Organic Chemistry

Chapter 1 

Structure and Properties

Organic chemistry

• Organic chemistry is the chemistry of the compounds of carbon. The misleading name "organic" is a relic of the days when chemical compounds were divided into two classes, inorganic and organic, depending upon where they had come from. Inorganic compounds were those obtained from minerals; organic compounds were those obtained from vegetable or animal sources, that is, from material produced by living organisms. Indeed, until about 1850 many chemists believed that organic compounds must have their origin in living organisms, and. consequently could never be synthesized from inorganic material.

• These compounds from organic sources had this in common: they all contained the element carbon. Even after it had become clear that these compounds did not have to come from living sources but could be made in the laboratory, it was convenient to keep the name organic to describe them and compounds like them. The division between inorganic and organic compounds has been retained to this day.

• Today, although many compounds of carbon are still most conveniently isolated from plant and animal sources, most of them are synthesized. They are sometimes synthesized from inorganic substances like carbonates or cyanides, but more often from other organic compounds. There are two large reservoirs of organic material from which simple organic compounds can be obtained '. petroleum and coal. (Both of these are "organic" in the old sense, being products of the decay of plants and animals.) These simple compounds are used as building blocks from which larger and more complicated compounds can be made.

• We recognize petroleum and coal as the fossil fuels, laid down over millennia and nonreplaceable, that are being consumed at an alarming rate to meet our constantly increasing demands for power. There is, fortunately, an alternative source of power nuclear energy but where are we to find an alternative reservoir of organic raw material.

• Organic chemistry is fundamental to biology and medicine. Aside from water, living organisms are made up chiefly of organic compounds; the molecules of "molecular biology" are organic molecules. Ultimately, biological processes are a matter of organic chemistry. 

The structural throe

• "Organic chemistry nowadays almost drives me mad. To me it appears like a primeval tropical forest full of the most remarkable things, a dreadful endless jungle into which one does not dare enter for there seems to be no way out.'* Friedrich Wohler, 1835.

• How can we even begin to study a subject of such enormous complexity? Is organic chemistry today as Wohler saw it a century and a half ago? The jungle is still there largely unexplored and in it are more remarkable things than Wohler ever dreamed of. But, so long as we do not wander too far too fast, we can enter without fear of losing our way, for we have a chart: the structural theory.

• The structural theory is the basis upon which millions of facts about hundreds of thousands of individual compounds have been brought together and arranged in a systematic way. It is the basis upon which these facts can best be accounted for and understood.

• The structural theory is the framework of ideas about how atoms are put together to make molecules. The structural theory has to do with the order in which atoms are attached to each other, and with the electrons that hold them together. It has to do with the shapes and sizes of the molecules that these atoms form, and with the way that electrons are distributed over them.

• A molecule is often represented by a picture or a model sometimes by several pictures or several models. The atomic nuclei are represented by letters or wooden balls, and the electrons that join them by lines or dots or wooden pegs. These crude pictures and models are useful to us only if we understand what they are intended to mean. Interpreted in terms of the structural theory, they tell us a good deal about the compound whose molecules they represent : how to go about making it; what physical properties to expect of it melting point, boiling point, specific gravity, the kind of solvents the compound will dissolve in, even whether it will be colored or not; what kind of chemical behavior to expect the kind of reagents the compound will react 'with and the kind of products that will be formed, whether it will react rapidly or slowly. We would know all this about a compound that we had never encountered before, simply on the basis of its structural formula and what we understand its structural formula. 

The chemical bond before 1926

• Any consideration of the structure of molecules must begin with a discussion of chemical bonds, the forces that hold atoms together in a molecule. We shall discuss chemical bonds first in terms of the theory as it had developed prior to 1926, and then in terms of the theory pf today. The introduction of quantum mechanics in 1926 caused a tremendous change in ideas about how molecules are formed. For convenience, the older, simpler language and pictorial representations are often still used, although the words and pictures are given a modern interpretation.

• In 1916 two kinds of chemical bond were described: the ionic bond by Walther Kossel (in Germany) and the covalent bond by G. N. Lewis (of the University of California). Both Kossel and Lewis based their ideas on the following concept the A positively charged nucleus is surrounded by electrons arranged in concentric shells or energy levels. There is a maximum number of electrons that can be accommodated in each shell: two in the first shell, eight in the second shell, eight or eighteen in the third shell, and so on. The greatest stability is reached when the outer shell is full, as in the noble gases. Both ionic and covalent bonds arise from the tendency of atoms to attain this stable configuration of electrons. 

• The covalent bond results from sharing of electrons, as, for example, in the formation of the hydrogen molecule. Each hydrogen atom has a single electron; by sharing a pair of electrons, both hydrogens can complete their shells of two. Two fluorine atoms, each with seven electrons in the valence shell, can complete their octets by sharing a pair of electrons. In a similar way we can visualize the formation of HF, H2O, NH3 , CH4 , and CF4 . Here, too, the bonding force is electrostatic attraction: this lime between each electron and both nuclei.

Quantum mechanics

• In 1926 there emerged the theory known as quantum mechanics, developed, in the form most useful to chemists, by Erwin Schrodinger (of the University of Zurich). He worked out mathematical expressions to describe the motion of an electron in terms of its energy. These mathematical expressions are called wave equations, since they are based upon the cjncept that electrons show properties not only of particles but also of waves. A wave equation has a series of solutions, called wave functions, each corresponding to a different energy level for the electron. For all but the simplest of systems, doing the mathematics is so time-consuming that at present and superhigh-speed computers will some day change this only approximate solutions can be obtained. Even so, quantum mechanics gives answers agreeing so well with the facts that it is accepted today as the most fruitful approach to an understanding of atomic and molecular structure. "Wave mechanics has shown us what is going on, a d at the deepest possible level ... it has taken the concepts of the experimental chemist the imaginative perception that came to those who had lived in their laboratories and allowed their minds to dwell creatively upon the facts that they had found and it has shown ho they all fit together; how, if you wish, they all have one single rationale; and how this hidden relationship to each other can be brought out.*' C. A. Coulson, London, 1951.

Atomic orbitals

• A wave equation cannot tell us exactly where an electron is at any particular moment, or how fast it is moving; it does not permit us to plot a precise orbital bout the nucleus. Instead, it tells us the probability of finding the electron at any particular.

• The region in space where an electron is likely to be found is called an orbital. There are different kinds of orbitals, which have different sizes and different shapes, and which are disposed about the nucleus in specific ways. The particular kind of orbital that an electron occupies depends upon the energy of the electron. It is the shapes of these orbitals and their disposition with respect to each other that we are particularly interested in, since these determine or, more precisely, can conveniently be thought of as determining the arrangement in space of the atoms of a molecule, and even help chemicbehavior.  

• It is convenient to picture an electron as being smeared out to form a cloud. We might think of this cloud as a sort of blurred photograph of the rapidly moving electron. The shape of the cloud is the shape of the orbital. The cloud is not uniform but is densest in those regions where the probability of finding the electron is highest, that is, in those regions where the average negative charge, or electron density, is greatest.

• No changes have been made just for the sake of change. In our rewriting and in the selection of new topics, we have stuck to the principle we have always held: these are beginning students, and they need all the help they can get. Discussion of neighboring group effects or the Woodward-Hoffmann rules is pitched' at the same level as the chlorination of methane in Chapter 2. New material is introduced at the rate at which we have found students can absorb it. Once presented, a principle is used and re-used. In a beginning course, we cannot hope to cover more than a tiny fraction of this enormous field; but what we can hope for is to make a good job of what we do.

Electronic configuration. Pauli exclusion principle

• There are a number of "rules" that determine the way in which the electrons of an atom may be distributed, that is, that determine the electronic configuration of an atom. The most fundamental of these rules is the Pauli exclusion principle: only two electrons can occupy any atomic orbital, and to do so these two must have opposite spins. These electrons of opposite spins are said to be paired. Electrons of like spin tend to get as far from each other as possible. This tendency is the most important of all the factors that determine the shapes and properties of molecule.

• The exclusion principle, advanced in 1925 by Wolfgang Pauli, Jr. (of the Institute for Theoretical Physics, Hamburg, Germany), has been called the cornerstone of chemistry. The first ten elements of the Periodic Table have the electronic configurations shown in Table 1.1. We see that an orbital becomes occupied only if the orbitals. 

• of lower energy are filled (e.g., 2s after Is, 2p after 2s). We see that an orbital is not occupied by a pair of electrons until other orbitals of equal energy are each occupied by one electron (e.g., the 2p orbitals). The Is electrons make up the first shell of two, and the 2s and 2p electrons make up the second shell of eight. For elements beyond the first ten, there is a third shell containing a 3s orbital, 3p orbitals, and so on.

Molecular orbital 

• In molecules, as in isolated atoms, electrons occupy orbitals, and in accordance with much the same "rules." These molecular orbitals are considered to be centered about many nuclei, perhaps covering the entire molecule; the distribution of nuclei and electrons is simply the one that results in the most stable molecule. To make the enormously complicated mathematics more workable, two simplifying assumptions are commonly made: (a) that each pair of electrons is essentially localized near just two nuclei, and (b) that the shapes of these localized molecular orbitals and their disposition with respect to each other are related in a simple way to the shapes and disposition of atomic orbitals in the component atoms.

• The second assumption of a relationship between atomic and molecular orbitals, is a highly reasonable one, as discussed in the following section. It has proven so useful that, when necessary, atomic orbitals of certain kinds have been invented just so that the assumption can be retained.

The covalent bond

• Now let us consider the formation of a molecule. For convenience we shall picture this as happening by the coming together of the individual atoms, although most molecules are not actually made this way. We make physical models of molecules out of wooden or plastic balls that represent the various atoms; the location of holes or snap fasteners tells us how to put them together. In the same way, we shall make mental models of molecules out of mental atoms; the location of atomic orbitals some of them imaginary will tell us how to put these together.

• For a covalent bond to form, two atoms must be located so that an orbital of one overlap an orbital of the other; each orbital must contain a single electron. When this happens, the two atomic orbitals merge to form a single bond orbital which is occupied by both electrons. The two electrons that occupy a bond orbital must have opposite spins, that is, must be paired. Each electron has available to it the entire bond orbital, and thus may be considered to "belong to" both atomic nuclei.

• As our first example, let us consider the formation of the hydrogen molecule, H2 , from two hydrogen atoms. Each hydrogen atom has one electron, which occupies the \s orbital. As we have seen, this Is orbital is a sphere with its center at the atomic nucleus. For a bond to form, the two nuclei must be brought closely enough together for overlap of the atomic orbitals to occur . For hydrogen, the system is most stable when the distance between the nuclei.

• this distance is called the bond length. At this distance the stabilizing effect of overlap is exactly balanced by repulsion between the similarly charged nuclei. The resulting hydrogen molecule contains 104 kcal/mole less energy than the hydrogen atoms from which it was made. We say that the hydrogen-hydrogen bond has a length of 0.74 A and a strength of 104 kcal.

• Next, let us consider the formation of the fluorine molecule, F2 , from two fluorine atoms. As we can see from our table of electronic configurations, a fluorine atom has two electrons in the Is orbital, two electrons in the 2$ orbital, and two electrons in each of two 2p orbitals. In the third 2p orbital there is a single electron which is unpaired and available for bond formation. Overlap of this/? orbital with a similar p orbital of another fluorine atom permits electrons to pair and the bond to of.

• As the examples show, a covalent bond results from the overlap of two atomic orbitals to form a bond orbital occupied by a pair of electrons. Each kind of covalent bond has a characteristic length and straight.

Hybrid orbitals: sp

• How are we to account for its combining with two chlorine atoms? Bond formation is an energy-releasing (stabilizing) process, and the tendency is to form bonds and as many as possible even if this results in bond orbitals that bear little resemblance to the atomic orbitals we have talked about. If our method of mental molecule-building is to be applied here, it must be modified. We must invent an imaginary kind of beryllium atom, one that is about to become bonded to two chlorine atoms. 

• To arrive at this divalent beryllium atom, let us do a little electronic bookkeeping. First, we "promote" one of the 2s electrons to an empty/?

• This provides two unpaired electrons, which are needed for bonding to two chlorine atoms. We might now expect beryllium to form one bond of one kind, using the p orbital, and one bond of another kind, using the s orbital. Again, this is contrary to fact: the two bonds in beryllium chloride are known to be equivalent* Next, then, we hybridize the* orbitals. Various combinations.

• one p orbitals are taken mathematically, and the mixed (hybrid) orbitals with the greatest degree of directional character are found (Fig. 1 .5). The more an atomic orbital is concentrated in the direction of the bond, the greater the overlap and the stronger the bond it can form. Three highly significant results emerge from the

• Atomic orbitals: hybrid sp orbitals. (a) Cross-section and approximate shape of a single orbital. Strongly directed along one axis. (b) Representation as a sphere, with small back lobe omitted, (c) Two orbitals, with axes lying along a straight line. 

Hybrid orbitals: s/> 

• Now, let us turn to one of the simplest of organic molecules, methane, CH4 . Carbon (Table 1.1) has an unpaired electron in each of the two p orbitals, and on this basis might be expected to form a compound CH2 . (It docs, but

• CH2 is a highly reactive molecule whose properties center about the need to provide carbon v it two more bonds.) Again, we see the tendency to form as many bonds as Posible: in this case, to combine with four hydrogen atoms. To provide; four unpaired electrons, we promote one of the 2s electrons to the empty p 

• Once more the most strongly directed orbitals are hybrid orbitals: this time, sp* orbital's, from the mixing of one 'orbital and three p orbitals. Each one 

• Thus, in these last three sections, we have seen that there are i covalent bonds not only characteristic bond lengths and bond Diji gees but also characteristic bond angles. These bond angles can relate to the arrangement of atomic orbitals including hybrid! w>,Wy volved in bond formation; tyvm r**[ f,, ;^v go back to the Pauli we/i Cipla and the tender tor fudgier ^lacceroic to plot A% for, r from'db&fttt

• The sp 3 orbital occupied by the unshared pair of electrons is a region of high electron density. This region is a source of electrons for electron-seeking atoms and molecules, and thus gives ammonia its basic properties.

Intramolecular forces

• We must remember that the particular method of mentally building molecules that we are learning to use is artificial: it is a purely intellectual process involving imaginary overlap of imaginary orbitals. There are other, equally artificial ways that use different mental or physical models. Our method is the one. that so far has seemed to work out best for the organic chemist. Our kit of mental atomic models will contain just three "kinds" of carbon: tetrahedral C$/? 3 -hybridized), trigonal (sp 2-hybridized), and diagonal (,s/?-hybridized). By use of this kit, we shall find, one can do an amazingly good job of building hundreds of thousands of organic molecules.

• But, however we arrive at it, we see the actual structure of a molecule to be the net result of a combination of repulsive and attractive forces,
  which are related to charge and electron.

Bond dissociation energy. Homolysis and heterolysis

• We have seen that energy is liberated when atoms combine to form a molecule. For a molecule to break into atoms, an equivalent amount of energy must be consumed. The amount of energy consumed or liberated when a bond is broken or formed is known as the bond dissociation energy, D. It is characteristic of the particular bond: Table 1 .2 lists bond dissociation energies that have been measured for a number of bonds. As can be seen, they vary widely, from weak bonds like II (36 kcal/mole) to very strong bonds like H F (136 kcal/mole). Although the accepted values may change as experimental methods improve, certain trends are clear. We must not confuse bond dissociation energy (D) with another measure of bond strength called bond energy (E). If one begins with methane, for example, and breaks, successively, four carbon-hydrogen bonds, one finds four different bond dissociation energies: 

• We have seen that energy is liberated when atoms combine to form a molecule. For a molecule to break into atoms, an equivalent amount of energy must be consumed. The amount of energy consumed or liberated when a bond is broken or formed is known as the bond dissociation energy, D. It is characteristic of the particular bond: Table 1 .2 lists bond dissociation energies that have been measured for a number of bonds. As can be seen, they vary widely, from weak bonds like II (36 kcal/mole) to very strong bonds like H F (136 kcal/mole). Although the accepted values may change as experimental methods improve, certain trends are clear.

• So far, we have spoken of breaking a molecule into two atoms or into an atom and a group of at the two electrons making up the covalent bond, one goes to each fragment; such bond-breaking ishomolysis.We shall also encounter reactions involving bond-breaking of a different kind: heterolysis, in which both bonding electrons go to the same fragment.

Polarity of bonds

• Besides the properties already described^ certain covalent bonds have another property: polarity. Two atoms joined by a covalent bond share electron; their nuclei are held by the same electron cloud. But in most cases the two nuclei do not share ire electrons equally; the electron cloud is denser about one atom than the other. One end of the bond is thus relatively negative, and the other end is relatively positive; that is, there is a negative pole and a positive pole. Such a bond is said to be a polar bond, or to possess polarity.

• We can expect a covalent bond to be polar if it joins atoms that differ in their tendency to attract electrons, that is, atoms that deter in electrogenicity. Furthermore, the greater the difference in electronegativity, the more polar the bond will be.. The most electronegative elements are those located in the upper right-hand corner of the Periodic Table. Of the elements we are likely to encounter in organic chemistry, fluorine has the highest electronegativity, then oxygen, then nitrogen and chlorine, then bromine, and finally carbon. Hydrogen does not differ very much from carbon in electronegativity; it is not certain whether it is more or less electronegative. Electronegativity F > O > Cl, N > Br >.

• Bond polarities are intimately concerned with both physical and chemical properties. The polarity of bonds can lead to polarity of molecules, and thus profoundly affect melting point, boiling point, and solubility. The polarity of a bond determines the kind of reaction that can take place at t]ail bond, and even affects reactivity at nearby.

Polarity of molecules

• (A molecule is polar if the center of negative charge does not coincide with the

• center of positive charge. Such a molecule constitutes a dipole: two equal and

• opposite charges separated in space. A dipole is often symbolized by -f->, where

• the arrow points from positive to negative?) The molecule possesses a dipole

• moment, /z, which is equal to the magnitude of the charge, e, multiplied by the

• distance, </, between the centers of charge:

• Deby c.s.u. Angstrom

• In a way that cannot be gone into here, it is possible to measure the dipole moments of molecules; some of the values obtained are listed in.

• Methane and carbon tetrachloride, CC14 , have zero dipole moments. We certainly would expect the individual bonds of carbon tetrachloride at least to be molar; because of the very symmetrical tetrahedral arrangement, however, they exactly cancel each other out (Fig. 1.14). In methyl chloride, r Hecla, the polarity.

• of the carbon -chlorine bond is not canceled, however, and methyl chloride has a dipole moment of 1.86 r>. Thus, the polarity of a molecule depends not only upon the polarity of its individual bonds but also upon the way the bonds are directed, that is, upon the shape of the molecule.

Structure and physical properties

• We have just discussed one physical property of compounds: dipole moment. Other physical properties like melting point, boiling point, or solubility in a particular solvent are also of concern to us. The physical properties of a new compound give valuable clues about its structure. Conversely, the structure of a compound often tells us what physical properties to expect of i

• In attempting to synthesize a new compound, for example, we must plan a series of reactions to convert a compound that we have into the compound that we want. In addition, we must work out a method of separating our product from all the other compounds making up the reaction mixture: unconsumed reactants, solvent, catalyst, by-products. Usually, the isolation and purification of a product take much more time and effort than the actual making of it. The feasibility of isolating the product by distillation depends upon its boiling point and the boiling points of the contaminants; isolation by recrystallization depends upon its solubility in various solvents and the solubility of the contaminants. Success in the laboratory often depends upon making a good prediction of physical properties from structure. 

• We have seen that there are two extreme kinds' of chemical bonds: ionic bonds, formed by the transfer of electrons, and covalent bonds, formed by the sharing of electrons. The physical properties of a compound depend largely upon which kind of bonds hold its atoms together in the molecule.

Melting point

• In a crystalline solid the particles acting as structural units -ions or molecules-are arranged in some very regular, symmetrical \very, there is a geometric pattern repeated over and over \Vidhin a crystal. Melting is the change from ire highly ordered arrangement of particles in the cr>saline lattice to the more random arrangement that characterizes a liquid (see Figs. 1.16 and- 1.1 7). Melting occurs when a temperature is reached at which the thermal energy of the particles is great enough to overcome the intracrystalline forces that hold them in position. An Ionic compound forms crystals' in which the structural units are ions. Solid sodium chloride, for example, is made up of positive sodium ions and negative chloride ions alternating in a verve regular way. Surrounding each positive.

• ion and equidistant from it are six negative ions: one on each side of it, one above and one below, one in front and one in back. Each negative ion is surrounded in a similar wav by si\ positive ions. There is nothing that we can properly call a molecule of sodium chloride. A particular sodium ion does not "belong" to any one chloride ion; it is equally attracted to six chloride ions. The crystal is an extremely strong, rigid structure, since the electrostatic forces holding each ion in position are powerful. These powerful interionic forces are overcome only at a Verforces holding ions to each other. To melt sodium chloride we must supply enough energy to break ionic bonds between Na + and Cl~. To melt methane, CH4 , we do not need to supply enough energy to break covalent bonds between carbon and hvdrogen; we need only supply enough energy to break CH4 molecules av\ay from each other. In contrast to sodium chloride, methane melts at 183.> high temperature; sodium chloride has a melting point of 801 . Crystals of other ionic compounds resemble crystals of sodium chloride in having an ionic lattice, although the exact geometric arrangement may be different. As a result, these other ionic compounds, too, have high melting points. Many molecules contain both ionic and covalent bonds. Potassium nitrate, Konja, for example, is made up of K 4 ions and NO3 ions; the oxygen and nitrogen atoms of the NCV ion are he Ld to can owner ty covalent bonds. The physical properties of compounds like these are largely determined by the ionic bonds; potassium nitrate has very much the same sort of physical properties as sodium chloride. A non-ionic compound, one whose atoms are held to each other entirely by covalent bonds, forms crystals in which the structural units are molecules.

Intel-molecular forces

• What kind of forces hold neutral molecules to each other? Like interionic forces, these forces seem to be electrostatic in nature, involving attraction of positive charge for negative charge. There are two kinds of intermolecular forces: dipole-dipole interactions and van der Waals forces. Dipole- dipolc interaction is the attraction of the positive end of one polar molecule for the negative end of another polar molecule. In hydrogen chloride, for example, the relatively positive hydrogen of one molecule is attracted to the relatively negative chlorine of another.

• As a result of dipole- dipole interaction, polar molecules are generally held to each other more strongly than are non-polar molecules of comparable molecular weight; this difference in strength of intermolecular forces is reflected in the physical properties of the compounds concerned. An especially strong kind of dipole dipole attraction is hydrogen bonding, in which a hydrogen atom scree as a bridge between two electronegative atoms, holding one by a covalent bond and the other by purely electrostatic forces. When hydrogen is attached to a highly electronegative atom, the electron cloud is greatly distorted toward the electronegative atom, exposing the hydrogen nucleus. The strong positive charge of the thinly shielded hydrogen nucleus is strongly attracted by the negative charge of the' electronegative atom of a second molecule. This attraction has a strength of about 5 kcal/mole and is thus much weaker than the covalent bond about 50-100 kcal/mole that holds it to the first electronegative atom. It is, however, much stronger.

Solubility

• When a solid or liquid dissolves, the structural units ions or molecules become separated from each other, and the spaces in between become occupied by solvent molecules. In dissolution, as in melting and boiling, energy must be supplied to overcome the interionic or intermolecular forces. Where does the necessary energy come from? The energy required to break the bonds between solute particles is supplied by the formation of bonds between the solute particles and the solvent molecules: the old attractive forces are replaced by new ones. A great deal of energy is necessary to overcome the powerful electrostatic forces holding together an ionic lattice. Only water or other highly polar solvents are able to dissolve ionic compounds appreciably. What kind of bonds are formed between ions and a polar solvent? By definition, a polar molecule has a positive end and a negative end. Consequently, there is electrostatic attraction between a positive ion and the negative end of the solvent molecule, and between a negative ion and the positive end of the solvent molecule. These attractions are called iondipole bonds. Each ion-dipole bond is relatively weak, but in the aggregate they supply enough energy to overcome the interionic forces in the crystal. In solution each ion is surrounded by a cluster of solvent molecules and is said to be solvated; if the solvent happens to be water, the ion is said to be hydrated. In solution, as in the solid and liquid states, the unit of a substance like sodium chloride is the ion, although in this case it is a solvated ion (see Fig. 1.20). To dissolve ionic compounds a solvent must also have a high dielectric constant, that is, have high insulating properties to lower the attraction between oppositely charged ions once they are solvated.

• The solubility characteristics of non-ionic compounds are determined chiefly by their polarity. Non-polar or weakly polar compounds dissolve in non-polar or weakly polar solvents; highly polar compounds dissolve in highly polar solvents. "Like dissolves like" is an extremely useful rule of thumb. Methane dissolves in carbon tetrachloride because the forces holding methane molecules to each other and carbon tetrachloride molecules to each other are replaced by very similar forces holding methane molecules to carbon tetrachloride molecules.

Acids and base

• Turning from physical to chemical properties, let us review briefly one familiar topic that is fundamental to the understanding of organic chemistry: acidity and basicity. The terms acid and base have been defined in a number of ways, each definition corresponding to a particular way of looking at the properties of acidity and basicity. We shall find it useful to look at acids and bases from two of these viewpoints; the one we select will depend upon the problem at hand. According to the Lowry-Bronsted definition, an acid is a substance that gives up a proton, and a base is a substance that accepts a proton. When sulfuric acid dissolves in water, the acid H2 SO4 gives up a proton (hydrogen nucleus) to the base H2O to form a new acid H3O + and a new base HSO4 ~. When hydrogen chloride reacts with ammonia, the acid HC1 gives up a proton to the base NH3 to form the new acid NH4 + and the new base Cl".

• If aqueous H2SO4 is mixed with aqueous NaOH, the acid H3O* (hydronium ion) gives up a proton to the base OH" to form the new acid H2O and the new base H2O. When aqueous NH4C1 is mixed with aqueous NaOH, the acid NH4 +.

• pets, and the most general; it includes all the other concepts. A proton is an acid because it is deficient in electrons and needs an electron pair to complete its valence shell. Hydroxide ion, ammonia, and water are bases because they contain electron pairs available for sharing. In boron trifluoride, BF3 , boron has only six electrons in its outer shell and hence tends to accept another pair to complete its octet. Boron trifluoride is an acai.

• We write a formal negative charge on boron in these formulas because it has one more electron half-interest in the pair shared with nitrogen or oxygen than is balanced by the nuclear charge; correspondingly, nitrogen or oxygen is shown with a formal positive charge.

Electronic and steric effects

• Like acidity and basicity, other chemical properties, too, depend upon molecular structure. Indeed, most of this book will be concerned with finding out what this relationship is.

• A particular compound is found to undergo a particular reaction. Not surprisingly, other compounds of similar structure are also found to undergo the same reaction but faster or slower, or with the equilibrium lying farther to the right or left. We shall, first of all, try to see how it is that a particular kind of structure predisposes a compound to a particular reaction. Then and much of our time will be spent with this we shall try to see how it is that variations in molecular structure give rise to variations in reactivity: to differences in rate of reaction or in position of equilibrium.

• To do all this is a complicated business, and to help us we shall mentally analyze the molecule: we shall consider that a molecule consists of a reaction center to which are attached various substituents. The nature of the reaction center determines what reaction occurs. The nature of the substituents determines the reactivity.

• The resonance effect involves delocalization of electrons typically, those called TT (pi) electrons. It depends upon the overlap of certain orbitals, and therefore can operate only when the substituent is located in certain special ways relative to the charge center. By its very nature, as we shall see (Sec. 6.25), the resonance effect is a stabilizing effect, and so it amounts to electron withdrawal from a negatively charged center, and electron release to a positively charged center. A substituent can influence reactivity not only by its electronic effect (inductive and/or resonance), but also, in some cases, by its steric effect: an effect due to crowding at some stage of the reaction, and dependent therefore on the size of the substituent.

Isomerism

• Before we start our systematic study of the different kinds of organic compounds, let us look at one further concept which illustrates especially well the fundamental importance of molecular structure: the concept of isomerism. The compound ethyl alcohol is a liquid boiling at 78. Analysis (by the methods described later, Sec. 2.26) shows that it contains carbon, hydrogen, and oxygen in the proportions 2C:6H:1O. Measurement of its mass spectrum shows that it has a molecular weight of 46. The molecular formula of ethyl alcohol must therefore be C2 H6O. Ethyl alcohol is a quite reactive compound. For example, if a piece of sodium metal is dropped into a test tube containing ethyl alcohol, there is a vigorous bubbling, and the sodium metal is consumed; hydrogen gas is evolved and there is left behind a compound of formula C2 H5ONa. Ethyl alcohol reacts with hydriodic acid to form water and a compound of formula C2H5 I. The compound methyl ether is a gas with a boiling point of 24. It is clearly a different substance from ethyl alcohol, differing not only in its physical properties but also in its chemical properties. It does not react at all with sodium metal.

• ethyl alcohol, it reacts with hydriodic acid, but it yields a compound of formula CH3 I. Analysis of methyl ether shows that it contains carbon, hydrogen, and oxygen in the same proportions as ethyl alcohol, 2C:6H: 1O. It has the same molecular weight as ethyl alcohol, 46. We conclude that it has the same molecular formula, C2 H6O. Here we have two substances, ethyl alcohol and methyl ether, which have the same molecular formula, C2H6O, and yet quite clearly are different compounds. How can we account for the existence of these two compounds? The answer is: they differ in molecular structure. Ethyl alcohol has the structure represented by I, and methyl ether the structure represented by IT. As we shall see, the differences in physical and chemical properties of these two compounds can readily be accounted for on the basis of the difference in structure.