{"id":829,"date":"2018-11-28T15:37:18","date_gmt":"2018-11-28T15:37:18","guid":{"rendered":"https:\/\/courses.lumenlearning.com\/suny-potsdam-organicchemistry2\/?post_type=chapter&p=829"},"modified":"2023-03-29T05:08:32","modified_gmt":"2023-03-29T05:08:32","slug":"18-4-radical-reactions-in-practice","status":"publish","type":"chapter","link":"https:\/\/courses.lumenlearning.com\/suny-potsdam-organicchemistry2\/chapter\/18-4-radical-reactions-in-practice\/","title":{"raw":"18.4. Radical reactions in practice","rendered":"18.4. Radical reactions in practice"},"content":{"raw":"
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The Free-Radical Chain Reaction<\/span><\/div>\r\n<\/header>
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The three phases of radical chain reactions<\/h2>\r\nThe following video by Leah4Sci provides a good introduction to free radical reactions.\r\n\r\nhttps:\/\/www.youtube.com\/watch?v=Uz1_n9ZnksY<\/a>\r\n\r\nBecause of their high reactivity, free radicals have the potential to be both extremely powerful chemical tools and extremely harmful contaminants.\u00a0 Much of the power of free radical species stems from the natural tendency of radical processes to occur in a chain reaction fashion. Radical chain reactions<\/strong> have three distinct phases: initiation, propagation, and termination.\r\n\r\n\"image020.png\"\r\n\r\nThe initiation phase<\/strong> describes the step that initially creates a radical species. In most cases, this is a homolytic cleavage event, and takes place very rarely due to the high energy barriers involved. Often the influence of heat, UV radiation, or a metal-containing catalyst is necessary to overcome the energy barrier.\r\n\r\nMolecular chlorine and bromine will both undergo homolytic cleavage to form radicals when subjected to heat or light. Other functional groups which also tend to form radicals when exposed to heat or light are chlorofluorocarbons, peroxides, and the halogenated amide N-bromosuccinimide (NBS).\r\n\r\n\"image022.png\"\r\n\r\nThe propagation phase<\/strong> describes the 'chain' part of chain reactions. Once a reactive free radical is generated, it can react with stable molecules to form new free radicals. These new free radicals go on to generate yet more free radicals, and so on. Propagation steps often involve hydrogen abstraction or addition of the radical to double bonds.\r\n\r\n\"image024.png\"\r\n\r\nChain termination<\/strong> occurs when two free radical species react with each other to form a stable, non-radical adduct. Although this is a very thermodynamically downhill event, it is also very rare due to the low concentration of radical species and the small likelihood of two radicals colliding with one another. In other words, the Gibbs free energy barrier is very high for this reaction, mostly due to entropic rather than enthalpic considerations. The active sites of enzymes, of course, can evolve to overcome this entropic barrier by positioning two radical intermediates adjacent to one another.\r\n\r\n<\/div>\r\n
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Radical halogenation in the lab<\/h2>\r\nThe chlorination of an alkane provides a simple example of a free radical chain reaction.\u00a0 In the initiation phase, a chlorine molecule undergoes homolytic cleavage after absorbing energy from light:\r\n\r\n\"image026.png\"\r\n\r\nThe chlorine radical then abstracts a hydrogen, leading to an alkyl radical (step 2), which reacts with a second chlorine molecule (step 3) to form the chloroalkane product plus chlorine radical, which then returns to repeat step 2.\r\n\r\n\"image028.png\"\r\n\r\n\"image030.png\"\r\n\r\nLikely chain termination steps are the condensation of two alkyl radical intermediates or condensation of an alkane radical with a chlorine radical.\r\n\r\nAlkane halogenation reactions exhibit a degree of regiospecificity: if 2-methylbutane is subjected to a limiting amount of chlorine, for example, chlorination takes place fastest at the tertiary carbon.\r\n\r\n\"image032.png\"\r\n\r\nThis is because the tertiary radical intermediate is more stable than the secondary radical intermediate that results from abstraction of the proton on carbon #3, and of course both are more stable than a primary radical intermediate.\u00a0\u00a0 Recall that the Hammond postulate (section 6.2<\/a>, section 15.2B<\/a>) tells us that a lower-energy intermediate implies a lower-energy transition state, and thus a faster reaction.\r\n\r\nUnfortunately, chloroalkanes will readily undergo further chlorination resulting in polychlorinated products, so this is not generally a terribly useful reaction from a synthetic standpoint.\r\n\r\nAlkanes can be brominated by a similar reaction.\u00a0\u00a0 The regiochemical trends are the same as for chlorination, but significantly more pronounced (in other words, bromination is more regioselective).\u00a0 This is because hydrogen abstraction by bromine radical is much less exergonic than by chorine radical \u2013 and this in turn means that the transition state for abstraction by bromine resembles the resulting intermediate more closely than the transition state for abstraction by chlorine resembles its intermediate.\r\n\r\n\"image034.png\"\r\n\r\nAnother way of saying the same thing is that the bromination transition state has more \u2018radical character\u2019 than the chlorination transition state.\u00a0 Trends in radical stability thus have a greater influence on the speed of hydrogen abstraction.\r\n
\r\n\r\nThe reaction proceeds through the radical chain mechanism. The radical chain mechanism is characterized by three steps: initiation<\/strong>, propagation <\/strong>and termination<\/strong>. \u00a0Initiation requires an input of energy but after that the reaction is self-sustaining. \u00a0The first propagation step uses up one of the products from initiation, and the second propagation step makes another one, thus the cycle can continue until indefinitely.\r\n
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Step 1: Initiation<\/h3>\r\nInitiation breaks the bond between the chlorine molecule (Cl2<\/sub>). For this step to occur energy must be put in, this step is not energetically favorable. After this step, the reaction can occur continuously (as long as reactants provide) without input of more energy.\u00a0 It is important to note that this part of the mechanism cannot occur without some external energy input, through light or heat.\r\n\r\n\"\"\r\n\r\n<\/div>\r\n
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Step 2: Propagation<\/h3>\r\nThe next two steps in the mechanism are called propagation steps. In the first propagation step, a chlorine radical combines with a hydrogen on the methane. This gives hydrochloric acid (HCl, the inorganic product of this reaction) and the methyl radical. In the second propagation step more of the chlorine starting material (Cl2<\/sub>) is used, one of the chlorine atoms becomes a radical and the other combines with the methyl radical.\r\n\r\n\"Slide4.jpg\"\r\n\r\nThe first propagation step is endothermic, meaning it takes in heat (requires 2 kcal\/mol) and is not energetically favorable. In contrast the second propagation step is exothermic, releasing 27 kcal\/mol. Since the second propagation step is so exothermic, it occurs very quickly. The second propagation step uses up a product from the first propagation step (the methyl radical) and following Le Chatelier's principle<\/a>, when the product of the first step is removed the equilibrium is shifted towards it's products. This principle is what governs the unfavorable first propagation step's occurance.\r\n\r\n\"\"\r\n\r\n<\/div>\r\n
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Step 3: Termination<\/h3>\r\nIn the termination steps, all the remaining radicals combine (in all possible manners) to form more product (CH3<\/sub>Cl), more reactant (Cl2<\/sub>) and even combinations of the two methyl radicals to form a side product of ethane (CH3<\/sub>CH3<\/sub>).\r\n\r\n\"Slide5.jpg\"\r\n\r\n<\/div>\r\n<\/div>\r\n
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Problems with the chlorination of methane<\/h2>\r\nThe chlorination of methane does not necessarily stop after one chlorination. It may actually be very hard to get a monosubstituted chloromethane.\u00a0 Instead di-, tri- and even tetra-chloromethanes are formed. One way to avoid this problem is to use a much higher concentration of methane in comparison to chloride. This reduces the chance of a chlorine radical running into a chloromethane and starting the mechanism over again to form a dichloromethane. Through this method of controlling product ratios one is able to have a relative amount of control over the product.\r\n\r\n<\/div>\r\n
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Chlorination of other alkanes<\/h2>\r\nWhen alkanes larger than ethane are halogenated, isomeric products are formed. Thus chlorination of propane gives both 1-chloropropane and 2-chloropropane as mono-chlorinated products. Four constitutionally isomeric dichlorinated products are possible, and five constitutional isomers<\/strong> exist for the trichlorinated propanes. Can you write structural formulas for the four dichlorinated isomers?\r\n\r\n\\[CH_3CH_2CH_3 + 2Cl_2 \\rightarrow\u00a0\\text{Four} \\; C_3H_6Cl_2 \\; \\text{isomers} + 2 HCl\\]\r\n\r\nThe halogenation of propane discloses an interesting feature of these reactions. All the hydrogens in a complex alkane do not exhibit equal reactivity<\/strong>. For example, propane has eight hydrogens, six of them being structurally equivalent primary<\/strong>, and the other two being secondary<\/strong>. If all these hydrogen atoms were equally reactive, halogenation should give a 3:1 ratio of 1-halopropane to 2-halopropane mono-halogenated products, reflecting the primary\/secondary numbers. This is not what we observe. Light-induced gas phase chlorination at 25 \u00baC gives 45% 1-chloropropane and 55% 2-chloropropane.\r\n\r\nCH3<\/sub>-CH2<\/sub>-CH3<\/sub> \u00a0 + \u00a0 Cl2<\/sub><\/strong> \u00a0 \u2192\u00a0 45% CH3<\/sub>-CH2<\/sub>-CH2<\/sub>Cl<\/strong> \u00a0 + \u00a0 55% CH3<\/sub>-CHCl<\/strong>-CH3<\/sub>\r\n\r\nThe results of bromination ( light-induced at 25 \u00baC ) are even more surprising, with 2-bromopropane accounting for 97% of the mono-bromo product.\r\n\r\nCH3<\/sub>-CH2<\/sub>-CH3<\/sub> \u00a0 + \u00a0 Br2<\/sub><\/strong>\u00a0 \u2192\u00a0 3% CH3<\/sub>-CH2<\/sub>-CH2<\/sub>Br<\/strong> \u00a0 + \u00a0 97% CH3<\/sub>-CHBr<\/strong>-CH3<\/sub>\r\n\r\nThese results suggest strongly that 2\u00ba-hydrogens are inherently more reactive than 1\u00ba-hydrogens, by a factor of about 3:1. Further experiments showed that 3\u00ba-hydrogens are even more reactive toward halogen atoms. Thus, light-induced chlorination of 2-methylpropane gave predominantly (65%) 2-chloro-2-methylpropane, the substitution product of the sole 3\u00ba-hydrogen, despite the presence of nine 1\u00ba-hydrogens in the molecule.\r\n\r\n(CH3<\/sub>)3<\/sub>CH \u00a0 + \u00a0 Cl2<\/sub><\/strong> \u00a0 \u2192\u00a0 65% (CH3<\/sub>)3<\/sub>CCl<\/strong> \u00a0 + \u00a0 35% (CH3<\/sub>)2<\/sub>CHCH2<\/sub>Cl<\/strong>\r\n\r\nIf you are uncertain about the terms primary (1\u00ba), secondary (2\u00ba) & tertiary (3\u00ba) Click Here<\/a>.\r\n\r\nIt should be clear from a review of the two steps that make up the free radical chain reaction for halogenation that the first step (hydrogen abstraction) is the product determining step<\/strong>. Once a carbon radical is formed, subsequent bonding to a halogen atom (in the second step) can only occur at the radical site. Consequently, an understanding of the preference for substitution at 2\u00ba and 3\u00ba-carbon atoms must come from an analysis of this first step.\r\n\r\n
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Radical Allylic Halogenation<\/h2>\r\nWatch a video by Organic Chemistry Tutor on allylic halogenation:\r\n\r\nhttps:\/\/youtu.be\/EUp2c5aO4dk\r\n\r\n\"\"\r\n\r\n<\/header>
When halogens<\/a> are in the presence of unsaturated<\/a> molecules such as alkenes, the expected reaction is addition<\/a> to the double bond carbons resulting in a vicinal dihalide (halogens on adjacent carbons). However, when the halogen concentration is low enough, alkenes containing allylic hydrogens undergo substitution at the allylic position rather than addition at the double bond. The product is an allylic halide (halogen on carbon next to double bond carbons), which is acquired through a radical chain mechanism.\"Slide2\r\n
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Why substitution of allylic hydrogens?<\/h2>\r\nAs the table below shows, the dissociation energy for the allylic C-H bond is lower than the dissociation energies for the C-H bonds at the vinylic and alkylic positions. This is because the radical formed when the allylic hydrogen is removed is resonance<\/a>-stabilized. Hence,\u00a0given that the halogen concentration is low,\u00a0substitution at the allylic position is favored over competing reactions. However, when the halogen concentration is high, addition at the double bond is favored because a polar reaction outcompetes the radical chain reaction.\r\n\r\n\"Slide8.jpg\"\r\n\r\n<\/div>\r\n
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Radical Allylic Bromination (Wohl-Ziegler Reaction)<\/h2>\r\n
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Preparation of Bromine (low concentration)<\/h3>\r\nNBS (N-bromosuccinimide) is the most commonly used reagent to produce low concentrations of bromine. When suspended in tetrachloride (CCl4<\/sub>), NBS reacts with trace amounts of HBr to produce a low enough concentration of bromine to facilitate the allylic bromination reaction.\r\n\r\n\"Slide1\r\n\r\n<\/div>\r\n
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Allylic Bromination Mechanism<\/h3>\r\n
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Step 1: Initiation<\/h4>\r\nOnce the pre-initiation step involving NBS produces small quantities of Br2<\/sub>, the bromine molecules are homolytically cleaved by light to produce bromine radicals.\r\n\r\n\"Slide3.jpg\"\r\n\r\n<\/div>\r\n
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Step 2: Propagation<\/h4>\r\nOne bromine radical produced by homolytic cleavage in the initiation step removes an allylic hydrogen of the alkene molecule. A radical intermediate is generated, which is stabilized by resonance. The stability provided by delocalization<\/a> of the radical in the alkene intermediate is the reason that substitution at the allylic position is favored over competing reactions such as addition at the double bond.\r\n\r\n\"Slide4\r\n\r\nThe intermediate radical then reacts with a Br2<\/sub> molecule to generate the allylic bromide product and regenerate the bromine radical, which continues the radical chain mechanism. If the alkene reactant is asymmetric, two distinct product isomers are formed.\r\n\r\n\"Slide5\r\n\r\n<\/div>\r\n
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Step 3: Termination<\/h4>\r\nThe radical chain mechanism of allylic bromination can be terminated by any of the possible steps shown below.\r\n\r\n\"Slide6.jpg\"\r\n\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n
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Radical Allylic Chlorination<\/h2>\r\nLike bromination, chlorination at the allylic position of an alkene is achieved when low concentrations of Cl2<\/sub> are present. The reaction is run at high temperatures to achieve the desired results.\r\n
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Industrial Uses<\/h3>\r\nAllylic chlorination has important practical applications in industry. Since chlorine is inexpensive, allylic chlorinations of alkenes have been used in the industrial production of valuable products. For example, 3-chloropropene, which is necessary for the synthesis of products such as epoxy resin, is acquired through radical allylic chlorination (shown below).\r\n\r\n\"Slide7.jpg\"\r\n\r\n<\/div>\r\n<\/div>\r\n<\/section><\/div>\r\n<\/div>\r\n
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Useful polymers formed by radical chain reactions<\/h2>\r\nMany household polymeric materials with which you are probably familiar are made with a radical chain reaction process. Polyethylene (PET), the plastic material used to make soft drink bottles and many other kinds of packaging, is produced by the radical polymerization of ethylene (ethene in IUPAC nomenclature). A radical initiator such as benzoyl peroxide undergoes homolytic cleavage when subjected to high temperatures.\r\n\r\n\"image042.png\"\r\n\r\nIn the propagation phase, the benzoyl radical (X\u2022 in the figure below) adds to the double bond of ethylene, generating a new organic radical.\u00a0 Successive ethylene molecules add to the growing polymer, until termination occurs when two radicals happen to collide. In the figure below, the growing PET polymer is terminated by a benzoyl radical, but in an alternative termination step two growing PET radicals could condense.\r\n\r\n\"image044.png\"\r\n\r\nOther small substituted alkene monomers polymerize in a similar fashion to form familiar polymer materials. Two examples are given below.\r\n\r\n\"image046.png\"\r\n\r\n<\/div>\r\n
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Destruction of the ozone layer by CFC radicals<\/h2>\r\nThe high reactivity of free radicals and the multiplicative nature of radical chain reactions can be useful in the synthesis of materials such as polyethylene plastic - but these same factors can also result in dangerous consequences. You are probably aware of the danger posed to the earth's protective stratospheric ozone layer by the use of chlorofluorocarbons (CFCs) as refrigerants and propellants in aerosol spray cans.\u00a0 Freon-11, or CFCl3<\/sub>, is a typical CFC that was widely used until fairly recently.\u00a0 It can take months or years for a CFC molecule to drift up into the stratosphere from the surface of the earth, and of course the concentration of CFCs at this altitude is very low.\u00a0 Ozone, on the other hand, is continually being formed in the stratosphere.\u00a0 Why all the concern, then, about destruction of the ozone layer - how could such a small amount of CFCs possibly do significant damage?\u00a0 The problem lies in the fact that the process by which ozone is destroyed is a chain reaction, so that a single CFC molecule can initiate the destruction of many ozone molecules before a chain termination event occurs.\r\n\r\n<\/div>\r\n
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Harmful radical species in cells and natural antioxidants<\/h2>\r\nWhile the high reactivity of the hydroxide radical is a beneficial trait in the atmosphere, it is a harmful trait when the same hydroxide radical is present in a living cell.\u00a0 Hydroxide radical and other reactive oxygen species<\/strong> (ROS) such as superoxide (O2<\/sub>-<\/sup>) and peroxide (O2 <\/sub>2-<\/sup>) are continuously produced as minor side-products in the reduction of O2<\/sub> to H2<\/sub>O in respiration.\r\n\r\n\"image052.png\"\r\n\r\nThe ROS are highly reactive oxidizing agents, capable of inflicting damage to DNA, proteins, and the lipids of cell membranes - they are thought to play a major role in the natural aging process.\u00a0 Hydroxide radical, for example, will initiate a radical chain reaction with the hydrocarbon part of an unsaturated\u00a0 membrane lipid molecule that results in the formation of lipid hydroperoxide.\r\n\r\nOne important antioxidant that you are no doubt familiar with is ascorbic acid, or vitamin C.\u00a0 It reacts with harmful radicals to produce the ascorbyl radical, which is significantly more stable than most other radical species due to resonance delocalization.\u00a0 The end result of this first step is that a very reactive, potentially harmful radical (X\u2022) has been 'quenched', and replaced by a much less reactive (and thus less harmful) ascorbyl radical.is thus potentially able to scavenge two<\/em> harmful radical species.\r\n\r\nDehydroascorbate is subsequently either broken down and excreted, or else recycled (reduced) back to ascorbate.\u00a0 This can happen either in a direct, enzyme-free reaction with glutathione, or through the action of a specific glutathione\/NADH-dependant reductase enzyme.\u00a0 You were invited to propose a likely mechanism for the enzyme-free reaction in problem 16.14.\r\n
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Contributors<\/h2>\r\n