1 / 56

Chapter 16

Chapter 16. The Group 16 Elements. The “Chalcogens”. Group 16 Trends. Besides oxygen, all even oxidation states from +6 to –2 are observed. Anomalous Nature of Oxygen. More energetically favorable to form multiple bonds. Anomalous Nature of Oxygen. Inability to catenate. Oxygen Allotropes.

shelly
Download Presentation

Chapter 16

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Chapter 16 The Group 16 Elements

  2. The “Chalcogens”

  3. Group 16 Trends • Besides oxygen, all even oxidation states from +6 to –2 are observed

  4. Anomalous Nature of Oxygen • More energetically favorable to form multiple bonds

  5. Anomalous Nature of Oxygen • Inability to catenate

  6. Oxygen Allotropes • Dioxygen • colorless, odorless gas • does not burn, but supports combustion • all elements will combust with oxygen • except “noble metals” and noble gases • pyrophoricity

  7. Dioxygen • 21% of Earth’s atmosphere • Low solubility in water • 2 x 10-5 • DO (dissolved oxygen) vs. BOD (biological oxygen demand) • Fractional distillation of air • 109 tons used annually • steel industry • hospitals

  8. Dioxygen • Produced in the laboratory by the catalytic decomposition of hydrogen peroxide 2H2O2(aq)  2H2O(l) + O2(g) • Is oxygen diamagnetic or paramagnetic?

  9. Dioxygen

  10. Trioxygen • Ozone • diamagnetic gas with a strong odor • detected as low as 0.01 ppm • 0.1 ppm toxicity level

  11. Trioxygen • Produced by passing dioxygen through an electric field 3O2(g)  2O3(g) Hf° = +143 kJ/mol • Very strong oxidizing agent O3(g) + 2H+(aq) + 2e- O2(g) + H2O(l) E° = +2.07 V O2(g) + 4H+(aq) + 4e- 2H2O(l) E° = +1.23 V

  12. The Ozone Layer • Oxygen absorbs short wavelength UV radiation to produce atomic oxygen O2(g) + uv  2O(g) • Atomic oxygen reacts with dioxygen to give ozone O2(g) + O(g) O3(g) • Ozone absorbs longer wavelength UV radiation O3(g) + uv  O(g) + O2(g) O3(g) + O(g) 2O2(g)

  13. The Ozone Layer • Depleters O3(g) + X(g)  XO(g) + O2(g) O(g) + XO(g)  X(g) + O2(g) • where X = hydrogen atoms, hydroxyl radical, nitrogen monoxide, and chlorine atoms

  14. Trioxygen Complexes • Forms compounds with the alkali and alkaline earth metals • O3- anion (trioxide (1-)) • CsO3 • Ba(O3)2

  15. Bonding in Covalent Oxygen Compounds • H2O (104.5°), F2O (103°), and Cl2O (111°) • Bent rule • more electronegative elements have more “p” character • more electropositive elements have more “s” character

  16. Bonding in Covalent Oxygen Compounds • Forms coordinate, covalent bonds • can act as a Lewis acid or Lewis base • NF3O • PF3O • H2O

  17. Bonding in Covalent Oxygen Compounds • Oxygen gives “access” to higher oxidation state materials

  18. Oxide Trends • Properties of an oxide depend upon oxidation number of the element • Cr2O3 • m.p. 2266°C • CrO3 • m.p. 196°C

  19. Oxide Trends • Metals in the +2 oxidation state are typically basic MnO(s) + 2H+(aq)  Mn2+(aq) + H2O(l) • Metals in the +3 oxidation state are typically amphoteric Cr2O3(s) + 6H+(aq)  2Cr3+(aq) + 3H2O(l) Cr2O3(s) + 2OH-(aq)  2CrO2-(aq) + H2O(l)

  20. Oxide Trends • Metals in high oxidation states are typically acidic CrO3(s) + H2O(l)  H2CrO4(aq) • Oxides of nonmetals are always covalent, but the higher the oxidation state, the higher the acidity N2O(g) is neutral N2O5(g) + H2O(l)  2HNO3(l)

  21. Oxide Oxidation Numbers • O22- • dioxide(2-) or peroxide • H2O2 • O2- • dioxide(1-) or superoxide • CsO2 • O3- • trioxide(1-) • CsO3

  22. Mixed Metal Oxides • Spinels • Perovskites • ABO3 • A is a large, dipositive metal • B is a small, tetrapositive metal • CaTiO3

  23. Water • Only common liquid on the planet • essential to the chemical reactions of our life

  24. Water • Water has leeched ions from the Earth’s surface throughout the years • Mineral deposits are formed from deposition of aqueous solutions

  25. Water • Ion-dipole interactions account for the solubility of ionic compounds • Water is the basis of our acid-base system 2H2O(l)  H3O+(aq) + OH-(aq)

  26. Hydrogen Peroxide • H2O2 • almost colorless, viscous liquid • extensive hydrogen bonding • very corrosive • thermodynamically unstable to disproportionation H2O2(l)  H2O(l) + 1/2O2(g) G = -119.2 kJ/mol

  27. Hydrogen Peroxide • Prepared by reaction of sodium peroxide with water Na2O2(s) + 2H2O(l) 2NaOH(aq) + H2O2(aq) • Used to restore paintings PbS(s) + 4H2O2(aq) PbSO4(s) + 4H2O(l)

  28. Hydrogen Peroxide • Major industrial chemical • 106 tons produced annually • paper bleaching • household products • hair bleaching • chemical reagent

  29. Hydroxides • Strongest base in aqueous solution • forms compounds with every metallic element • very hazardous • contained in many cleaning products • prepared by the electrolysis of aqueous brine

  30. Hydroxides • Metal hydroxides prepared by mixing a metal salt with sodium hydroxide CuCl2(aq) + 2NaOH(aq) Cu(OH)2(s) + 2NaCl(aq) • Metal hydroxides are generally unstable Cu(OH)2(s) + heat  CuO(s) + H2O(l) • Mixing soluble hydroxides with an acidic oxide gives carbonates 2NaOH(aq) + CO2(g) Na2CO3(aq) + H2O(l)

  31. Hydroxyl Radical

  32. Sulfur • Can be in oxidation states from –2 to +6 • Tends to catenate • Exists as many different allotropic forms • S6 • S8 • S12 • S20

  33. S8 • Most common allotrope • Takes up a “crown” structure • Forms “polymorphs” • two different crystalline forms of the same material

  34. S6 • S6 was the second allotrope discovered • Prepared by mixing sodium thiosulfate with a strong acid 6Na2S2O3(aq) + 12HCl(aq)  S6(s) + 6SO2(g) + 12NaCl(aq) + 6H2O(l)

  35. S12 • Most thermodynamically stable allotrope to S8 • Prepared by reaction of a hydrogen sulfide with a sulfur chloride H2S8(eth.) + S4Cl2(eth.)  S12(s) + 2HCl(g)

  36. Other Sulfur Allotropes S7 S9 S10 S11 S13 S14 S18 S20

  37. Industrial Extraction of Sulfur • Elemental sulfur is found in large, underground deposits in the United States and Poland • around 150 to 750 m beneath the surface • Extracted using the “Frasch process” Herman Frasch (1851-1914)

  38. Frasch Process • Superheated water is pumped down the outermost of three, concentric pipes • Compressed air is pumped down the innermost pipe • A mixture of hot water, air, and molten sulfur comes up the middle pipe

  39. Natural Gas Deposits • Typically contaminated with hydrogen sulfide gas • low levels is called “sweet gas” • high levels is called “sour gas” • Production of sulfur from H2S is called the “Claus process”

  40. Claus Process • Amine Extraction HOCH2CH2NH2(l) + H2S(g)  HOCH2CH2NH3+(solvent) + HS-(solvent) • Thermal Step 3O2(g) + 2H2S(g)  2SO2(g) + 2H2O(g) • Catalytic Step 2SO2(g) + 4H2S(g)  6S(s) + 4H2O(g)

  41. Claus Process

  42. Hydrogen Sulfide • “rotten egg” smell • extremely toxic • can detect at 0.02 ppm • headaches and nausea at 10 ppm • death at 100 ppm • prepared by reaction of a metal sulfide with an acid FeS(s) + 2HCl(aq)  FeCl2(aq) + H2S(g)

  43. Hydrogen Sulfide • Burns in air O2(g) + 2H2S(g)  2S(s) + 2H2O(g) 3O2(g) + 2H2S(g)  2SO2(g) + 2H2O(l) • Used in the production of deuterium oxide H2O(l) + HDS(g)  HDO(l) + H2S(g) HDO(l) + HDS(g)  D2O(l) + H2S(g)

  44. Sulfides • Group 1 and 2 metals plus aluminum form basic, soluble sulfides H2O(l) + S2-(aq)  HS-(aq) + OH-(aq) H2O(l) + HS-(aq)  H2S(g) + OH-(aq) • All other metal sulfides are insoluble Cinnabar Pyrite Stibnite

  45. Industrial Sulfides • Na2S • 105-106 tons annually Na2SO4(s) + 2C(s)  Na2S(l) + 2CO2(g) • tanning of leathers • ore separation • dyes • metal ion precipitation

  46. Laboratory Sulfides • Used in the qualitative testing for metal ions CH3CSNH2(aq) + 2H2O(l)  CH3CO2-(aq) + NH4+(aq) + H2S(g) Cu2+(aq) + S2-(aq)  CuS(s)

  47. Sulfur Oxides • SO2 • colorless, dense, toxic gas • prepared by the addition of sulfite or hydrogen sulfite to an acid SO32-(aq) + 2H+(aq)  H2O(l) + SO2(g) HSO3-(aq) + H+(aq)  H2O(l) + SO2(g) • reducing agent SO2(aq) + 2H2O(l)  SO42-(aq)+ 4H+(aq) + 2e-

  48. Sulfur Oxides • Acid rain problems SO2(g) + OH(g)  HOSO2(g) HOSO2(g)+ O2(g)  HO2(g) + SO3(g) SO3(g) + H2O(g)  H2SO4(aq) • Limestone is used to reduce emissions CaCO3(s) + heat  CaO(s) + CO2(g) 2SO2(g) + 2CaO(s) + O2(g)  2CaSO4(s)

  49. Sulfur Trioxide • SO3 • colorless liquid • very acidic and deliquescent SO3(g) + H2O(l)  H2SO4(l) • prepared from SO2 and oxygen 2SO2(g) + O2(g)  2SO3(g) • forms a trimer upon solidification

  50. Sulfuric Acid • H2SO4 • oily, dense liquid • concentrated solution is 18M • is conductive 2H2SO4(l)  H2O(H2SO4) + H2S2O7(H2SO4) H2O(H2SO4) + 2H2SO4(l)  H3O+(H2SO4) + HSO4-(H2SO4) H2S2O7(H2SO4) + 2H2SO4(l)  H3SO4+(H2SO4) + HS2O7-(H2SO4)

More Related