Presentation on theme: "The Elements: The Last Four Main Groups The firing of the powerful space shuttle booster rockets is an impressive display of the vigor of the reactions."— Presentation transcript:
The Elements: The Last Four Main Groups The firing of the powerful space shuttle booster rockets is an impressive display of the vigor of the reactions that take place between some of the elements on the right of the periodic table. These elements play an important role in many aspects of space flight, not only as fuel. Their reactions are also important in maintaining life on our planet.
Figure 20.1 The elements of Group 15: (back row, from left to right) nitrogen (cooled to the liquid), red phosphorus, arsenic; (front row, from left to right) antimony and bismuth.
Figure 20.2 The bacteria that inhabit these nodules on the roots of a pea plant are able to fix atmospheric nitrogen and make it available to the plant.
Figure 20.3 The minerals (from left to right) orpiment, As 2 S 3 ; stibnite, Sb 2 S 3 ; and realgar, As 4 S 4.
Figure 20.4 When aqueous ammonia is added to a copper(II) sulfate solution, first a light-blue precipitate of Cu(OH) 2 forms. The precipitate disappears when more ammonia is added to form the dark blue complex [Cu(NH 3 ) 4 ] 2 by a Lewis acid-base reaction.
Figure 20.5 Magnesium nitride is formed when magnesium burns in an atmosphere of nitrogen.
Figure 20.6 When sodium azide (left) is heated in a sealed, evacuated tube, it decomposes into sodium metal and nitrogen gas. The sodium condenses on the walls of the tube (right).
Figure 20.7 Dinitrogen trioxide, N 2 O 3, condenses to a deep blue liquid that freezes at 100°C to a pale blue solid. On standing, it turns green as a result of partial decomposition into nitrogen dioxide (not shown), a yellow-brown gas.
Figure 20.8 Liquid oxygen is pale blue (the gas itself is colorless).
Figure 20.9 Ozone is a blue gas that condenses to a dark blue, highly unstable liquid.
Figure A collection of sulfide ores (from left to right): galena, PbS; cinnabar, HgS; pyrite, FeS 2 ; sphalerite, ZnS. Pyrite has a lustrous golden color and has frequently been mistaken for gold; hence it is also known as fool’s gold. Gold and fool’s gold are readily distinguished by their densities.
Figure One of the two most common forms of sulfur is the blocklike rhombic form (a). It differs from the needlelike monoclinic sulfur (b) in the manner in which the S 8 rings are stacked together.
Figure Two of the Group 16 elements: selenium (left) and tellurium (right).
Figure The electronegativities of the Group 16 elements decrease down the group.
Figure The atomic and ionic radii of the Group 16 elements increase down the group. The values shown are in picometers, and the anion (shown in green in each case) is substantially larger than the neutral parent atom.
Figure Chemiluminescence, the emission of light as the result of a chemical reaction, occurs when hydrogen peroxide is added to a solution of the organic compound perylene. Although hydrogen peroxide itself can fluoresce, in this case the light is emitted by the perylene.
Figure The blue stones in this ancient Egyptian ornament are lapis lazuli. This semiprecious stone is an aluminosilicate colored by S 2 and S 3 impurities. The blue color is due to S 3 , and its hint of green to S 2 .
Figure Sulfur dioxide is a reducing agent. When it is bubbled through an aqueous solution of bromine (left), it reduces the Br2 to colorless bromide ions (right). The SO 2 is oxidized to H 2 SO 4.
Figure Sulfuric acid is an oxidizing agent. When concentrated acid is poured onto solid sodium bromide, NaBr, the bromide ions are oxidized to bromine, which colors the solution red-brown.
Figure Sulfuric acid is a dehydrating agent. (a) When concentrated sulfuric acid is poured onto sucrose, (b) the sucrose, a carbohydrate, is dehydrated, (c) leaving a frothy black mass of carbon.
Figure The electronegativities of the halogens decrease steadily down the group.
Figure The atomic and ionic radii of the halogens increase steadily down the group as electrons occupy outer shells of the atoms further from the nuclei. The values shown are in picometers. In all cases, ionic radii (represented by the green spheres) are larger than atomic radii.
Figure Fluorine is prepared on a large scale by an adaptation of the electrolytic method that was used to isolate it originally. These electrolytic cells are producing fluorine in a commercial preparation plant.
Figure Iron reacts vigorously and exothermically with chlorine to form anhydrous iron(III) chloride.
Figure Chlorine is an oxidizing agent. When chlorine is bubbled through a colorless solution of bromide ions, it oxidizes them to bromine, which colors the solution red-brown.
Figure Solutions of iodine in a variety of solvents. From left to right, the solvents are tetrachloromethane (carbon tetrachloride), water, and potassium iodide solution. In the solution at the far right, a little starch has been added to a solution of iodine in potassium iodide solution; starch acts as an indicator for the presence of iodine.
Figure When a mixture of hydrofluoric acid and ammonium fluoride is swirled inside a flask (a), the reaction with the silica in the glass frosts the glass surface (b).
Case Study 20 (a) Powdered reactants are mixed with a liquid polymer base and hardened inside the space shuttle booster rocket shell.
Case Study 20 (b) The white smoke emitted by the space shuttle booster rockets consists of powdered aluminum oxide and aluminum chloride.
Case Study 20 (c) The Apollo lunar lander was powered by a mixture of hydrazine derivatives and dinitrogen tetroxide.
Figure The colors of fluorescent lighting art by Tom Anthony are due to emissions from noble-gas atoms. Neon is responsible for the red light; when it is mixed with a little argon, the color becomes blue-green. The yellow color is achieved by coating the inside of the glass with substances that give off yellow light when excited.
Figure The ionization energies of the noble gases decrease steadily down the group. The values shown are in kilojoules per mole.
Figure Crystals of xenon tetrafluoride, XeF 4. This compound was first prepared in 1962 by the reaction of xenon and fluorine at 6 atm and 400°C.