COLUMN CHROMATOGRAPHY

NMR

HPLC

UV VIS SPECTROSCOPY

AAS

ADDITION OF ALKENE REACTION

CARBON AND METHANE

Methane has four covalent C-H bonds. Because all four bonds have the same length and all the bond angles are the same (109.5°),we can conclude that the four C-H bonds in methane are identical.
The potential map of methane shows that neither carbon nor hydrogen carries much of a charge: There are neither red areas, representing partially negatively charged atoms, nor blue areas, representing partially positively charged atoms. The absence of partially charged atoms can be explained by the similar electronegativities of carbon and hydrogen, which cause carbon and hydrogen to share their bonding electrons relatively equally. Methane is a nonpolar molecule.
You may be surprised to learn that carbon forms four covalent bonds since you know that carbon has only two unpaired electrons in its ground-state electronic configuration. But if carbon were to form only two covalent bonds, it would not complete its octet. Now we need to come up with an explanation that accounts for carbon’s forming four covalent bonds. If one of the electrons in the 2s orbital were promoted into the empty 2p atomic orbital, the new electronic configuration would have four unpaired electrons; thus, four covalent bonds could be formed.
Because a p orbital is higher in energy than an s orbital, promotion of an electron from an s orbital to a p orbital requires energy. The amount of energy required is 96 kcal/mol. The formation of four C-H bonds releases 402 kcal/mol of energy because the bond dissociation energy of a single C-H bond is 105 kcal/mol If the electron were not promoted, carbon could form only two covalent bonds, which would release only 201 kcal/mol So, by spending 96 kcal/mol (or 402 Kj/mol) to promote an electron, an extra (or 879 kJ/mol) is released. In other words, promotion is energetically advantageous.

Keep a marshmallow warm with a fire

PENICILLIN

Penicillin contains a strained ring. The strain in the four-membered ring increases the amide’s reactivity. It is thought that the antibiotic activity of penicillin results from its ability to acylate (put an acyl group on) a group of an enzyme that is involved in the synthesis of bacterial cell walls. Acylation inactivates the enzyme,and actively CH2OH growing bacteria die because they are unable to produce functional cell walls. Penicillin has no effect on mammalian cells because mammalian cells are not enclosed by cell walls. To minimize hydrolysis of the ring during storage,penicillins are refrigerated. Bacteria that are resistant to penicillin secrete penicillinase, an enzyme that catalyzes the hydrolysis of the ring of penicillin. The ring-opened product has no antibacterial activity.

More than 10 different penicillins are currently in clinical use. They differ only in the group (R) attached to the carbonyl group. Some of these penicillins are shown here. In addition to their structural differences, the penicillins differ in the organisms against which they are most effective. They also differ in their resistance to penicillinase. For example, ampicillin,a synthetic penicillin, is clinically effective against bacteria that are resistant to penicillin G, a naturally occurring penicillin. Almost 19% of humans are allergic to penicillin G. Penicillin V is a semisynthetic penicillin that is in clinical use. It is not a naturally occurring penicillin; nor is it a true synthetic penicillin because chemists don’t synthesize it. The Penicillium mold synthesizes it after the mold is fed 2-phenoxyethanol, the compound it needs for the side chain.

KIRCHHOFF'S LAW DERIVATION

DALMATIAN AND CHEMISTRY

When amino acids are metabolized, the excess nitrogen is concentrated into uric acid, a compound with five amide bonds. A series of enzyme-catalyzed reactions degrades uric acid to ammonium ion. The extent to which uric acid is degraded in animals depends on the species. Birds, reptiles, and insects excrete excess nitrogen as uric acid. Mammals excrete excess nitrogen as allantoin. Excess nitrogen in aquatic animals is excreted as allantoic acid, urea, or ammonium salts. Dalmatians,unlike other mammals, excrete high levels of uric acid. The reason for this is that breeders of Dalmatians select dogs that have no white hairs in their black spots,and the gene that causes the white hairs is linked to the gene that causes uric acid to be converted to allantoin. Dalmatians, therefore, are susceptible to gout (painful deposits of uric acid in joints).

Benzo[a]pyrene Carcinogenic

Benzo[a]pyrene is one of the most carcinogenic of the aromatic hydrocarbons. This hydrocarbon is formed whenever an organic compound undergoes incomplete combustion. For example, benzo[a]pyrene is found in cigarette smoke, automobile exhaust, and charcoal-broiled meat. Several arene oxides can be formed from benzo[a]pyrene. The two most harmful are the 4,5-oxide and the 7,8-oxide. It has been suggested that people who develop lung cancer as a result of smoking may have a higher than normal concentration of cytochrome P450 in their lung tissue. The 4,5-oxide is harmful because it forms a carbocation that cannot be stabilized by electron delocalization without destroying the aromaticity of an adjacent benzene ring. Thus, the carbocation is relatively unstable, so the epoxide will tend not to open until it is attacked by a nucleophile. The 7,8-oxide is harmful because it reacts with water to form a diol, which then forms a diol epoxide. The diol epoxide does not readily undergo rearrangement (the harmless pathway), because it opens to a carbocation that is destabilized by the electron-withdrawing OH groups. Therefore, the diol epoxide can exist long enough to be attacked by nucleophiles (the carcinogenic pathway).

SURVIVAL COMPOUNDS

Several marine organisms, including sponges, corals,and algae, synthesize organohalides (halogen-containing organic compounds) that they use to deter predators. For example,red algae synthesize a toxic,foultasting organohalide that keeps predators from eating them. One predator, however, that is not deterred is a mollusk called a sea hare. After consuming red algae,a sea hare converts the original organohalide into a structurally similar compound it uses for its own defense. Unlike other mollusks, a sea hare does not have a shell. Its method of defense is to surround itself with a slimy material that contains the organohalide, thereby protecting itself from carnivorous fish.

KEKULE'S DREAM

Friedrich August Kekulé von Stradonitz (1829–1896) was born in Germany. He entered the University of Giessen to study architecture, but switched to chemistry after taking a course in the subject. He was a professor of chemistry at the University of Heidelberg, at the University of Ghent in Belgium, and then at the University of Bonn. In 1890,he gave an extemporaneous speech at the twenty-fifth-anniversary celebration of his first paper on the cyclic structure of benzene. In this speech, he claimed that he had arrived at the Kekulé structures as a result of dozing off in front of a fire while working on a textbook. He dreamed of chains of carbon atoms twisting and turning in a snakelike motion, when suddenly the head of one snake seized hold of its own tail and formed a spinning ring. Recently, the veracity of his snake story has been questioned by those who point out that there is no written record of the dream from the time he experienced it in 1861 until the time he related it in 1890. Others counter that dreams are not the kind of evidence one publishes in scientific papers, although it is not uncommon for scientists to report moments of creativity through the subconscious, when they were not thinking about science. Also, Kekulé warned against publishing dreams when he said, “Let us learn to dream, and perhaps then we shall learn the truth. But let us also beware not to publish our dreams until they have been examined by the wakened mind.” In 1895,he was made a nobleman by Emperor William II. This allowed him to add “von Stradonitz” to his name. Kekulé’s students received three of the first five Nobel Prizes in chemistry: van’t Hoff in 1901, Fischer in 1902, and Baeyer in 1905.

MECHANISM REACTION: USE THE ARROW

VON BAEYER AND BARBITURIC ACID

Johann Friedrich Wilhelm Adolf von Baeyer (1835–1917) was born in Germany. He discovered barbituric acid—the first of a group of sedatives known as barbiturates—in 1864 and named it after a woman named Barbara. Who Barbara was is not certain. Some say she was his girlfriend, but because Baeyer discovered barbituric acid in the same year that Prussia defeated Denmark, some believe he named the acid after Saint Barbara, the patron saint of artillerymen. Baeyer was the first to synthesize indigo,the dye used in the manufacture of blue jeans. He was a professor of chemistry at the University of Strasbourg and later at the University of Munich. He received the Nobel Prize in chemistry in 1905 for his work in synthetic organic chemistry.

DIAMOND AND GRAPHITE

Diamond is the hardest of all substances. Graphite, in contrast, is a slippery, soft solid most familiar to us as the “lead” in pencils. Both materials, in spite of their very different physical properties, contain only carbon atoms. The two substances differ solely in the nature of the carbon–carbon bonds holding them together. Diamond consists of a rigid three-dimensional network of atoms, with each carbon bonded to four other carbons via orbitals. The carbon atoms in graphite, on the other hand,are hybridized,so each bonds to only three other carbon atoms. This trigonal planar arrangement causes the atoms in graphite to lie in flat, layered sheets that can shear off of neighboring sheets. You experience this when you write with a pencil: Sheets of carbon atoms shear off,leaving a thin trail of graphite. There is a third substance found in nature that contains only carbon atoms: buckminsterfullerene. Like graphite, buckminsterfullerene contains only hybridized carbons, but instead of forming planar sheets,the carbons in buckminsterfullerene form spherical structures.

WATER IN CHEMISTRY

Water is the predominant chemical component of living organisms. Its unique physical properties, which include the ability to solvate a wide range of organic and inorganic molecules, derive from water’s dipolar structure and exceptional capacity for forming hydrogen bonds. The manner in which water interacts with a solvated biomolecule influences the structure of each. An excellent nucleophile, water is a reactant or product in many metabolic reactions. Water has a slight propensity to dissociate into hydroxide ions and protons. The acidity of aqueous solutions is generally reported using the logarithmic pH scale. Bicarbonate and other buffers normally maintain the pH of extracellular fluid between 7.35 and 7.45. Suspected disturbances of acid-base balance are verified by measuring the pH of arterial blood and the CO2 content of venous blood. Causes of acidosis (blood pH <> 7.45) may, for example, follow vomiting of acidic gastric contents. Regulation of water balance depends upon hypothalamic mechanisms that control thirst, on antidiuretic hormone (ADH), on retention or excretion of water by the kidneys, and on evaporative loss. Nephrogenic diabetes insipidus, which involves the inability to concentrate urine or adjust to subtle changes in extracellular fluid osmolarity, results from the unresponsiveness of renal tubular osmoreceptors to ADH.

WATER IS AN IDEAL BIOLOGIC SOLVENT Water Molecules Form Dipoles
A water molecule is an irregular, slightly skewed tetrahedron with oxygen at its center (Figure 2–1). The two hydrogens and the unshared electrons of the remaining two sp3 hybridized orbitals occupy the corners of the tetrahedron. The 105-degree angle between the hydrogens differs slightly from the ideal tetrahedral angle, 109.5 degrees. Ammonia is also tetrahedral, with a 107-degree angle between its hydrogens. Water is a dipole, a molecule with electrical charge distributed asymmetrically about its structure. The strongly electronegative oxygen atom pulls electrons away from the hydrogen nuclei, leaving them with a partial positive charge, while its two unshared electron pairs constitute a region
of local negative charge. Water, a strong dipole, has a high dielectric constant. As described quantitatively by Coulomb’s law, the strength of interaction F between oppositely charged particles is inversely proportionate to the dielectric constant ε of the surrounding medium. The dielectric constant for a vacuum is unity; for hexane it is 1.9; for ethanol it is 24.3; and for water it is 78.5. Water therefore greatly decreases the force of attraction between charged and polar species relative to water-free environments with lower dielectric constants. Its strong dipole and high dielectric constant enable water to dissolve large quantities of charged compounds such as salts.

Water Molecules Form Hydrogen Bonds
An unshielded hydrogen nucleus covalently bound to an electron-withdrawing oxygen or nitrogen atom can interact with an unshared electron pair on another oxygen or nitrogen atom to form a hydrogen bond. Since water molecules contain both of these features, hydrogen bonding favors the self association of water molecules into ordered arrays (Figure 2–2). Hydrogen bonding profoundly influences the physical properties of water and accounts for its exceptionally high viscosity, surface tension, and boiling point. On average, each molecule in liquid water associates through hydrogen bonds with 3.5 others. These bonds are both relatively weak and transient, with a half-life of about one microsecond. Rupture of a hydrogen bond in liquid water requires only about 4.5 kcal/mol, less than 5% of the energy required to rupture a covalent OH bond. Hydrogen bonding enables water to dissolve many organic biomolecules that contain functional groups which can participate in hydrogen bonding. The oxygen atoms of aldehydes, ketones, and amides provide pairs of electrons that can serve as hydrogen acceptors. Alcohols and amines can serve both as hydrogen acceptors and as donors of unshielded hydrogen atoms for formation of hydrogen bonds.

INTERACTION WITH WATER INFLUENCES THE STRUCTURE OF BIOMOLECULES

Covalent & Noncovalent Bonds Stabilize Biologic Molecules

The covalent bond is the strongest force that holds molecules together (Table 2–1). Noncovalent forces, while of lesser magnitude, make significant contributions to the structure, stability, and functional competence of macromolecules in living cells. These forces, which can be either attractive or repulsive, involve interactions both within the biomolecule and between it and the water that forms the principal component of the surrounding environment.

Biomolecules Fold to Position Polar & Charged Groups on Their Surfaces Most biomolecules are amphipathic; that is, they possess regions rich in charged or polar functional groups as well as regions with hydrophobic character. Proteins tend to fold with the R-groups of amino acids with hydrophobic side chains in the interior. Amino acids with charged or polar amino acid side chains (eg, arginine, glutamate, serine) generally are present on the surface in contact with water. A similar pattern prevails in a phospholipid bilayer, where the charged head groups of

phosphatidyl serine or phosphatidyl ethanolamine contact water while their hydrophobic fatty acyl side chains cluster together, excluding water. This pattern maximizes the opportunities for the formation of energetically favorable charge-dipole, dipole-dipole, and hydrogen bonding interactions between polar groups on the biomolecule and water. It also minimizes energetically unfavorable contact between water and hydrophobic groups.

ORGANIC CHEMISTRY AND YOU

You are already a highly skilled organic chemist. As you read these words, your eyes are using an organic compound (retinal) to convert visible light into nerve impulses. When you picked up this book, your muscles were doing chemical reactions on sugars to give you the energy you needed. As you understand, gaps between your brain cells are being bridged by simple organic molecules (neuro transmitter amines) so that nerve impulses can be passed around your brain. And you did all that without consciously thinking about it. You do not yet understand these processes in your mind as well as you can carry them out in your brain and body. You are not alone there. No organic chemist, however brilliant, understands the detailed chemical working of the human mind or body very well.

Organic chemistry began as a tentative attempt to understand the chemistry of life. It has grown into the confident basis of vast multinational industries that feed, clothe, and cure millions of people without their even being aware of the role of chemistry in their lives. Chemists cooperate with physicists and mathematicians to understand how molecules behave and with biologists to understand how molecules determine life processes. The development of these ideas is already a revelation at the beginning of the twenty-first century, but is far from complete. We aim not to give you the measurements of the skeleton of a dead science but to equip you to understand the conflicting demands of an adolescent one.

Like all sciences, chemistry has a unique place in our pattern of understanding of the universe. It is the science of molecules. But organic chemistry is something more. It literally creates itself as it grows. Of course we need to study the molecules of nature both because they are interesting in their own right and because their functions are important to our lives. Organic chemistry often studies life by making new molecules that give information not available from the molecules actually present in living things.

This creation of new molecules has given us new materials such as plastics, new dyes to colour our clothes, new perfumes to wear, new drugs to cure diseases. Some people think that these activities are unnatural and their products dangerous or unwholesome. But these new molecules are built by humans from other molecules found on earth using the skills inherent in our natural brains. Birds build nests; man makes houses. Which is unnatural? To the organic chemist this is a meaningless distinction. There are toxic compounds and nutritious ones, stable compounds and reactive ones—but there is only one type of chemistry: it goes on both inside our brains and bodies and also in our flasks and reactors, born from the ideas in our minds and the skill in our hands. We are not going to set ourselves up as moral judges in any way. We believe it is right to try and understand the world about us as best we can and to use that understanding creatively.

PHENYL CINNAMATE

Into a 250 ml. Claisen flask place 72 g. of cinnamic acid and 60 g. (37 ml.) of redistilled thionyl chloride. Stopper the side arm, fit the flask with a reflux condenser the top of which is connected to a gas absorption device, and mount the entire apparatus at an angle so that the condensate will not run into the side arm. Heat the mixture on a water bath, cautiously at first, until hydrogen chloride ceases to be evolved (about 1 hour), allow to cool, and add 47 g. of pure (e.g., A.R.) phenol. Heat the mixture on a water bath until no further evolution of hydrogen chloride is observed (about 1 hour). Then place the apparatus on an asbestos-centred wire gauze and heat the flask until the contents are brought just to the reflux temperature in order to complete the reaction : do not heat unduly long as prolonged heating leads to loss of product due to decomposition and polymerisation. Allow the reaction -mixture to cool and distil under diminished pressure ; collect the fraction of b.p. 190-210°/15 mm. This solidifies to a pale yellow solid, m.p. 66-69°, weighing 98 g. Grind it to a powder in a glass mortar and wash the powder with 250 ml. of cold 2 per cent, sodium bicarbonate solution. Recrystallise from rectified spirit (150 ml.) : 81 g. of pure phenyl cinnamate (white crystals) of m.p. 75-76° are obtained.

PHENYLHYDRAZINE

Place 130 ml. of concentrated hydrochloric acid in a 1 -5 litre round-bottomed flask, equipped with a mechanical stirrer and immersed in a freezing mixture of ice and salt. Start the stirrer and, when the temperature has fallen to about 0°, add 60 g. of finely-crushed ice (1), run in 47-5 g. (46-5 ml.) of pure aniline during about 5 minutes, and then add another 60 g. of crushed ice.
Dissolve 35 g. of sodium nitrite in 75 ml. of water, cool to 0-3°, and run in the cold solution from a separatory funnel, the stem of which reaches nearly to the bottom of the flask. During the addition of the nitrite solution (ca. 20 minutes), stir vigorously and keep the temperature as near 0° as possible by the frequent addition of crushed ice. There should be a slight excess of nitrous acid (potassium iodide- starch paper test) at the end of 10 minutes after the last portion of nitrite is added.

In the meantime, prepare a sodium sulphite solution as follows. In a 2-litre beaker or bolt-head flask place 50 g. of sodium hydroxide (2) and add 500 ml. of water. When the sodium hydroxide has dissolved, add 112-5 g. of recrystallised sodium bisulphite (3), and stir mechanically until the solid has dissolved. Cool the resulting solution to about 25° and add a few drops of phenolphthalein indicator solution. Introduce small quantities of sodium bisulphite until the pink colour of the solution just disappears, then stir in a further 12 g. of sodium bisulphite (the total weight required should not exceed 135-140 g.). Cool this solution, with stirring, to about 5° by immersion in an ice bath, then add about 60 g. of crushed ice. Run in the ice-cold diazonium solution as rapidly as possible, while stirring vigorously.

The reaction mixture immediately acquires a bright orange-red colour. Slowly heat the solution to 60-70° on a water bath and maintain this temperature for 30-60 minutes, i.e., until the colour becomes quite dark. Acidify the solution to litmus with con- centrated hydrochloric acid (40-50 ml. are required) ; continue the heating on a boiling water bath until the colour becomes much lighter and in any case for 4-6 hours. If any solid is present, filter the solution. To the hot, clear solution add, with stirring, 500 ml. of concentrated hydrochloric acid ; cool, first in running water, and then in a freezing mixture to 0°. The phenylhydrazine hydrochloride separates as yellowish or pinkish crystals. Collect them on a Buchner funnel, drain, wash with 25 ml. of dilute hydrochloric acid (1:3) , and press well with a large glass stopper (4). Liberate the free base by adding to the phenylhydrazine hydrochloride 125 ml. of 25 per cent, sodium hydroxide solution. Extract the phenyl- hydrazine with two 40 ml. portions of benzene, dry the extracts with 25 g. of sodium hydroxide pellets or with anhydrous potassium carbonate : thorough drying is essential if foaming in the subsequent distillation is to be avoided. Most of the benzene may now be distilled under atmo- spheric pressure, and the residual phenylhydrazine under reduced pressure. For this purpose, fit a small dropping funnel to the main neck of a 100 ml. Claisen flask (which contains a few fragments of porous porcelain) and assemble the rest of the apparatus but do not connect the " Perkin triangle " to the pump. Run in about 40 ml. of the benzene, solution into the flask, heat the latter in an air bath so that the benzene distils over steadily. Allow the remainder of the benzene solution to run in from the dropping funnel as fast as the benzene itself distils over. When all the benzene solution has been introduced into the flask, close the stopcock on the funnel, and continue the heating until the temperature on the thermometer reads about 90°. Allow to cool. Replace the dropping funnel by a rubber stopper carrying a capillary tube reaching to the bottom of the flask, and distil under diminished pressure.
Collect the phenylhydrazine at 137-138°/18 mm. (or at 119-120°/12 mm.). The yield of almost colourless liquid is 70 g. ; it crystallises on cooling in ice and then melts at 23°. Phenylhydrazine slowly darkens on exposure to light.

CAUTION. Phenylhydrazine is highly poisonous and produces unpleasant burns in contact with the skin. Wash off immediately any liquid which has come into contact with the skin first with 2 per cent, acetic acid, then with soap and water.

Notes.
(1) External cooling may be dispensed with if more ice is added.
(2) This weight assumes 100 per cent, purity of the sodium hydroxide. If the commercial solid is used, its purity should be determined and a corresponding adjustment made in the weight.
(3) The sodium sulphite solution may also be prepared by dissolving 100 g. of pure (or a corresponding quantity of commercial) sodium hydroxide in about 125 ml. of water, and then diluting to 750 ml. The flask is cooled in running water, a few drops of phenolphthalein indicator are added, and sulphur dioxide passed in until the pink colour just disappears (it is advisable to add a further 1-2 drops of the indicator at this point) and then for 2-3 minutes longer. It is best to remove a sample for test from time to time, dilute with 3-4 volumes of water, and test with 1 drop of phenolphthalein.
(4) If desired, the phenylhydrazine hydrochloride may be purified by recrystallisation. The crude hydrochloride is boiled with 6 times its weight of water and a few grams of decolourising carbon. After filtering, a volume of concentrated hydrochloric acid equal in volume to one-third of the solution is added, and the mixture cooled to 0°. Pure white crystals are obtained in 85-90 per cent, yield.