezes. Gradually impurities from each zone move into the adjacent zone, producing pure solid at the end of the rod.

37 22-37 Table 22.4 Some Important Alloys and their Composition

38 22-38 Figure Three binary alloys. vanadium carbideCu 3 Au  -brass

39 22-39 Producing Sodium and Potassium Sodium ore is halite, which is obtained either by the evaporation of brines or by mining salt deposits. Na is extracted and purified in an electrolytic apparatus called the Downs cell. Sylvite (mostly KCl) is the major ore of potassium. Chemical reduction of K + ions by Na at high temperature produces the metal: Na(l) + K + (l) Na + (l) + K(g) The reaction is carried out at 850°C so that gaseous K is produced. As the K(g) is removed the equilibrium shifts to produce more.

40 22-40 Figure 22.16The Downs cell for production of sodium. CaCl 2 is mixed with the NaCl to lower the melting point.

41 22-41 Evaporative salt beds near San Francisco Bay. Firefighter with KO 2 breathing mask.

42 22-42 Metallurgy of Iron Iron ores contain iron oxides, carbonates and sulfides. Iron is recovered from its ores by reduction with C in a blast furnace. –This process produces impure pig iron. Pig iron is converted to steel by means of the basic- oxygen process. –High pressure O 2 is blown over and through the molten iron, oxidizing the impurities, which are converted to a molten slag and removed. Steel is produced by alloying iron with other metals to prevent corrosion and increase both strength and flexibility.

43 22-43 Table 22.5 Important Minerals of Iron Mineral TypeMineral, Formula OxideHematite, Fe 2 O 3 Magnetite, Fe 3 O 4 Ilmenite, FeTiO 3 CarbonateSiderite, FeCO 3 SulfidePyrite, FeS 2 Pyrrhotite, FeS

44 22-44 Figure 22.17The major reactions in a blast furnace.

45 22-45 Figure 22.18The basic-oxygen process for making steel. Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

46 22-46 Isolation of Copper The most common copper ore is chalcopyrite, a mixed ore of CuFeS 2, which contains 0.5% Cu by mass. Flotation of the ore removes some of the iron, enriching the mass % of Cu. Roasting oxidizes the FeS to FeO but leaves the CuS unreacted. –2FeCuS 2 (s) + 3O 2 (g) → 2CuS(s) + 2FeO(s) + 2SO 2 (g) Heating in sand converts FeO to FeSiO 3, a molten slag. –FeO(s) + SiO 2 (g) → FeSiO 3 (l) –CuS → Cu 2 S at this high temperature.

47 22-47 Isolation of Copper The final smelting step converts some Cu 2 S to Cu 2 O by roasting in air. Cu 2 S and Cu 2 O react: –2Cu 2 S(s) + 3O 2 (g) →2Cu 2 O(s) + 2SO 2 (g) –Cu 2 S(s) + 2Cu 2 O(s) → 6Cu(l) + SO 2 (g)

48 22-48 Mining chalcopyrite.

49 22-49 Figure 22.19The electrorefining of copper. Copper for wiring needs to be 99.99% pure. This is achieved by electrorefining. The impure Cu is used as a series of anodes, and the cathodes are made from purified Cu. An acidic solution of CuSO 4 serves as the electrolyte. Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

50 22-50 Isolation of Aluminum Al is the most abundant metal in the Earth’s crust. –The major ore of Al is bauxite, which contains Al 2 O 3. The Bayer process is the pretreatment of bauxite by boiling with NaOH. –Al 2 O 3 and SiO 2 are amphoteric and dissolve in the base, while Fe 2 O 3 and TiO 2 do not. –The silicate is precipitated by heating, and the filtrate acidified to precipitate Al 3+ as Al(OH) 3. –Drying at high temperature converts this to Al 2 O 3. Al 2 O 3 is converted to the free metal by the Hall-Heroult electrolytic process.

51 22-51 Mining bauxite.

52 22-52 Figure 22.20The electrolytic cell in the manufacture of aluminum.

53 22-53 Figure 22.21The many familiar and essential uses of aluminum. Although Al is a very active metal, it does not corrode readily because when Al 2 O 3 forms on the metal surface it adheres to the metal and prevents more O 2 from penetrating. Al can be anodized to form a thicker protective oxide layer.

54 22-54 Isolation of Magnesium Magnesium is readily obtained from sea water, although it is also abundant on land. The Dow process is used to isolate Mg. –Ca(OH) 2 is generated from crushed seashells, then added to the seawater. –Ca(OH) 2 (aq) + Mg 2+ (aq) → Mg(OH) 2 (aq) + Ca 2+ (aq) –Mg(OH) 2 (s) + 2HCl(aq) → MgCl 2 (aq) + 2H 2 O(l) –Heating evaporates the water to give hydrated MgCl 2 ·nH 2 O(s). –The solid is heated and melted, and electrolysis is used to convert the MgCl 2 to molten Mg and Cl 2 (g).

55 22-55 Figure 22.22The Dow process for isolation of elemental Mg from seawater.

56 22-56 Sources of Hydrogen Very pure H 2 is prepared through electrolysis of water with Pt (or Ni) electrodes: 2H 2 O(l) + 2e - → H 2 (g) + 2OH - (aq) E = V [cathode; reduction] H 2 O(l) → ½O 2 (g) + 2H + (aq) + 2e - E = 0.82 V [anode; oxidation] H 2 O(l) → H 2 (g) + ½O 2 (g) E = V H 2 can also be produced by thermal methods. The most common of these use H 2 O and a simple alkane like CH 4. H 2 O(g) + CH 4 (g) → CO(g) + 3H 2 (g)  H° = 206 kJ H 2 O(g) + CO(g) CO 2 (g) + H 2 (g)  H° = -41 kJ This second reaction, called the water-gas shift reaction, is carried out using an iron or cobalt oxide catalyst and high temperatures.

57 22-57 Uses of Hydrogen Ammonia is synthesized using N 2 and H 2. H 2 is used in the hydrogenation of C=C double bonds in liquid oils to form C-C bonds in solid fats and margarine. –Partially hydrogenated vegetable fats and oils are present in a wide variety of foods. H 2 is essential in the manufacture of “bulk” chemicals, those produced in large amounts. –Methanol, used as a gasoline additive, is produced by the reaction of CO(g) and H 2 (g) over a copper-zinc catalyst. Cu-ZnO catalyst CO(g) + 2H 2 (g) CH 3 OH(l)

58 22-58 Hydrogen Isotopes Hydrogen has three isotopes; – 1 H (protium), 2 H or 2 D (deuterium), and 3 H or 3 T (tritium). The relative difference in the mass of these isotopes is very large, leading to significant differences in bond energies and reactivities. –A bond to H is much weaker and easier to break than a bond to deuterium. Any reaction that includes a bond to H or D in the rate- determining step will occur much faster with H than with D. This is called the kinetic isotope effect. –This effect is exploited in the isolation of deuterium, and is also used to study reaction mechanisms.

59 22-59 Table 22.6 Some Molecular and Physical Properties of Diatomic Protium, Deuterium, and Tritium. Property H 2 D 2 T 2 Molar mass (g/mol) Bond length (pm) Melting point (K) Boiling point (K)  H° fus (kJ/mol)  H° vap (kJ/mol) Bond energy (kJ/mol at 298 K)

60 22-60 Producing Sulfuric Acid Sulfur can be produced by the Claus process from H 2 S in “sour” natural gas: 2H 2 S(g) + 2O 2 (g) → ⅛S 8 (g) + SO 2 (g) + 2H 2 O(g) at low temperature 2H 2 S(g) + SO 2 (g) → ⅛S 8 (g) + 2H 2 O(g) using an Fe 2 O 3 catalyst The Frasch process is used when natural deposits of S are found: Superheated water is pumped into the deposit to melt the sulfur, which is forced to surface by compressed air. Sulfur is burned in air to form SO 2 : ⅛S(s) + O 2 (g) → SO 2 (g)  H° = -297 kJ

61 22-61 Figure 22.23The Frasch process for mining elemental sulfur.

62 22-62 Producing Sulfuric Acid The contact process oxidizes SO 2 to SO 3 : SO 2 (g) + ½O 2 (g) SO 3 (g)  H° = -99 kJ This reaction is exothermic and very slow at room temperature. A V 2 O 5 catalyst is used to optimize rate, and the yield of SO 3 is increased by increasing the pressure and removing the SO 3 as it forms. SO 3 cannot be added directly to water to form H 2 SO 4 because it will polymerize. A small amount of H 2 SO 4 is used instead: SO 3 (g) → H 2 S 2 O 7 (l) H 2 S 2 O 7 (l) + H 2 O(l) → 2H 2 SO 4 (l)

63 22-63 Figure 22.24The many indispensable applications of sulfuric acid.

64 22-64 The Chlor-Alkali Process The chlor-alkali process yields chlorine, which ranks among the top 10 chemicals in the United States. This process involves the electrolysis of concentrated aqueous NaCl. Diaphragm-Cell Method: 2Cl - (aq) → Cl 2 (g) + 2e - E° = 1.36 V [anode; oxidation] 2H 2 O(l) + 2e - → 2OH - (aq) + H 2 (g) + 2e - E° = -1.0 V [cathode; reduction] 2Cl - (aq) + 2H 2 O(l) → 2OH - (aq) + H 2 (g) + Cl 2 (g) E cell = -2.4 V To maximize the yield of Cl 2, a voltage almost twice this value and a current in excess of 3x10 4 A are used.

65 22-65 Figure 22.25A diaphragm cell for chlor-alkali process. Overall reaction: 2Na + (aq) + 2Cl - (aq) + 2H 2 O(l) → 2Na + (aq) + 2OH - (aq) + H 2 (g) + Cl 2 (g)

66 22-66 Mercury-Cell Method: High purity NaOH can be obtained with this method, which uses a mercury cathode. This creates such a large overvoltage for the reduction of H 2 O that Na + is reduced instead: 2Cl - (aq) → Cl 2 (g) + 2e - [anode; oxidation] 2Na + (aq) + 2e - → 2Na(Hg)[cathode; reduction] The sodium amalgam is pumped out of the system and treated with H 2 O, which is reduced by the Na: 2Na(Hg) + 2H 2 O(l) → 2Na + (aq) + 2OH - (aq) + H 2 (g)

67 22-67 Membrane-Cell Method: This method replaces the diaphragm with a polymeric membrane to separate the cell compartments. The membrane allows only cations to move through it, and only from the anode to the cathode compartment As Cl - ions are oxidized at the anode, Na + cations in the anode compartment move through the membrane to the cathode compartment and form an NaOH solution.


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