Alkyl Halides

<body> <p align="center"><font size="7"><strong>Alkyl Halides </strong></font></p> <div align="center"><center><table border="1" width="500"> <tr> <td align="center"><a href="#halide">Alkyl Halides</a></td> <td align="center"><a href="#rxn">Chain Reaction Mechanism</a></td> <td align="center"><a href="#use">Uses of Halogenated Compounds</a></td> </tr> </table> </center></div><p align="center"><a name="halide"><em><strong>Alkyl Halides </strong></em></a></p> <p>Imagine that a pair of crystallizing dishes are placed on an overhead projector as shown in the figure below. An <em>alkene</em> is added to the dish in the upper-left corner of the projector and an <em>alkane</em> is added to the dish in the upper-right corner. A few drops of bromine dissolved in chloroform (CHCl<sub>3</sub>) are then added to each of the crystallizing dishes. </p> <p align="center"><img src="graphics/18.gif" width="446" height="142"> </p> <p>The characteristic red-orange color of bromine disappears the instant this reagent is added to the alkene in the upper-left corner as the Br<sub>2</sub> molecules add across the C=C double bond in the alkene. </p> <p align="center"><img src="graphics/19.gif" width="376" height="50"> </p> <p>The other crystallizing dish picks up the characteristic color of a dilute solution of bromine because this reagent doesn't react with alkanes under normal conditions. </p> <p>If the crystallizing dish in the upper-right corner is moved into the center of the projector, however, the color of the bromine slowly disappears. This can be explained by noting that alkanes react with halogens at high temperatures or in the presence of light to form alkyl halides. </p> <p align="center"><img src="graphics/20.gif" width="389" height="43"></p> <p>The light source in an overhead projector is intense enough to initiate this reaction, although the reaction is still significantly slower than the addition of Br<sub>2</sub> to an alkene. </p> <p>The reaction between an alkane and one of the halogens (F<sub>2</sub>, Cl<sub>2</sub>, Br<sub>2</sub>, or I<sub>2</sub>) can be understood by turning to a simpler example. </p> <p align="center">CH<sub>4</sub>(<em>g</em>) + Cl<sub>2</sub>(<em>g</em>) <img src="graphics/arrowr.gif" width="68" height="12">CH<sub>3</sub>Cl(<em>g</em>) + HCl(<em>g</em>) </p> <p>This reaction has the following characteristic properties. </p> <ul> <li>It doesn't take place in the dark or at low temperatures.</li> <li>It occurs in the presence of ultraviolet light or at temperatures above 250ºC.</li> <li>Once the reaction gets started, it continues after the light is turned off.</li> <li>The products of the reaction include CH<sub>2</sub>Cl<sub>2</sub> (dichloromethane), CHCl<sub>3</sub> (chloroform), and CCl<sub>4</sub> (carbon tetrachloride), as well as CH<sub>3</sub>Cl (chloromethane).</li> <li>The reaction also produces some C<sub>2</sub>H<sub>6</sub>.</li> </ul> <p>These facts are consistent with a <strong>chain-reaction</strong> mechanism that involves three processes: chain initiation, chain propagation, and chain termination. </p> <p align="center"><a href="#top"><img src="graphics/button.gif" alt="return to top" border="0" width="226" height="36"></a></p> <p align="center"><a name="rxn"><em><strong>Chain Reaction Mechanism</strong></em></a></p> <p align="center"><strong>Chain Initiation</strong></p> <p>A Cl<sub>2</sub> molecule can dissociate into a pair of chlorine atoms by absorbing energy in the form of either ultraviolet light or heat. </p> <div align="center"><center><table border="0"> <tr> <td>Cl<sub>2</sub></td> <td><img src="graphics/arrowr.gif" width="68" height="12"></td> <td>2 Cl ·</td> <td> </td> <td> </td> <td><em><img src="graphics/delta.gif" width="11" height="10">H</em><sup>o</sup> = 243.4 kJ/mol<sub>rxn</sub></td> </tr> </table> </center></div><p>The chlorine atom produced in this reaction is an example of a <strong>free radical</strong> <img src="graphics/em.gif" alt="--" width="22" height="9"> an atom or molecule that contains one or more unpaired electrons. </p> <p align="center"><strong>Chain Propagation</strong></p> <p>Free radicals, such as the Cl· atom, are extremely reactive. When a chlorine atom collides with a methane molecule, it can abstract a hydrogen atom to form HCl and a CH<sub>3</sub>· radical. </p> <div align="center"><center><table border="0"> <tr> <td>CH<sub>4</sub> + Cl· </td> <td><img src="graphics/arrowr.gif" width="68" height="12"></td> <td>CH<sub>3</sub>· + HCl</td> <td> </td> <td> </td> <td><em><img src="graphics/delta.gif" width="11" height="10">H</em><sup>o</sup> = -16 kJ/mol<sub>rxn</sub> </td> </tr> </table> </center></div><p>If the CH<sub>3</sub>· radical then collides with a Cl<sub>2</sub> molecule, it can remove a chlorine atom to form CH<sub>3</sub>Cl and a new Cl· radical. </p> <div align="center"><center><table border="0"> <tr> <td>CH<sub>3</sub>· + Cl<sub>2</sub></td> <td><img src="graphics/arrowr.gif" width="68" height="12"></td> <td>CH<sub>3</sub>Cl + Cl· </td> <td> </td> <td> </td> <td><em><img src="graphics/delta.gif" width="11" height="10">H</em><sup>o</sup> = -87 kJ/mol<sub>rxn</sub> </td> </tr> </table> </center></div><p>Because a Cl· atom is generated in the second reaction for every Cl· atom consumed in the first, this reaction continues in a chain-like fashion until the radicals involved in these chain-propagation steps are destroyed. </p> <p align="center"><strong>Chain Termination </strong></p> <p>If a pair of the radicals that keep the chain reaction going collide, they combine in a chain-terminating step. Chain termination can occur in three ways. </p> <div align="center"><center><table border="0"> <tr> <td>2 Cl ·</td> <td><img src="graphics/arrowr.gif" width="68" height="12"></td> <td>Cl<sub>2</sub></td> <td> </td> <td> </td> <td><em><img src="graphics/delta.gif" width="11" height="10">H</em><sup>o</sup> = -243.4 kJ/mol<sub>rxn</sub></td> </tr> <tr> <td>CH<sub>3</sub>· + Cl ·</td> <td><img src="graphics/arrowr.gif" width="68" height="12"></td> <td>CH<sub>3</sub>Cl </td> <td> </td> <td> </td> <td><em><img src="graphics/delta.gif" width="11" height="10">H</em><sup>o</sup> = -330 kJ/mol<sub>rxn</sub> </td> </tr> <tr> <td>2 CH<sub>3</sub>·</td> <td><img src="graphics/arrowr.gif" width="68" height="12"></td> <td>CH<sub>3</sub>CH<sub>3</sub> </td> <td> </td> <td> </td> <td><em><img src="graphics/delta.gif" width="11" height="10">H</em><sup>o</sup> = -350 kJ/mol<sub>rxn</sub> </td> </tr> </table> </center></div><p>Because the concentration of the radicals is relatively small, these chain-termination reactions are relatively infrequent. </p> <p>This chain-reaction mechanism for free-radical reactions explains the observations listed for the reaction between CH<sub>4</sub> and Cl<sub>2</sub>. </p> <ul> <li>The reaction doesn't occur in the dark or at low temperatures because energy must be absorbed to generate the free radicals that carry the reaction.</li> </ul> <div align="center"><center><table border="0"> <tr> <td>Cl<sub>2</sub> </td> <td><img src="graphics/arrowr.gif" width="68" height="12"></td> <td>2 Cl· </td> <td> </td> <td> </td> <td><em><img src="graphics/delta.gif" width="11" height="10">H</em><sup>o</sup> = 243.4 kJ/mol<sub>rxn</sub> </td> </tr> </table> </center></div><ul> <li>The reaction occurs in the presence of ultraviolet light because a UV photon has enough energy to dissociate a Cl<sub>2</sub> molecule to a pair of Cl· atoms. The reaction occurs at high temperatures because Cl<sub>2</sub> molecules can dissociate to form Cl· atoms by absorbing thermal energy.</li> <li>The reaction continues after the light has been turned off because light is only needed to generate the Cl· atoms that start the reaction. The chain reaction then converts CH<sub>4</sub> into CH<sub>3</sub>Cl without consuming these Cl· atoms.</li> </ul> <div align="center"><center><table border="0"> <tr> <td>CH<sub>4</sub> + Cl·</td> <td><img src="graphics/arrowr.gif" width="68" height="12"></td> <td>CH<sub>3</sub>· + HCl </td> </tr> <tr> <td>CH<sub>3</sub>· + Cl<sub>2</sub></td> <td><img src="graphics/arrowr.gif" width="68" height="12"></td> <td>CH<sub>3</sub>Cl + Cl·</td> </tr> </table> </center></div><ul> <li>The reaction doesn't stop at CH<sub>3</sub>Cl because the Cl· atoms can abstract additional hydrogen atoms to form CH<sub>2</sub>Cl<sub>2</sub>, CHCl<sub>3</sub>, and eventually CCl<sub>4</sub>.</li> </ul> <div align="center"><center><table border="0"> <tr> <td>CH<sub>3</sub>Cl + Cl·</td> <td><img src="graphics/arrowr.gif" width="68" height="12"></td> <td>CH<sub>2</sub>Cl· + HCl </td> <td> </td> </tr> <tr> <td>CH<sub>2</sub>Cl· + Cl<sub>2</sub></td> <td><img src="graphics/arrowr.gif" width="68" height="12"></td> <td>CH<sub>2</sub>Cl<sub>2</sub> + Cl·,</td> <td>and so on</td> </tr> </table> </center></div><ul> <li>The formation of C<sub>2</sub>H<sub>6</sub> is a clear indication that the reaction proceeds through a free-radical mechanism. When two CH<sub>3</sub>· radicals collide, they combine to form a ethane molecule.</li> </ul> <div align="center"><center><table border="0"> <tr> <td>2 CH<sub>3</sub>· </td> <td><img src="graphics/arrowr.gif" width="68" height="12"></td> <td>CH<sub>3</sub>CH<sub>3</sub> </td> </tr> </table> </center></div><p>Free-radical halogenation of alkanes provides us with another example of the role of atom-transfer reactions in organic chemistry. The net effect of this reaction is to oxidize a carbon atom by removing a hydrogen from this atom. </p> <div align="center"><center><table border="0"> <tr> <td>CH<sub>4</sub></td> <td>+</td> <td>Cl<sub>2</sub></td> <td><img src="graphics/arrowr.gif" width="68" height="12"></td> <td>CH<sub>3</sub>Cl</td> <td>+</td> <td>HCl</td> </tr> <tr> <td>-4</td> <td> </td> <td> </td> <td> </td> <td>-2</td> <td> </td> <td> </td> </tr> </table> </center></div><p>The reaction, however, doesn't involve the gain or loss of electrons. It occurs by the transfer of a hydrogen atom in one direction and a chlorine atom in the other. </p> <p align="center"><a href="#top"><img src="graphics/button.gif" alt="return to top" border="0" width="226" height="36"></a></p> <p align="center"><a name="use"><em><strong>Uses of Halogenated Compounds</strong></em> </a></p> <p>The chlorinated derivatives of methane have been known for so long that they are frequently referred to by the common names shown in the figure below.</p> <p align="center"><img src="graphics/img31.gif" align="middle" width="464" height="79"> </p> <div align="center"><center><table border="0"> <tr> <td><em>Methyl chloride</em></td> <td> </td> <td> </td> <td><em>Methylene chloride</em></td> <td> </td> <td> </td> <td><em>Chloroform</em></td> <td> </td> <td> </td> <td><em>Carbon tetrachloride </em></td> </tr> <tr> <td align="center"><em>BP</em> = - 24.2ºC</td> <td align="center"> </td> <td align="center"> </td> <td align="center"><em>BP</em> = 40ºC </td> <td align="center"> </td> <td align="center"> </td> <td align="center"><em>BP</em> = 61.7ºC </td> <td align="center"> </td> <td align="center"> </td> <td align="center"><em>BP </em>= 76.5ºC </td> </tr> </table> </center></div><p>The first member of this series is a gas at room temperature; the other three are liquids. These chlorinated hydrocarbons make excellent solvents for the kind of nonpolar solutes that would dissolve in hydrocarbons. They have several advantages over hydrocarbons; they are less volatile and significantly less flammable. </p> <p>Chloroform (CHCl<sub>3</sub>) and carbon tetrachloride (CCl<sub>4</sub>) react with hydrogen fluoride to form a mixture of chlorofluorocarbons, such as CHCl<sub>2</sub>F, CHClF<sub>2</sub>, CCl<sub>3</sub>F, CCl<sub>2</sub>F<sub>2</sub>, and CClF<sub>3</sub>, which are sold under trade names such as Freon and Genetron. The freons are inert gases with high densities, low boiling points, low toxicities, and no odor. As a result, they once found extensive use as propellants in antiperspirants and hair sprays. Controversy over the role of chlorofluorocarbons in the depletion of the Earth's ozone layer led the Environmental Protection Agency to ban the use of CCl<sub>2</sub>F<sub>2</sub> and CCl<sub>3</sub>F in aerosols in 1978. CCl<sub>2</sub>F<sub>2</sub>, CCl<sub>3</sub>F and CHFCl<sub>2</sub> are still used as refrigerants in the air-conditioning industry, however. </p> <p>The structures of a variety of more complex halogenated compounds are shown in the figure below. Halothane is an anesthetic and thyroxine is a thyroid hormone. DDT is an insecticide that has been shown to accumulate in the fatty tissue of birds and is therefore only used as a last resort. Chlordane is a potent insecticide that is still used to control termites. </p> <div align="center"><center><table border="0"> <tr> <td align="center"><img src="graphics/2_4a.gif" width="109" height="96"></td> </tr> <tr> <td align="center"><img src="graphics/2_4b.gif" width="421" height="108"></td> </tr> <tr> <td align="center"><img src="graphics/2_4c.gif" width="261" height="97"></td> </tr> <tr> <td align="center"><img src="graphics/2_4d.gif" width="190" height="134"></td> </tr> </table> </center></div><p align="center"><a href="#top"><img src="graphics/button.gif" alt="return to top" border="0" width="226" height="36"></a></p> <hr> <p><img src="http://chemed.chem.purdue.edu/Count.cgi?df=toprev_org2_alkyl.dat" width="32" height="32"></p> </body>