Organic chemistry-1-.pdf

ORGANIC CHEMISTRY-1- For 1st stage pharmacy students Assistant Prof. Dr. Ibrahim. J. Sahib 2023-2024 Development of Organic Chemistry as a Science In the past …, Development of Organic Chemistry as a Science In 1828, Wohler (a German chemist) Organic Chemistry: Organ...

ORGANIC CHEMISTRY-1- For 1st stage pharmacy students Assistant Prof. Dr. Ibrahim. J. Sahib 2023-2024 Development of Organic Chemistry as a Science In the past …, Development of Organic Chemistry as a Science In 1828, Wohler (a German chemist) Organic Chemistry: Organic chemistry forms the basis of biochemistry, in which various aspects of health and diseases are studied. Biochemical knowledge is very important for practicing nutritional, medical, and related life sciences. In addition, organic chemistry paved the way for the development of medicinal chemistry, pharmaceutical organic chemistry, bioinformatics, biotechnology, gene therapy, pharmacology, pathology, chemical engineering, dental science, and so on. Organic substances play such a vital role in our daily lives that all of us should know about organic chemistry to understand how it influences our life processes. 1 Organic Compounds contain mainly the elements carbon and hydrogen, its obtained from vegetable and animal sources, and can be synthesized from inorganic material. Organic compounds are divided into two main classes: First: hydrocarbons, contain only two elements, carbon and hydrogen. It can be classified into Aliphatic and Aromatic compounds. Aliphatic compounds are open-chain compounds such as alkanes, alkenes, alkynes, and their cyclic analogs. Aromatic compounds are benzene and compounds that resemble benzene in chemical behavior. The Unique Nature of Carbon Ability to form four strong covalent bonds Electronic configuration of carbon (ground state) : 1s22s22p2 2 Hund’s Rule Every orbital in a sub shell is singly occupied one electron before any one orbital is doubly occupied All electrons in singly occupied orbitals have the same spin Ability to form four strong covalent bonds Each carbon atom has four unpaired electrons when excited Tend to form four strong covalent bonds Ability to Form Multiple Bonds 3  Organic chemists, with carbon chemistry as their subject, have developed all kinds of shorthand phrases to describe structures and phenomena that might otherwise take a sentence of two to explain. Here‟s today‟s example: the terminology of carbon-containing functional groups: primary, secondary, tertiary, quaternary.  Primary carbons, are carbons attached to one other carbon. (Hydrogens – although usually 3 in number in this case – are ignored in this terminology, as we shall see).  Secondary carbons are attached to two other carbons.  Tertiary carbons are attached to three other carbons.  quaternary carbons are attached to four other carbons. You can‟t go higher than that. To have five substitutents, you‟d need 10 electrons around carbon, a clear violation of the octet rule. When people do write 5 covalent bonds around carbon, it‟s a mistake. (In the trade, these are often called Texas carbons – 1) because it resembles a star, 2) because everything is bigger in Texas, and 3) because the only man who can put five bonds on carbon is also known as Walker, Texas Ranger.) 4 We use the same terminology for carbocations. A primary carbocation is attached to one other carbon, a secondary to two, and a tertiary to three. You can‟t have a quaternary carbocation without violating the octet rule either (you‟d need an extra empty p orbital for that, bringing the total to 5). Alcohols also follow the primary/secondary/tertiary nomenclature. The rule for alcohols is that they are named according to the number of carbons attached to the carbon bearing the hydroxyl group: in other words, whether the hydroxyl bound to a primary, secondary, or tertiary carbon. You can‟t have a quaternary alcohol – again, that would involve breaking the octet rule. [A bit of non-essential nomenclature: the carbon attached to the OH is sometimes referred to as the “carbinol” carbon]. 5 A related category of compounds are the alkyl halides, which we encounter in SN1/SN2/E1/E2 reactions (among many others). The naming for alkyl halides is similar to that for alcohols: they are named according to the number of carbons attached to the halogen, where halogen is fluorine, chlorine, bromine, or iodine. Next, we come to amines, which are a little bit different. They are named according to the number of carbons attached to nitrogen. Primary, secondary, and tertiary amines are nitrogens bound to one, two and three carbons, respectively. Since the nitrogen has a lone pair, it is still possible to form another bond to carbon. These are called quaternary amines, although they bear a positive charge on nitrogen and are not at all basic. They are often referred to as quaternary ammonium salts. You‟ll see the -ium ending quite a bit – it designates a positively charged species. 6 Finally, amides also fall into this category. A primary amide is bound to one carbon – the carbonyl carbon. Successive substitutions of hydrogen for carbon turn the amide into secondary and tertiary amides. [You might ask – can you have quaternary amides? Well, yes. Except nobody calls them that – they‟re quite unstable, and go by another name. Why? Because. Why is the plural of goose, “geese”, but the plural of moose is “moose”? There is no satisfying answer to nomenclature questions] Second: contains function groups in addition to carbon and hydrogen contain other groups such as halides (F, Cl, Br, I), hydroxyl groups (-OH), carboxyl group (-COOH), carbonyl group (C=O) …. etc. 7  PHYSICAL PROPERTIES OF DRUG MOLECULES 1. Physical State Drug molecules exist in various physical states, for example, amorphous solids, crystalline solids, hygroscopic solids, liquids, or gases. The physical state of drug molecules is an important factor in the formulation and delivery of drugs. 2. Melting Point and Boiling Point The melting point is the temperature at which a solid becomes a liquid, while the boiling point is the temperature at which the vapor pressure of the liquid is equal to the atmospheric pressure. The melting point is used to characterize organic compounds and to confirm their purity. The melting point of a pure compound is always higher than that of a compound mixed with an impurity. The melting point increases as the molecular weight increases, and the boiling point increases as the molecular size increases. The more highly branched alkane has a lower boiling point than alkane with the same molecular weight. 3. Polarity and Solubility Polarity is a physical property of a compound, that relates to other physical properties, for example, melting and boiling points, solubility, and intermolecular interactions 8 between molecules. Generally, there is a direct correlation between the polarity of a molecule and the number and types of polar and nonpolar covalent bonds. In a few cases, a molecule having polar bonds, but in a symmetrical arrangement, may give rise to a nonpolar molecule, for example, carbon dioxide (CO2). The term bond polarity describes the sharing of electrons between atoms. In a nonpolar covalent bond, two atoms share the electrons equally. A polar covalent bond is one in which one atom has a greater attraction for the electrons than the other atom. When this relative attraction is strong, the bond is ionic. The polarity of a bond arises from the different electronegativities of the two atoms that participate in bond formation. The greater the difference in electronegativity between the bonded atoms, the greater the bond polarity. Thus, the electronegativity of an atom is related to bond polarity. For example, water is a polar molecule, whereas cyclohexane is nonpolar. More examples of polar and nonpolar molecules are shown in the following Table. Life occurs exclusively in water. Solutions in which water is the dissolving medium are called aqueous solutions. In aqueous solutions, the polar parts are hydrated and the nonpolar parts are excluded. Hydrogen bonding is a consequence of the basic molecular 9 structure of water. Water has a very high boiling point compared with small organic molecules due to hydrogen bonding. Examples of some common solvents and their boiling points are compared with the boiling points of water in the following table. Solubility is the amount of solute dissolved in a specific solvent under given conditions. The process of dissolving a solute in a solvent is called solvation. The interaction between a dissolved species and the molecules of a solvent is called solvation. The process of mixing a solute and solvent to form a solution is called dissolution. Solubility largely depends on temperature, polarity, molecular size, and stirring. Temperature always affects solubility, and an increasing temperature usually increases the solubility of most solids in a liquid solvent. The polarity of the solute and solvent also affects their solubility. The stronger the attraction between solute and solvent molecules, the greater the solubility. In general, like dissolves like; that is, materials with similar polarity are soluble in each other. The size matters. Organic molecules with a branching carbon increase their solubility more than a long-chain carbon because branching reduces the size of the molecule and makes it easier to solvate. For example, isobutanol is more soluble in water than butanol. ACID-BASE PROPERTIES Drug molecules contain various types of functional groups, and these functional groups contribute to the overall acidity or basicity of drug molecules. One of the adverse effects of aspirin is stomach bleeding, which is partly due to its acidic nature. In the stomach, aspirin is hydrolyzed to salicylic acid and acetic acid. The carboxylic acid group (─COOH) and a phenolic hydroxyl group (─OH) present in salicylic acid, make this molecule acidic. Moreover, acetic acid is formed, and that is also moderately acidic. Thus, the intake of aspirin increases the acidity of the stomach significantly, and if this increased 10 acidic condition stays in the stomach for a long period, it may cause stomach bleeding. Like aspirin, some other drug molecules are acidic. Similarly, there are basic and neutral drugs as well. Now, let us see what the terms acid, base, and neutral compounds mean, and how these parameters are measured. Most drugs are organic molecules and can be acidic, basic, or neutral. 2. Electronegativity and Acidity Electronegativity is a measure of the tendency of an atom to attract a bonding pair of electrons. The stability of A− determines the relative acidity of HA within a period. The greater the electronegativity, the greater the stability of A−. We know that carbon is less electronegative than nitrogen, which in turn is less electronegative than oxygen, and that oxygen is less electronegative than fluorine. Therefore, the strength of acidity increases from methane to hydrogen fluoride, as shown next. A molecule is said to have resonance when its structure cannot be adequately described by a single Lewis structure (resonance is possible whenever a Lewis structure has a multiple bond and an adjacent atom with at least one lone pair of electrons). Resonance may delocalize the electron pair that A− needs to form the new bond with a proton. Delocalization increases the stability of A− that also decreases the reactivity. A base that has resonance delocalization of the electron pair that is shared with the proton will therefore be less basic than a base without this feature. Since a weaker base has a 11 stronger conjugate acid, a compound whose conjugate base has resonance stabilization will be more acidic. Both carboxylic acids and alcohols contain an ─OH group, but a carboxylic acid is a stronger acid than an alcohol. As we can see, the deprotonation of ethanol (CH3CH2OH) affords the ethoxide ion (CH3CH2O−), which has no resonance (only one Lewis structure can be drawn), but deprotonation of acetic acid (CH3CH2CO2H) affords an acetate ion (CH3CH2CO2−) that has resonance (two contributing Lewis structures can be drawn). TYPES OF CHEMICAL BONDING A chemical bond is the attractive force that holds two atoms together. Valence electrons take part in bonding. An atom that gains electrons becomes an anion, a negatively charged ion, and an atom that loses electrons becomes a cation, a positively charged ion. Metals lose electrons and become electropositive; non-metals gain electrons and become electronegative. While cations are smaller than atoms, anions are larger. The energy required for removing an electron from an atom or ion is called ionization energy. Atoms can form either ionic or covalent bonds to attain a complete outer shell electronic configuration. 12 1. Ionic Bonds Ionic bonds result from the transfer of one or more electrons between atoms. The more electronegative atom always gains one or more valence electrons, and hence becomes an anion. The less electronegative atom always loses one or more valence electrons and becomes a cation. The attraction of opposite charges holds ionic compounds together. Therefore, ionic bonds consist of the electrostatic attraction between positively and negatively charged ions. Ionic bonds involve the complete transfer of electrons between two atoms of widely different electronegativities, and both atoms obtain a stable octet outermost shell of electrons. Thus, ionic bonds are commonly formed between reactive metals, electropositive elements on the left side of the periodic table, and non-metals, electronegative elements on the right side of the periodic table. For example, Na (electronegativity: 0.9) easily gives up an electron, and Cl (electronegativity: 3.0) readily accepts an electron to form an ionic bond. In the formation of the ionic compound Na+Cl−, the single valence electron of Na is transferred to the partially filled valence shell of chlorine. 2. Covalent Bonds In general, most bonds within organic molecules, including various drug molecules, are covalent. The exceptions are compounds that possess metal atoms where the metal atoms should be treated as ions. Covalent bonds are formed by the sharing of the electron pairs between bonded atoms instead of giving up or gaining electrons. In this case, an atom can obtain a filled valence shell by sharing electrons. There are two types of covalent bonds: First, nonpolar covalent bond, the electrons are shared equally between two atoms. An ideal example is the bonding between two identical atoms, for example, H 2, O2, N2, Cl2, and F2. Nonpolar covalent bonds can also occur between different atoms that have identical electronegativity values, for example, CH4. 13  Second, polar covalent bond, the electrons are attracted to one atom more strongly than the other. Such an unequal sharing of the pair of bonding electrons results in a polar covalent bond, for example, HF, HCl, and H2O. There are three types of non-bonding intermolecular interactions: dipole-dipole interactions, van der Waals forces, and hydrogen bonding. Dipole-Dipole Interactions: The interactions between the positive end of one dipole and the negative end of another dipole are called dipole-dipole interactions. As a result of dipole-dipole interactions, polar molecules are held together more strongly than nonpolar molecules. Dipole-dipole interactions arise when electrons are not equally shared in the covalent bonds because of the difference in electronegativity. For example, hydrogen fluoride has a dipole moment of 1.98 D, which lies along the H─F bond. As the fluorine atom has greater electronegativity than the hydrogen atom, electrons are therefore pulled strongly towards fluorine, as shown. Dipole-dipole interactions are stronger than van der Waals forces, but not as strong as ionic or covalent bonds. Van der Waals Forces: Relatively weak forces of attraction that exist between nonpolar molecules are called van der Waals forces or London dispersion forces. These forces are distance-dependent interactions between atoms or molecules. van der Waals forces are the weakest of all the intermolecular interactions. Electrons move continuously within bonds and molecules, so at any time, one side of the molecule can have more electron density than the other side, which gives rise to a temporary dipole. Because the dipoles in the molecules are induced, the interactions between the molecules are also called induced dipole–induced dipole interactions. Alkenes are nonpolar molecules because the electronegativities of carbon and hydrogen are similar. Consequently, there are no 14 significant partial charges on any of the atoms in an alkane. Therefore, the size of the van der Waals forces that hold alkane molecules together depends on the area of contact between the molecules. The greater the area of contact, the stronger the van der Waals forces, and the greater the amount of energy required to overcome these forces. For example, the isobutane (bp: −10.2 °C) and butane (bp: −0.6 °C), both with the molecular formula C4H10, have different boiling points. Isobutane is a more compact molecule than butane. Thus, butane has a greater surface area for interaction with each other than isobutane. The stronger interactions that are possible for n-butane are reflected in its boiling point, which is higher than the boiling point of isobutane. Hydrogen Bonding: Hydrogen bonding is the attractive force between the hydrogen attached to an electronegative atom of one molecule and an electronegative atom of the same (intramolecular) or a different molecule (intermolecular). It is an unusually strong force of attraction between highly polar molecules in which hydrogen is covalently bonded to nitrogen, oxygen, or fluorine. Therefore, a hydrogen bond is a special type of interaction between atoms. A hydrogen bond is formed whenever a polar covalent bond involving a hydrogen atom is near an electronegative atom such as O or N. The attractive forces of hydrogen bonding are usually indicated by a dashed line rather than the solid line used for a covalent bond. For example, water molecules form intermolecular hydrogen bonding, and the hydrogen bonding in the H2O graphic shows a cluster of water molecules in the liquid state. 15  The hydrogen bond is of fundamental importance in biology. The hydrogen bond is said to be the „bond of life‟. The double helix structure of DNA is formed and held together with the hydrogen bonds. The nature of the hydrogen bonds in proteins dictates their properties and behavior. Intramolecular hydrogen bonds (within the molecule) in proteins result in the formation of globular proteins, for example, enzymes or hormones. On the other hand, intermolecular hydrogen bonds (between different molecules) tend to give insoluble proteins, such as fibrous proteins. In cellulose, a polysaccharide, molecules are held together through hydrogen bonding, which provides plants with rigidity and protection. In drug-receptor binding, hydrogen bonding often plays an important role. 16