In chemistry, a radical (more precisely, a free radical ) is an atom, molecule, or ion that has unpaired valence electrons or an open electron shell, and therefore may be seen as having one or more "dangling" covalent bonds. With some exceptions, these "dangling" bonds make free radicals highly chemically reactive towards other substances, or even towards themselves: their molecules will often spontaneously dimerize or polymerize if they come in contact with each other. Most radicals are reasonably stable only at very low concentrations in inert media or in a vacuum. Free radicals may be created in a number of ways, including synthesis with very dilute or rarefied reagents, reactions at very low temperatures, or breakup of larger molecules. The latter can be affected by any process that puts enough energy into the parent molecule, such as ionizing radiation, heat, electrical discharges, electrolysis, and chemical reactions. Indeed, radicals are intermediate stages in many chemical reactions. History The first organic free radical identified was triphenylmethyl radical. This species was discovered by Moses Gomberg in 1900 at the University of Michigan USA. Historically, the term radical in radical theory was also used for bound parts of the molecule, especially when they remain unchanged in reactions. These are now called functional groups. For example, methyl alcohol was described as consisting of a methyl "radical" and a hydroxyl "radical". Neither are radicals in the modern chemical sense, as they are permanently bound to each other, and have no unpaired, reactive electrons; however, they can be observed as radicals in mass spectrometry when broken apart by irradiation with energetic electrons. Depiction in chemical reactions In this chapter, we will learn about some reactions in which the key steps involve the movement of single electrons. You may recall from way back in section 6.1A that single electron movement is depicted by a single-barbed 'fish-hook ' arrow (as opposed to the familiar double-barbed arrows that we have been using throughout the book to show two-electron movement). Single-electron mechanisms involve the formation and subsequent reaction of free radical species, highly unstable intermediates that contain an unpaired electron. We will learn in this chapter how free radicals are often formed from homolytic cleavage , an event where the two electrons in a breaking covalent bond move in opposite directions. (In contrast, essentially all of the reactions we have studied up to now involve bond-breaking events in which both electrons move in the same direction: this is called heterolytic cleavage ). We will also learn that many single-electron mechanisms take the form of a radical chain reaction, in which one radical causes the formation of a second radical, which in turn causes the formation of a third radical, and so on. The high reactivity of free radical species and their ability to initiate chain reactions is often beneficial - we will learn in this chapter about radical polymerization reactions that form useful materials such as plexiglass and polyproylene fabric. We will also learn about radical reactions that are harmful, such as the degradation of atmospheric ozone by freon, and the oxidative damage done to lipids and DNA in our bodies by free radicals species. Finally, we will see how some enzymes use bound metals to catalyze high e The geometry and relative stability of carbon radicals As organic chemists, we are particularly interested in radical intermediates in which the unpaired electron resides on a carbon atom. Experimental evidence indicates that the three bonds in a carbon radical have trigonal planar geometry, and therefore the carbon is considered to be sp 2 -hybridized with the unpaired electron occupying the perpendicular, unhybridized 2p z orbital. Contrast this picture with carbocation and carbanion intermediates, which are both also trigonal planar but whose 2p z orbitals contain zero or two electrons, respectively. The trend in the stability of carbon radicals parallels that of carbocations (section 8.4B ): tertiary radicals, for example, are more stable than secondary radicals, followed by primary and methyl radicals. This should make intuitive sense, because radicals, like carbocations, can be considered to be electron deficient, and thus are stabilized by the electron-donating effects of nearby alkyl groups. Benzylic and allylic radicals are more stable than alkyl radicals due to resonance effects - an unpaired electron can be delocalized over a system of conjugated pi bonds. An allylic radical, for example, can be pictured as a system of three parallel 2p z orbitals sharing three electrons. Trends in radical stability Allyic & Benzlic > 3 o > 2 o > 1 o > Methyl Template:ExampleStart Exercise 17.1: Draw the structure of a benzylic radical compound, and then draw a resonance form showing how the radical is stabilized. Template:ExampleEnd With enough resonance stabilization, radicals can be made that are quite unreactive. One example of an inert organic radical structure is shown below. In this molecule, the already extensive resonance stabilization is further enhanced by the ability of the chlorine atoms to shield the radical center from external reagents. The radical is, in some sense, inside a protective 'cage'. Organic Chemistry With a Biological Emphasis by Tim Soderberg (University of Minnesota, Morris) Prof. Steven Farmer ( Sonoma State University ) The Relationship between Bond Order and Bond Energy Triple bonds between like atoms are shorter than double bonds, and because more energy is required to completely break all three bonds than to completely break two, a triple bond is also stronger than a double bond. Similarly, double bonds between like atoms are stronger and shorter than single bonds. Bonds of the same order between different atoms show a wide range of bond energies, however. Table 8.6 lists the average values for some commonly encountered bonds. Although the values shown vary widely, we can observe four trends: Table 8.6 Average Bond Energies (kJ/mol) for Commonly Encountered Bonds at 273 K Single Bonds Single Bonds.1 Single Bonds.2 Single Bonds.3 Single Bonds.4 Single Bonds.5 Single Bonds.6 Single Bonds.7 Single Bonds.8 Single Bonds.9 Multiple Bonds Multiple Bonds.1 H–H 432 C–C 346 N–N ≈167 O–O ≈142 F–F 155.0 C=C 602 H–C 411 C–Si 318 N–O 201 O–F 190 F–Cl 249.0 C≡C 835 H–Si 318 C–N 305 N–F 283 O–Cl 218 F–Br 249.0 C=N 615 H–N 386 C–O 358 N–Cl 313 O–Br 201 F–I 278.0 C≡N 887 H–P ≈322 C–S 272 N–Br 243 O–I 201 Cl–Cl 240.0 C=O 749 H–O 459 C–F 485 P–P 201 S–S 226 Cl–Br 216.0 C≡O 1072 H–S 363 C–Cl 327 NaN NaN S–F 284 Cl–I 208.0 N=N 418 H–F 565 C–Br 285 NaN NaN S–Cl 255 Br–Br 190.0 N≡N 942 H–Cl 428 C–I 213 NaN NaN S–Br 218 Br–I 175.0 N=O 607 H–Br 362 Si–Si 222 NaN NaN NaN NaN I–I 149.0 O=O 494 H–I 295 Si–O 452 NaN NaN NaN NaN NaN NaN S=O 532 Source: Data from J. E. Huheey, E. A. Keiter, and R. L. Keiter, Inorganic Chemistry , 4th ed. (1993). Bonds between hydrogen and atoms in the same column of the periodic table decrease in strength as we go down the column. Thus an H–F bond is stronger than an H–I bond, H–C is stronger than H–Si, H–N is stronger than H–P, H–O is stronger than H–S, and so forth. The reason for this is that the region of space in which electrons are shared between two atoms becomes proportionally smaller as one of the atoms becomes larger (part (a) in Figure 8.11). Bonds between like atoms usually become weaker as we go down a column (important exceptions are noted later). For example, the C–C single bond is stronger than the Si–Si single bond, which is stronger than the Ge–Ge bond, and so forth. As two bonded atoms become larger, the region between them occupied by bonding electrons becomes proportionally smaller, as illustrated in part (b) in Figure 8.11. Noteworthy exceptions are single bonds between the period 2 atoms of groups 15, 16, and 17 (i.e., N, O, F), which are unusually weak compared with single bonds between their larger congeners. It is likely that the N–N, O–O, and F–F single bonds are weaker than might be expected due to strong repulsive interactions between lone pairs of electrons on adjacent atoms. The trend in bond energies for the halogens is therefore Cl–Cl > Br–Br > F–F > I–I Similar effects are also seen for the O–O versus S–S and for N–N versus P–P single bonds. Note Bonds between hydrogen and atoms in a given column in the periodic table are weaker down the column; bonds between like atoms usually become weaker down a column. 3. Because elements in periods 3 and 4 rarely form multiple bonds with themselves, their multiple bond energies are not accurately known. Nonetheless, they are presumed to be significantly weaker than multiple bonds between lighter atoms of the same families. Compounds containing an Si=Si double bond, for example, have only recently been prepared, whereas compounds containing C=C double bonds are one of the best-studied and most important classes of organic compounds. Figure 8.11 The Strength of Covalent Bonds Depends on the Overlap between the Valence Orbitals of the Bonded Atoms. The relative sizes of the region of space in which electrons are shared between (a) a hydrogen atom and lighter (smaller) vs. heavier (larger) atoms in the same periodic group; and (b) two lighter versus two heavier atoms in the same group. Although the absolute amount of shared space increases in both cases on going from a light to a heavy atom, the amount of space relative to the size of the bonded atom decreases; that is, the percentage of total orbital volume decreases with increasing size. Hence the strength of the bond decreases. 4. Multiple bonds between carbon, oxygen, or nitrogen and a period 3 element such as phosphorus or sulfur tend to be unusually strong. In fact, multiple bonds of this type dominate the chemistry of the period 3 elements of groups 15 and 16. Multiple bonds to phosphorus or sulfur occur as a result of d -orbital interactions, as we discussed for the SO 4 2− ion in Section 8.6. In contrast, silicon in group 14 has little tendency to form discrete silicon–oxygen double bonds. Consequently, SiO 2 has a three-dimensional network structure in which each silicon atom forms four Si–O single bonds, which makes the physical and chemical properties of SiO 2 very different from those of CO 2 . Note Bond strengths increase as bond order increases, while bond distances decrease. Table : Average bond energies: 0 1 Bond (kJ/mol) C-F 485 C-Cl 328 C-Br 276 C-I 240 NaN NaN C-C 348 C-N 293 C-O 358 C-F 485 NaN NaN C-C 348 C=C 614 C=C 839 Contributors Kim Song (UCD), Donald Le (UCD)