The Carbonyl Group

<body> <p align="center"><font size="6"><strong>The Carbonyl Group</strong></font></p> <div align="center"><center><table border="1" width="500"> <tr> <td align="center"><a href="#describe">Description of the Carbonyl Group</a></td> <td align="center"><a href="#react">Reactions at the Carbonyl Group</a> </td> <td align="center" colspan="2"><a href="#acid">Carboxylic Acids and Carboxylate Ions </a></td> </tr> <tr> <td align="center"><a href="#name">Naming Carboxylic Acids</a></td> <td align="center"><a href="#dicarbox">Dicarboxylic and Tricarboxylic Acids</a></td> <td align="center"><a href="#esters">Esters</a></td> <td align="center"><a href="#fat">Fats and Oils </a></td> </tr> </table> </center></div><p align="center"><a name="describe"><em><strong>Description of the Carbonyl Group</strong></em></a></p> <p>It is somewhat misleading to write the carbonyl group as a covalent C=O double bond. The difference between the electronegativities of carbon and oxygen is large enough to make the C=O bond moderately polar. As a result, the carbonyl group is best described as a hybrid of the following resonance structures. </p> <p align="center"><img src="graphics/40.gif" width="205" height="82"> </p> <p>We can represent the polar nature of this hybrid by indicating the presence of a slight negative charge on the oxygen (<font face="Symbol">d</font><sup>-</sup>) and a slight positive charge (<font face="Symbol">d</font><sup>+</sup>) on the carbon of the C=O double bond. </p> <p align="center"><img src="graphics/41.gif" width="52" height="94"> </p> <p align="center"><a name="react"><em><strong>Reactions at the Carbonyl Group </strong></em></a></p> <p align="left">Reagents that attack the electron-rich <font face="Symbol">d</font> <sup>-</sup> end of the C=O bond are called <strong>electrophiles</strong> (literally, "lovers of electrons"). Electrophiles include ions (such as H<sup>+</sup> and Fe<sup>3+</sup>) and neutral molecules (such as AlCl<sub>3</sub> and BF<sub>3</sub>) that are <em>Lewis acids</em>, or <em>electron-pair acceptors</em>. Reagents that attack the electron-poor <font face="Symbol">d</font><sup>+</sup> end of this bond are <strong>nucleophiles</strong> (literally, "lovers of nuclei"). Nucleophiles are <em>Lewis bases</em> (such as NH<sub>3</sub> or the OH<sup>-</sup> ion). </p> <p align="center"><img src="graphics/42.gif" width="172" height="134"> </p> <p>The polarity of the C=O double bond can be used to explain the reactions of carbonyl compounds. Aldehydes and ketones react with a source of the hydride (H<sup>-</sup>) ion because the H<sup>-</sup> ion is a Lewis base, or nucleophile, that attacks the <font face="Symbol">d</font><sup>+</sup> end of the C=O bond. When this happens, the two valence electrons on the H<sup>-</sup> ion form a covalent bond to the carbon atom. Since carbon is tetravalent, one pair of electrons in the C=O bond is displaced onto the oxygen to form an intermediate with a negative charge on the oxygen atom. </p> <p align="center"><img src="graphics/43.gif" width="257" height="117"> </p> <p>This alkoxide ion can then remove an H<sup>+</sup> ion from water to form an alcohol. </p> <p align="center"><img src="graphics/44.gif" width="433" height="96"> </p> <p>Common sources of the H<sup>-</sup> ion include lithium aluminum hydride (LiAlH<sub>4</sub>) and sodium borohydride (NaBH<sub>4</sub>). Both compounds are ionic. </p> <div align="center"><center><table border="0"> <tr> <td>LiAlH<sub>4</sub>:</td> <td> </td> <td> </td> <td>[Li<sup>+</sup>][AlH<sub>4</sub><sup>-</sup>] </td> </tr> <tr> <td> </td> <td> </td> <td> </td> <td> </td> </tr> <tr> <td>NaBH<sub>4</sub>: </td> <td> </td> <td> </td> <td>[Na<sup>+</sup>][BH<sub>4</sub><sup>-</sup>] </td> </tr> </table> </center></div><p>The aluminum hydride (AlH<sub>4</sub><sup>-</sup>) and borohydride (BH<sub>4</sub><sup>-</sup>) ions act as if they were complexes between an H<sup>-</sup> ion, acting as a Lewis base, and neutral AlH<sub>3</sub> or BH<sub>3</sub> molecules, acting as a Lewis acid. </p> <p align="center"><img src="graphics/45.gif" width="282" height="87"></p> <p>LiAlH<sub>4</sub> is such as good source of the H<sup>-</sup> ion that it reacts with the H<sup>+</sup> ions in water or other protic solvents to form H<sub>2</sub> gas. The first step in the reduction of a carbonyl with LiAlH<sub>4</sub> is therefore carried out using an ether as the solvent. The product of the hydride reduction reaction is then allowed to react with water in a second step to form the corresponding alcohol. </p> <p align="center"><img src="graphics/img62.gif" align="middle" width="255" height="67"> </p> <p>NaBH<sub>4</sub> is less reactive toward protic solvents, which means that borohydride reductions are usually done in a single step, using an alcohol as the solvent. </p> <p align="center"><img src="graphics/img63.gif" align="middle" width="263" height="61"> </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="acid"><em><strong>Carboxylic Acids and Carboxylate Ions </strong></em></a></p> <p>When one of the substituents on a carbonyl group is an OH group, the compound is a <strong>carboxylic acid</strong> with the generic formula RCO<sub>2</sub>H. These compounds are acids, as the name suggests, which form <strong>carboxylate ions</strong> (RCO<sub>2</sub><sup>-</sup>) by the loss of an H<sup>+</sup> ion. </p> <p align="center"><img src="graphics/46.gif" width="317" height="57"> </p> <p>The carboxylate ion formed in this reaction is a hybrid of two resonance structures. </p> <p align="center"><img src="graphics/47.gif" width="297" height="90"> </p> <p align="left">Resonance delocalizes the negative charge in the carboxylate ion, which makes this ion more stable than the alkoxide ion formed when an alcohol loses an H<sup>+</sup> ion. By increasing the stability of the conjugate base, resonance increases the acidity of the acid that forms this base. Carboxylic acids are therefore much stronger acids than the analogous alcohols. The value of <em>K</em><sub><em>a</em></sub> for a typical carboxylic acid is about 10<sup>-5</sup>, whereas alcohols have values of <em>K</em><sub><em>a</em></sub> of only 10<sup>-16</sup>. </p> <p align="center"><img src="graphics/48.gif" width="448" height="106"> </p> <p>Carboxylic acids were among the first organic compounds to be discovered. As a result, they have well-established common names that are often derived from the Latin stems of their sources in nature. Formic acid (Latin <em>formica</em>, "an ant") and acetic acid (Latin <em>acetum</em>, "vinegar") were first obtained by distilling ants and vinegar, respectively. Butyric acid (Latin <em>butyrum</em>, "butter") is found in rancid butter, and caproic, caprylic, and capric acids (Latin <em>caper</em>, "goat") are all obtained from goat fat. A list of common carboxylic acids is given in the table below. </p> <p align="center"><a name="soltable"></a><em>Common Carboxylic Acids </em></p> <div align="center"><center><table border="0"> <tr> <td> </td> <td> </td> <td><em>Common Name</em></td> <td> </td> <td><em>Formula</em></td> <td> </td> <td colspan="8"><em>Solubility in H</em><sub><em>2</em></sub><em>O<br> (g/100 mL)</em> </td> </tr> <tr> <td colspan="3">Saturated carboxylic acids <br> and fatty acids</td> <td> </td> <td> </td> <td> </td> <td colspan="8"> </td> </tr> <tr> <td> </td> <td> </td> <td>Formic acid</td> <td> </td> <td>HCO<sub>2</sub>H </td> <td> </td> <td> </td> <td><img src="graphics/infinity.gif" width="16" height="7"></td> <td> </td> <td> </td> <td> </td> <td> </td> <td> </td> <td> </td> </tr> <tr> <td> </td> <td> </td> <td>Acetic acid </td> <td> </td> <td>CH<sub>3</sub>CO<sub>2</sub>H </td> <td> </td> <td> </td> <td><img src="graphics/infinity.gif" width="16" height="7"></td> <td> </td> <td> </td> <td> </td> <td> </td> <td> </td> <td> </td> </tr> <tr> <td> </td> <td> </td> <td>Proprionic acid </td> <td> </td> <td>CH<sub>3</sub>CH<sub>2</sub>CO<sub>2</sub>H </td> <td> </td> <td> </td> <td><img src="graphics/infinity.gif" width="16" height="7"></td> <td> </td> <td> </td> <td> </td> <td> </td> <td> </td> <td> </td> </tr> <tr> <td> </td> <td> </td> <td>Butyric acid </td> <td> </td> <td>CH<sub>3</sub>(CH<sub>2</sub>)<sub>2</sub>CO<sub>2</sub>H</td> <td> </td> <td> </td> <td><img src="graphics/infinity.gif" width="16" height="7"></td> <td> </td> <td> </td> <td> </td> <td> </td> <td> </td> <td> </td> </tr> <tr> <td> </td> <td> </td> <td>Caproic acid </td> <td> </td> <td>CH<sub>3</sub>(CH<sub>2</sub>)<sub>4</sub>CO<sub>2</sub>H </td> <td> </td> <td colspan="8">0.968 </td> </tr> <tr> <td> </td> <td> </td> <td>Caprylic acid </td> <td> </td> <td>CH<sub>3</sub>(CH<sub>2</sub>)<sub>6</sub>CO<sub>2</sub>H </td> <td> </td> <td colspan="8">0.068</td> </tr> <tr> <td> </td> <td> </td> <td>Capric acid </td> <td> </td> <td>CH<sub>3</sub>(CH<sub>2</sub>)<sub>8</sub>CO<sub>2</sub>H </td> <td> </td> <td colspan="8">0.015 </td> </tr> <tr> <td> </td> <td> </td> <td>Lauric acid </td> <td> </td> <td>CH<sub>3</sub>(CH<sub>2</sub>)<sub>10</sub>CO<sub>2</sub>H </td> <td> </td> <td colspan="8">0.0055 </td> </tr> <tr> <td> </td> <td> </td> <td>Myristic acid </td> <td> </td> <td>CH<sub>3</sub>(CH<sub>2</sub>)<sub>12</sub>CO<sub>2</sub>H </td> <td> </td> <td colspan="8">0.0020</td> </tr> <tr> <td> </td> <td> </td> <td>Palmitic acid </td> <td> </td> <td>CH<sub>3</sub>(CH<sub>2</sub>)<sub>14</sub>CO<sub>2</sub>H </td> <td> </td> <td colspan="8">0.00072</td> </tr> <tr> <td> </td> <td> </td> <td>Stearic acid </td> <td> </td> <td>CH<sub>3</sub>(CH<sub>2</sub>)<sub>16</sub>CO<sub>2</sub>H </td> <td> </td> <td colspan="8">0.00029 </td> </tr> <tr> <td colspan="3">Unsaturated fatty acids </td> <td> </td> <td> </td> <td> </td> <td colspan="8"> </td> </tr> <tr> <td> </td> <td> </td> <td>Palmitoleic acid </td> <td> </td> <td colspan="10">CH<sub>3</sub>(CH<sub>2</sub>)<sub>5</sub>CH=CH(CH<sub>2</sub>)<sub>7</sub>CO<sub>2</sub>H </td> </tr> <tr> <td> </td> <td> </td> <td>Oleic acid </td> <td> </td> <td colspan="10">CH<sub>3</sub>(CH<sub>2</sub>)<sub>7</sub>CH=CH(CH<sub>2</sub>)<sub>7</sub>CO<sub>2</sub>H </td> </tr> <tr> <td> </td> <td> </td> <td>Linoleic acid </td> <td> </td> <td colspan="10">CH<sub>3</sub>(CH<sub>2</sub>)<sub>4</sub>CH=CHCH<sub>2</sub>CH=CH(CH<sub>2</sub>)<sub>7</sub>CO<sub>2</sub>H</td> </tr> <tr> <td> </td> <td> </td> <td>Linolenic acid </td> <td> </td> <td colspan="10">CH<sub>3</sub>CH<sub>2</sub>CH=CHCH<sub>2</sub>CH=CHCH<sub>2</sub>CH=CH(CH<sub>2</sub>)<sub>7</sub>CO<sub>2</sub>H </td> </tr> </table> </center></div><p align="center"><a name="name"><em><strong>Naming Carboxylic Acids</strong></em></a></p> <p>The systematic nomenclature of carboxylic acids is easy to understand. The ending -<em>oic acid</em> is added to the name of the parent alkane to indicate the presence of the <img src="graphics/em.gif" alt="--" width="22" height="9">CO<sub>2</sub>H functional group. </p> <div align="center"><center><table border="0"> <tr> <td>HCO<sub>2</sub>H</td> <td> </td> <td> </td> <td><em>Methanoic acid</em> </td> </tr> <tr> <td>CH<sub>3</sub>CO<sub>2</sub>H</td> <td> </td> <td> </td> <td><em>Ethanoic acid</em> </td> </tr> <tr> <td>CH<sub>3</sub>CH<sub>2</sub>CO<sub>2</sub>H</td> <td> </td> <td> </td> <td><em>Propanoic acid</em> </td> </tr> </table> </center></div><p>Unfortunately, because of the long history of their importance in biology and biochemistry, you are more likely to encounter these compounds by their common names. </p> <p>Formic acid and acetic acid have a sharp, pungent odor. As the length of the alkyl chain increases, the odor of carboxylic acids becomes more unpleasant. Butyric acid, for example, is found in sweat, and the odor of rancid meat is due to carboxylic acids released as the meat spoils. </p> <p>The <a href="#soltable">solubility data</a> in the table above show that carboxylic acids also become less soluble in water as the length of the alkyl chain increases. The <img src="graphics/em.gif" alt="--" width="22" height="9">CO<sub>2</sub>H end of this molecule is polar and therefore soluble in water. As the alkyl chain gets longer, the molecule becomes more nonpolar and less soluble in water. </p> <p align="center"><a name="dicarbox"><em><strong>Dicarboxylic and Tricarboxylic Acids</strong></em></a></p> <p>Compounds that contain two <img src="graphics/em.gif" alt="--" width="22" height="9">CO<sub>2</sub>H functional groups are known as<strong> dicarboxylic acids</strong>. A number of dicarboxylic acids (see table below) can be isolated from natural sources. Tartaric acid, for example, is a by-product of the fermentation of wine, and succinic, fumaric, malic, and oxaloacetic acid are intermediates in the metabolic pathway used to oxidize sugars to CO<sub>2</sub> and H<sub>2</sub>O. </p> <p align="center"><em>Common Dicarboxylic Acids </em></p> <div align="center"><center><table border="0"> <tr> <td><img src="graphics/49.gif" width="73" height="17"></td> <td> </td> <td valign="bottom">oxalic acid</td> </tr> <tr> <td><img src="graphics/50.gif" width="98" height="17"></td> <td> </td> <td valign="bottom">malonic acid</td> </tr> <tr> <td><img src="graphics/52.gif" width="117" height="53"></td> <td> </td> <td valign="bottom">malic acid</td> </tr> <tr> <td><img src="graphics/51.gif" width="123" height="17"></td> <td> </td> <td valign="bottom">succinic acid</td> </tr> <tr> <td><img src="graphics/55.gif" width="136" height="56"></td> <td> </td> <td valign="bottom">maleic acid</td> </tr> <tr> <td><img src="graphics/56.gif" width="136" height="59"></td> <td> </td> <td valign="bottom">fumaric acid</td> </tr> <tr> <td><img src="graphics/53.gif" width="112" height="52"></td> <td> </td> <td valign="bottom">tartaric acid</td> </tr> <tr> <td><img src="graphics/54.gif" width="107" height="53"></td> <td> </td> <td valign="bottom">oxaloacetic acid</td> </tr> </table> </center></div><p>Several tricarboxylic acids also play an important role in the metabolism of sugar. The most important example of this class of compounds is the citric acid that gives so many fruit juices their characteristic acidity. </p> <p align="center"><img src="graphics/57.gif" width="195" height="77"> </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="esters"><em><strong>Esters </strong></em></a></p> <p>Carboxylic acids (-CO<sub>2</sub>H) can react with alcohols (ROH) in the presence of either acid or base to form <strong>esters</strong> (-CO<sub>2</sub>R). Acetic acid, for example, reacts with ethanol to form ethyl acetate and water. </p> <p align="center"><img src="graphics/58.gif" width="319" height="51"> </p> <p>This isn't an efficient way of preparing an ester, however, because the equilibrium constant for this reaction is relatively small (<em>K</em><sub>c</sub> <img src="graphics/congrnt.gif" width="20" height="11">3). Chemists tend to synthesize esters in a two-step process. They start by reacting the acid with a chlorinating agent such as thionyl chloride (SOCl<sub>2</sub>) to form the corresponding <strong>acyl chloride</strong>. </p> <p align="center"><img src="graphics/59.gif" width="374" height="59"> </p> <p>They then react the acyl chloride with an alcohol in the presence of base to form the ester. </p> <p align="center"><img src="graphics/60.gif" width="397" height="53"> </p> <p>The base absorbs the HCl given off in this reaction, thereby driving it to completion. </p> <p>As might be expected, esters are named as if they were derivatives of a carboxylic acid and an alcohol. The ending -<em>ate</em> or -<em>oate</em> is added to the name of the parent carboxylic acid, and the alcohol is identified using the "alkyl alcohol" convention. The following ester, for example, can be named as a derivative of acetic acid (CH<sub>3</sub>CO<sub>2</sub>H) and ethyl alcohol (CH<sub>3</sub>CH<sub>2</sub>OH). </p> <p align="center"><img src="graphics/61.gif" width="211" height="47"> </p> <p>Or it can be named as a derivative of ethanoic acid (CH<sub>3</sub>CO<sub>2</sub>H) and ethyl alcohol (CH<sub>3</sub>CH<sub>2</sub>OH). </p> <p align="center"><img src="graphics/62.gif" width="227" height="47"> </p> <p>The term <em>ester</em> is commonly used to describe the product of the reaction of any strong acid with an alcohol. Sulfuric acid, for example, reacts with methanol to form a diester known as dimethyl sulfate. </p> <p align="center"><img src="graphics/img74.gif" align="middle" width="326" height="75"> </p> <p>Phosphoric acid reacts with alcohols to form triesters such as triethyl phosphate. </p> <p align="center"><img src="graphics/img75.gif" align="middle" width="350" height="84"> </p> <p>Compounds that contain the <img src="graphics/em.gif" alt="--" width="22" height="9">CO<sub>2</sub>R functional group might therefore best be called <strong>carboxylic acid esters</strong>, to indicate the acid from which they are formed. </p> <p>Carboxylic acid esters with low molecular weights are colorless, volatile liquids that often have a pleasant odor. They are important components of both natural and synthetic flavors (see figure below). </p> <p align="center"><img src="graphics/29a.gif" width="402" height="297"></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="fat"><em><strong>Fats and Oils</strong></em> </a></p> <p>Long-chain carboxylic acids such as stearic acid [CH<sub>3</sub>(CH<sub>2</sub>)<sub>16</sub>CO<sub>2</sub>H] are called <strong>fatty acids</strong> because they can be isolated from animal fats. These fatty acids are subdivided into the two categories on the basis of whether they contain C=C double bonds: <strong>saturated fatty acids</strong> and <strong>unsaturated fatty acids</strong>. </p> <p>There are four important unsaturated fatty acids. One of them is a derivative of palmitic acid, and is known as palmitoleic acid. </p> <p align="center"><img src="graphics/29b.gif" width="288" height="57"> </p> <p>The other three are derivatives of stearic acid. Oleic acid has a single C=C double bond in the center of the fatty acid chain. </p> <p align="center"><img src="graphics/29c.gif" width="254" height="57"> </p> <p>Linoleic acid, has another C=C double bond in the nonpolar half of the fatty acid chain. </p> <p align="center"><img src="graphics/30a.gif" width="360" height="56"> </p> <p>Linolenic acid has one more C=C double bond in the same half of the fatty acid chain. </p> <p align="center"><img src="graphics/30b.gif" width="453" height="56"></p> <p>There are several regularities in the chemistry of these unsaturated fatty acids. First, they contain cis double bonds. Second, the double bonds are always isolated from each other by a CH<sub>2</sub> group. </p> <p>So much attention is paid to the structures of the fatty acids in discussions of these compounds that it is easy to miss an important point: Free fatty acids are seldom found in nature. They are usually tied up with alcohols to form esters (RCO<sub>2</sub>R). The most abundant of these esters is the triester formed when a molecule of glycerol (HOCH<sub>2</sub>CHOHCH<sub>2</sub>OH) combines with three fatty acids, as shown in the figure below. These lipids have been known by a variety of names, including <em>fat</em>, <em>neutral fat</em>, <em>glyceride</em>, <em>triglyceride</em>, and <em>triacylglycerol</em>. </p> <p align="center"><img src="graphics/30c.gif" width="287" height="134"> </p> <p>Most animal fats are complex mixtures of different triglycerides. As the percentage of unsaturated fatty acids increases, the melting point of these triesters decreases until they eventually become an oil at room temperature. Beef fat, which is roughly one-third unsaturated fatty acids, is a solid. Olive oil, which is roughly 80% unsaturated, is a liquid. </p> <p>The effect of unsaturated fatty acids on the melting point of a triglyceride can be understood by recognizing that the cis C=C double bond introduces a rigid 30 bend in the hydrocarbon chain. This bend or "kink" increases the average distance between triglyceride molecules, which decreases the van der Waals interactions between neighboring molecules. Thus, the introduction of unsaturated fatty acids into a triglyceride increases the fluidity of the lipid. The table below compares the relative abundance of the common fatty acids in a typical animal fat (butter) and a vegetable oil (olive oil). </p> <p align="center"><em>Relative Abundance of Fatty Acids in a Typical Fat and a Typical Oil </em></p> <div align="center"><center><table border="0"> <tr> <td><em>Fatty Acid</em></td> <td> </td> <td><em>Butter </em></td> <td> </td> <td><em>Olive Oil</em> </td> </tr> <tr> <td>Butyric </td> <td> </td> <td>3-4% </td> <td> </td> <td> </td> </tr> <tr> <td>Caproic </td> <td> </td> <td>1-2% </td> <td> </td> <td> </td> </tr> <tr> <td>Caprylic </td> <td> </td> <td><1% </td> <td> </td> <td> </td> </tr> <tr> <td>Capric</td> <td> </td> <td>2-3% </td> <td> </td> <td> </td> </tr> <tr> <td>Lauric </td> <td> </td> <td>2-5% </td> <td> </td> <td> </td> </tr> <tr> <td>Myristic </td> <td> </td> <td>8-15% </td> <td> </td> <td><1% </td> </tr> <tr> <td>Palmitic </td> <td> </td> <td>25-29% </td> <td> </td> <td>5-15% </td> </tr> <tr> <td>Stearic </td> <td> </td> <td>9-12% </td> <td> </td> <td>1-4% </td> </tr> <tr> <td>Palmitoleic </td> <td> </td> <td>4-6% </td> <td> </td> <td> </td> </tr> <tr> <td>Oleic </td> <td> </td> <td>18-33% </td> <td> </td> <td>67-84%</td> </tr> <tr> <td>Linoleic </td> <td> </td> <td>2-4% </td> <td> </td> <td>8-12% </td> </tr> <tr> <td>Linolenic</td> <td> </td> <td><1% </td> <td> </td> <td> </td> </tr> </table> </center></div><p>Fats and oils are used by living cells for only one purpose<img src="graphics/em.gif" alt="--" width="22" height="9">to store energy. They are a far more efficient storage system than carbohydrates such as glycogen or starch because they give off between two and three times as much energy when they are burned. (Glycogen releases 15.7 kJ per gram of carbohydrate consumed, whereas lipids give approximately 40 kJ/g.) This explains why the seeds of many plants are relatively rich in oils, which provide the energy the seed needs to grow until the leaves can begin to produce energy by photosynthesis. </p> <p>The average human contains enough fat (21% of the body weight for men, 26% for women) to provide the energy they need to survive for up to 3 months. But there is only enough glycogen stored in the human body at any time to provide enough energy for one day. Thus, glycogen is only used for the short-term storage of food energy. In "times of plenty," the body stores energy in the form of fat to compensate for "times of shortage" to come. </p> <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_carbonyl.dat"></p> </body>