Minimum -E

Internuclear distance -^^

Figure 3. Energy of two hydrogens as a function of internuclear distance.

A large amount of energy is required to break the H-H bond of H2 - 104 kcal/mol - (a mol is defined as 6.02x1023 atoms, Avogadro's number). In the pure state the H2 molecule can be heated to 500 °C without any decomposition.

The bonding diagram for H2 allows us to explain why H2 does not react with another H atom to form H3 as a stable molecule. The bonding orbital MO(1) of H2 is filled since no more than two electrons can occupy a MO. No significant bonding can be attained by combining MO(2) of H2 with the AO of the H atom.

Most of the molecules in this book contain hydrogen. In every case each hydrogen forms one and only one electron-pair bond with its partner leading us to the first rule:

Rule #1: H forms just one electron-pair bond — called a covalent single bond.

This rule is useful in deriving the actual structures of molecules. Thus, the formula of water, H20, and Rule #1 lead to the structure H-O-H, rather than O-H-H, or something else. Similarly, since methane has the composition CH4, we can draw the structure:

Three Depictions of H2

There are a number of alternative ways of representing H2. In addition to a line drawing H-H, we can use a ball-and-stick drawing or a space-filling diagram, as follows:

Line Ball-and-Stick Space-filling drawing Model Model

Figure 4. Three different representations of the hydrogen molecule.

The space-filling representation contains additional information because it tells us about the shape and size of a molecule. These features are exceedingly important in determining how two molecules can fit together or interact and also how close they can get. With molecules, as with macroscopic objects, two things cannot occupy the same space at the same time.

Bonding in Carbon Compounds

Carbon, the element with atomic number six has six protons in the nucleus and six electrons surrounding it. Two of these fill the low energy 1s orbital and play no role in chemical bonding. The remaining four electrons can form bonds utilizing a spherical 2s orbital and three dumbbell-shaped, mutually perpendicular 2p orbitals, the shapes of which are illustrated in Figure 5. The boundary of each dumbbell encloses about 90% of the total electron cloud. The three p orbitals are usually designated as 2px, 2py and 2pz because they can be placed along the axes of a rectilinear coordinate system. The sum of the electron densities for all three 2p orbitals is a spherical cloud of larger diameter than the 2s orbital.

1 s orbital 2s orbital

1 s orbital 2s orbital

2px orbital 2py orbital 2pz orbital

2px orbital 2py orbital 2pz orbital

Superposition of

Figure 5. Representation of s and p atomic orbitals. The boundary of each sphere and dumbbell encloses about 90% of the total electron cloud.

z sp

Figure 6. Combination of one 2s and one 2px atomic orbital (AO) to form two sp hybrid orbitals.

The 2s orbital can mix with three 2p orbitals to form hybrid atomic orbitals in three different ways. A sequence for generating the three different 2s/2p hybrids is shown in Figures 68. The number of hybrid orbitals always equals the number of AOs from which they derive.

7fk sp

Figure 7. Combination of two sp hybrid orbitals and a 2py atomic orbital to form three sp2 hybrid orbitals.

Figure 8. Combination of three sp hybrid orbitals and a 2py orbital to form four sp3 hybrid orbitals.

Hybrid Orbitals for Tetracoordinate Carbon

If all three of the 2p orbitals and the 2s orbital are hybridized, a set of 4 equivalent orbitals results, the axes of which are directed at the vertices of a tetrahedron (Figure 9a). This arrangement is used in forming methane and other compounds having four atoms attached to carbon (tetracoordinate carbon). The attachment of four hydrogen atoms (4 H) to a carbon atom results in the formation of four electron-pair bonds and the complete filling of the 4 bonding MOs by eight electrons (4 from C and 4 from H).

T sp3


Ball-and-Stick Model Space-filling Model

Figure 9a. Formation of methane from four sp3 hybrid orbitals of carbon and four 1s orbitals of hydrogen.

The angle between any two of the C-H bonds in methane is 109.5°, the angle between the center of a tetrahedron and any two of the vertices (Figure 9b). The internuclear distance of each C-H bond is 1.54 x 10"8 cm or 1.54 angstroms (A). A second useful bonding rule emerges from these facts.

Rule #2: Carbon (C) can bond to a maximum of four atoms (tetracoor-dination). The preferred angle between any two of the bonds is 109.5°. Tetracoordinate carbon utilizes sp3 hybrid orbitals.

Figure 9b. The angle between the center of a tetrahedron and any two of its vertices is 109.5°.

Deviation of the bond angles at an sp3-hybridized carbon atom from the preferred value of 109.5° leads to a higher energy (i.e., less stable) structure. The destabilization increases to a value of about 6 kcal/mol for an angle of 90°. This destabilization, called angle strain, influences the chemistry and properties of a compound.

Carbon forms strong bonds to most atoms, including H, oxygen (O), nitrogen (N), chlorine (CI), and by no means least, to itself. Thus, methane is just the first in a large family of compounds of carbon and hydrogen (hydrocarbons). That family includes the straight-chain saturated hydrocarbons, the first five members of which are shown in Figure 10.


H H Ethane

H H H Propane



Figure 10. The simplest straight-chain hydrocarbons.

Branching of Carbon Chains

Carbon chains can also be branched.


H H-Ç H H ¿-H H-C-Ç-Ç-C H 1 Ù C^ H-<?s H ^H L, H H


Figure 11a. Two simple branched hydrocarbons, isobutane and trimethylpentane.

H H-Ç H H ¿-H H-C-Ç-Ç-C H 1 Ù C^ H-<?s H ^H L, H H


Figure 11a. Two simple branched hydrocarbons, isobutane and trimethylpentane.

There is a simpler notation for depicting the structures of carbon compounds in which the hydrogens are omitted. This shorthand notation leaves it to the reader to add the number of hydrogens corresponding to tetracoor-dination (Figure 11b).

Isobutane is equivalent to is equivalent to



h3c ch3 ch

Figure 11b. Simplified notation of carbon compounds In which the hydrogens are omitted.

Cyclic Structures

Carbon can form rings by bonding with itself in a cycle as well as chains. The simplest members of the family of cyclic hydrocarbons are shown in Figure 12.




Tricoordinate Carbon Compounds. The Double Bond

Carbon can form compounds in which three atoms are linked to it using hybrid orbitals generated from the combination of the 2s atomic orbital with two of the 2p atomic orbitals. The orbitals of trigonally hybridized carbon are shown in Figure 13.

Side-on view

Side-on view

Figure 13. Side-on and Top-on views of the orbitals of a trigonally hybridized (sp2-hybridized) carbon atom.

Two of the simplest carbon compounds that involve tricoordinate (trigonally hybridized) carbon atoms are formaldehyde (H2C=0) and ethylene (H2C=CH2). These are planar molecules that contain a double bond to carbon as well as single (electron-pair) bonds to the hydrogens (Figure 14).

Formaldehyde m o or h,c



Figure 12. The first four members of cyclic hydrocarbons: cyclopropane, cyclobutane, cyclopentane and cyclohexane.


Figure 14. Formaldehyde and ethylene contain trigonally hybridized carbon.

Bonding to tricoordinate carbon utilizes three sp2 hybrid orbitals and the remaining 2p AO. These orbitals allow the derivation of the correct geometry of molecules with tricoordinate carbon. For example, the planar structure of ethylene results because overlap of the p-orbitals is maximum when they are parallel, as shown in Figure 15a.

n bond


n bond


a bond

k bond a bond k bond

Figure 15a. Formation of ethylene by the combination of two carbon atoms and four hydrogen atoms.

The linkage between the two carbons of ethylene is called a double bond because it involves four electrons. The double bond can be represented by two lines, as in the drawing in Figure 14, or by a and n bonds as shown in Figure 15a. The k bond, formed by the side-by-side combination of two parallel p atomic orbitals (shown in blue in Figure 15a), has two lobes, one above and one below the molecular plane. The o bond, formed by the combination of two colinear sp2 orbitals, is symmetric about the C-C axis (axial symmetry).

c-c o bond

a bond

Figure 15b. rc-Bond of ethylene.

Rule #3: Tricoordinate carbon is connected to each of the three attached atoms in a planar arrangement. The bonding involves three in-plane hybrid sp2 orbitals and an orthogonal p atomic orbital.

Dicoordinate Triple Bond.

Carbon Compounds. The

Dicoordinate carbon compounds utilize two sp orbitals formed from the hybridization of the 2s orbital with one 2p orbital, and also the remaining two 2p orbitals. The orbitals for this type of carbon are shown in Figure 16; note that the angle between the two sp orbitals is 180°. The py and p2 orbitals that are not involved in hybridization remain unchanged.

Side view

View along the Z axis

Side view

View along the Z axis

Figure 16. Two views of the orbitals of a dicoordinate (sp-hybridized) carbon atom

Two simple examples of dicoordinate carbon compounds are hydrogen cyanide and acetylene, both of which possess triple bonds and linear geometry because of the 180° angle between the sp orbitals of carbon (Figure 17).

acetylene hydrogen cyanide m t»

Figure 17. The two simplest dicoordinate carbon compounds, acetylene and hydrogen cyanide.

The triple bond consists of six electrons; two of these are in an axially symmetric MO formed from the combination of two sp AOs. The remaining 4 electrons are in two bonding 7t-MOs formed from overlap of the four 2p AOs.

The linear structure of acetylene follows from the use of the sp hybrid orbitals and p orbitals of each carbon and two H 1s orbitals to assemble the molecule, as shown in Figure 18a.

An electron cloud representation of the two n bonds of acetylene is shown in Figure 18b. These two rc-bonds and the sp-sp a bond of acetylene hold six electrons and constitute a C-C triple bond.

c-c o bond

a bond

it bond

C-H a bond ji bond

Figure 15b. rc-Bond of ethylene.

it bond

C-H a bond ji bond

it1 bond k2 bond

ti1 bond ji2 bond

Figure 18a. Formation of acetylene by the combination of two sp-hybridized carbon atoms and two hydrogen atoms.

it1 bond k2 bond ti1 bond ji2 bond

Figure 18a. Formation of acetylene by the combination of two sp-hybridized carbon atoms and two hydrogen atoms.

ji1 bond ji1 bond

k2 bond

Figure 18b. The two it-bonds of acetylene k2 bond

Figure 18b. The two it-bonds of acetylene

Rule #4: Dicoordinate carbon forms bonds to the two attached atoms in a colinear arrangement. The bonding involves two colinear sp orbitals and two p atomic orbitals at carbon.

Carbon dioxide (C02) is another linear molecule in which carbon is sp-hybridized (Figure 19).

o=c=o carbon dioxide

Figure 19. Structure and shape of carbon dioxide.

The Common Chemical Elements in Living Systems

Most of the common elements that make life possible fall within the first three rows of the Periodic Table of Elements. These are shown in Figure 20 along with the corresponding atomic numbers. The atomic number of an atom is identical to the number of protons in the nucleus or the number of orbiting electrons.



















10 Ne







16 S


18 Ar



20 Ca




Figure 20. A portion of the Periodic Table of Elements

Figure 20. A portion of the Periodic Table of Elements

Some Simple Compounds of Hydrogen and Non-Carbon Elements

All the elements shown in Figure 20 combine with hydrogen, with the exception of the inert gases He, Ne and Ar. Some examples of the simplest of these are the following (Figure 21).



Hydrogen Fluoride

Hydrogen Sulfide

Hydrogen Chloride

Figure 21. Simple compounds of hydrogen and noncarbon elements.

The dot pairs in the above structures represent electrons in the outer valence shell that are not needed in bonding. The structure of water, for example, involves an sp3 hybridized oxygen atom connected to two hydrogen atoms. The single bonds to the hydrogens use up two out of the six available electrons of an O atom. The remaining four oxygen electrons are located in the remaining (nonbonding) sp3 orbitals. Since the four sp orbitals are filled by eight electrons, no further electrons, for instance from a hydrogen atom (H ), can be added. (Reminder: each orbital can hold only two electrons.)

Carbon Bonding to Elements Other than Hydrogen

Carbon can also bond to most of the elements, for instance replacing hydrogen in the compounds shown in Figure 21.


Dimethyl Ether

Methyl Fluoride

HjC CH3 Dimethyl Sulfide

Methyl Chloride



Bromide vi

Rule #5: Carbon can bond to itself to form either straight or branched chains or rings. Carbon can also bond to many other atoms.

Ionic Bonds

The metallic elements at the left of the Periodic Table lose an electron very readily and tend to form positively-charged ions (cations), rather than covalently bonded compounds. The elements F, CI, Br and I at the right of the Periodic Table, in contrast, have high electron affinity and readily accept an electron to form negative ions (anions). In the case of fluoride ion (F") the available orbitals are filled, as with the inert gas neon (Ne) with which it is isoelectronic (i.e., both F" and Ne have a total of 8 outer shell electrons). The bonding between sodium and chlorine, for example, is essentially electrostatic, and the bond is described as ionic in character. It is the extreme of the covalent bond of H2 which involves no charge separation. Sodium chloride (NaCI, salt) in solid form is a crystalline structure containing Na+ and CI" ions in an indefinitely repeating lattice in which each Na+ is surrounded by 6 CI", and vice versa. It is so stable that the melting point of salt is about 800 °C. The energy that holds Na*cr in the crystal lattice is 187 kcal/mol, much greater than the H—H covalent bond energy (104 kcal/mol).

Bonds of Intermediate Polarity

Hydrogen chloride (HCI) is a gas at room temperature, in contrast to the ionic solid sodium chloride. The bonding in H-CI is best described as a covalent bond with appreciable (but far from full) ionic character or charge separation. The electron pair between H and CI is not shared equally. It is a polarized molecule with more electron density at the CI end and less at the H end (Figure 23).

Less electron density

Less electron density

More electron density

Figure 23. The electron density around HCl. In this computer-generated electron-density map, the blue color represents the lowest electron density whereas the red color represents the highest electron density.

More electron density

Figure 22. Simple compounds of carbon with noncarbon elements.

Figure 23. The electron density around HCl. In this computer-generated electron-density map, the blue color represents the lowest electron density whereas the red color represents the highest electron density.

In aqueous solution HCI ionizes to form a hydrated proton and a hydrated chloride ion. Thus, it is a strong acid.

Equation 2. Dissociation of HCI in water

The polarization of the bond in gaseous hydrogen chloride, often indicated using the notation Hs+—Cl°", is a consequence of greater electron affinity of a chlorine atom as compared to a hydrogen atom. Expressed in another way, chlorine is more strongly electron-attracting than hydrogen.

Polarization of covalent bonds is very common. Four examples are shown in Figure 24.


H H Water


H H Water


8* 8" H3C-CI Methyl Chloride




Figure 24. Electron density maps of four simple compounds. The highest electron density is shown in red whereas the lowest electron density is shown in blue.

primarily because of electrostatic forces, forming an extended three-dimensional network, a small part of which is shown in Figure 25.

Figure 25. An extended three-dimensional network of molecules in liquid water involves "hydrogen bonds" (blue dashes) as shown.

The bonds between molecules of H20 in the liquid - called "hydrogen bonds" - are much weaker than the H—H bond (104 kcal/mol), or the C—H bond in methane (105 kcal/mol). The bond dissociation energy of the covalent H—O bond in gaseous H20 is about 117 kcal/mol, whereas the energy of attraction of an H in water with an O atom of a neighboring water molecule is about 6 kcal/mol. Intermolecular hydrogen bonds between water molecules in the liquid are about 90% electrostatic and 10% covalent. The hydrogen bonds in liquid water stabilize it by an energy of cohesion that is responsible for the unusually high heat of vaporization (9.7 kcal/mol, or 538 kcal/liter).

Aqueous Solvation of Ions

Molecular Polarity and Hydrogen Bonding

Oil and water do not mix because neither can dissolve the other. The former is essentially hydrocarbon-like and nonpolar, whereas water is polarized with the oxygen relatively negative and the two hydrogens positive.

The polarity of the O—H bonds in water causes the boiling point of water (100 °C) to be much higher than that, for instance, of methane (CH4, -161 °C) which has about the same size, or ammonia (NH3, -33 °C). The O—H bond polarity of water causes molecules of H20 to associate with one another,

Another unique property of water is the existence of ionized species. In pure, neutral water, the concentration of the hydrated proton H30+(H20)n and hydrated hydroxide ion H0"(H20)n are each ca. 10"7 mol/liter. These species are essentially hydrogen-bonded clusters.

The polarity of water makes it a good solvent for polar ionic molecules because water can form electrostatic or hydrogen bonds to the dissolved species. For example:

Equation 3. Formation of hydrated ions in water

One of the simplest indications that solutions of NaCI or HCI in water contain ions is their high electrical conductivity. Pure water is only a weak conductor of electricity because the concentrations of the ions H30+ and HO" are only 10"7 mol/liter. Seawater is a much better conductor because it contains 0.1 mol/liter of Na+ and CI" ions.

atom, and the result is a transient attraction (Figure 26).

Solvation Energies in Water

Water is unique as a solvent.

Despite the fact that solid NaCI is bound in the crystal lattice with an energy of 187 kcal/mol, it dissolves in water to give solutions of hydrated Na+ and CI". The reason for this is that the energy of solvation of these ions in water (191 kcal/mol) overcomes the high lattice energy by 4 kcal/mol. The sum of energies of solvation by water of a proton (269 kcal/mol) and of a chloride ion (89 kcal/mol) are greater than the dissociation energy of gaseous hydrogen chloride (358 kcal/mol i/s 103 kcal/mol), and so it is clear why HCI is both soluble in water and fully ionized. The take-home lesson is that solvation by water strongly stabilizes both positive and negative ions.

The high solvation energy of sodium chloride in water is largely electrostatic. Solvated Na+ is surrounded by a cluster of at least six H20 molecules with oxygen in proximity to Na+. Solvation of CI similarly involves a cluster of H?0 molecules with hydrogen in proximity to CI".

Interactions Between Nonpolar Molecules

There are also attractive forces that operate between nonpolar molecules such as straight-chain hydrocarbons. These intermolecular attractions, sometimes called van der Waals forces, are very much weaker than covalent or ionic bonds, or even hydrogen bonds. An instructive example is the attraction between two atoms of the inert gas argon (Ar) which has been measured as ca. 0.28 kcal/mol. This attraction arises not from covalent bonding (because the atomic orbitals of Ar are filled), but from fluctuations in electron density around Ar that create transient imbalance, with one side of the atom being more negative than the other. This polarity induces an opposite distribution of charge on a nearby Ar

Figure 26. Electron density fluctuation in Ar(1) induces opposite electron density in Ar(2), leading to net attraction between them.

The attraction between two adjacent nonpolar molecules increases in proportion to the area of contact and is usually on the order of ca. 1 kcal per square A of close contact. One manifestation of van der Waals attraction is the steady increase in boiling point temperatures (in °C) with increasing molecular size for the series of straight-chain hydrocarbons (Figure 27).

Octane (126 X)

Decane (174 X)

Figure 27. Boiling points (red) of straight-chain hydrocarbons (blue).

Although van der Waals attractions between molecules are weaker than hydrogen bonding or electrostatic interactions, they can become significant when two nonpolar molecular surfaces are complimentary in shape of sizeable area. These forces play a major role in determining three-dimensional protein geometry and specificity of drug action. In addition, this type of interaction is what makes it possible for gecko lizards to walk across smooth ceilings or vertical walls.

Functional Groups, Subunits Within Structures that Confer Characteristic Properties and Reactivity.

Compounds containing a carbon-carbon double bond within the structure show characteristic chemical behavior. For instance, ethylene, 1-pentene and cyclohexene all react with H2 in the presence of finely divided nickel (Ni) as a catalyst to form products in which hydrogen has been added to the sp2 carbons (Figure 28). Such addition reactions to C=C are so common that these compounds are called unsaturated.

There are many other characteristic reactions of the C=C subunit, often called an olefinic functional group (or alkene). In addition, there are many other types of functional groups in organic compounds. A tabulation of some of the most common ones are shown in Figures 29a and 29b





Ni catalyst

Ni catalyst

Ni catalyst

H H ethane pentane


Figure 28. Reaction of compounds containing a C=C double bond with hydrogen gas (H2) in the presence of Ni catalyst.


c=c double bond


H2C = CH2 ethylene





■•■F —Cl ...Br ---l halogen acetone acetaldehyde

H3C-N02 nitromethane

Cl chloroform

Figure 29b. A small sample of functional groups in organic compounds.

Many of the compounds shown in Figures 29a and 29b are familiar to most people, especially ethanol (alcohol), butanethiol (butylmercaptan, the essence of the odor of skunks), and acetic acid (vinegar). They are also widely useful articles of commerce. For instance ethylene is the building block from which giant molecules of the plastic polyethylene are made.

Often, combinations of directly connected functional groups occur in molecules. Some examples of such combinations are shown in Figure 30.

•i^ci vinyl chloride

ethyl acetate (an ester)

O ii

-"^O-H carboxyl

amine hc=c-ch3 methylacetylene h2

h3c oh ethanol h2 h2 ,cx ,cv


butanethiol h3c o-h acetic acid methylamine

Figure 29a. A small sample of functional groups in organic compounds.

h2n nh2 urea

x nh2

acetamide (an amide)

^OH acrylic acid h3 S

dimethyl disulfide o o

CH/^0 CH3 acetic anhydride (an anhydride)

acrylonitrile (a nitrile)

tertiary-butyl hydroperoxide

Figure 30. Compounds containing a combination of functional groups shown in Figures 29a and 29b.

Functional groups containing sulfur and phosphorus can also exist in states of higher oxygenation since these elements have five 3d-orbitals available for bond formation in addition to the 3s and three 3p orbitals. Some examples of these more highly oxidized functional groups are shown in Figure 31.

ch3x ch3

dimethyl sulfoxide (a sulfoxide)

o dimethyl sulfone (a sulfone)

ch3ch2-o-p-o-ch2ch3 o diethylphosphoric acid (a phosphoric acid ester)

o o dimethylpyrophosphoric acid (a pyrophosphate)

Figure 32. Three different representations of the delocalization of electrons in formate ion In formula a, atomic orbitals combine to form two 3-center bonding MOs. In formula B, the delocalization over 0-C-0 is illustrated with dotted lines and the oxygen atoms share the negative charge equally. Formula C is the computer-generated electron density map of formate ion in which red indicates high electron density.

ch,-s-nh2 il o methanesulfonamide (a sulfonamide)

methanesulfonic acid (a sulfonic acid)

Figure 31. Compounds containing highly oxidized functional groups derived from sulfur and phosphorus.

Carboxylic Acids. Part I. Acidity

The classification of structural subunits as functional groups is a very useful technique for organizing the enormous amount of information that is encompassed by organic chemistry, the chemistry of the vast family of carbon compounds (carbogens). In this section the carboxylate functional group is examined in some detail using a few of the smaller molecules of the class (carboxylic acids) as representative.

Formic acid, the simplest carboxylic acid, is a proton donor. In aqueous solution it ionizes partially to hydrated negative formate ion and hydrated H30+, as Equation 4 indicates.

H-^oh formic acid

H-^oe formate ion (hydrated)


Equation 4. Dissociation of formic acid in water.

The reaction is rapid (submillisecond time scale) and reversible, i.e., the components are in very fast equilibrium with one another. One reason for the stability of formate ion is that the negative charge is spread equally between the two oxygens. The carbon is sp2 hybridized and there are two delocalized 7t-bonding MOs as expressed in formulas A, B and C in Figure 32.

Delocalization of electrons or charge in organic molecules is, in general, strongly stabilizing because it leads to a lower energy structure than the hypothetical electron-localized version(s).

Formic acid is the acidic ingredient that causes the immediate sting in the bite of a bee or hornet. It can be neutralized by a base such as ammonia (NH3) or sodium hydroxide (NaOH) to form formate salts. (See Figure 33.)

Figure 33. Two simple salts of formic acid, ammonium and sodium formate.

Carboxylic Acids. Part II. General Reactions and Derivatives

The conversion of a carboxylic acid to a carboxylate salt by treatment with a base is a general property of this functional group class. As might be expected, the salts are generally much more soluble in water than the corresponding carboxylic acids. Solutions of carboxylic acids in water are acidic, just as for formic acid, although the degree of acidity varies from one compound to another. Acetic acid, CH3COOH, is less acidic than formic acid. Trifluoroacetic acid, CF3COOH, is a much stronger acid than either formic or acetic acid. The reason for this is that fluorine is powerfully electron attracting, and considerable negative charge is delocalized to the three fluorines in CF3COO", conferring extra stabilization. Figure 34 shows the electron density maps for acetate, formate and trifluoroacetate.

Hj H




Figure 34. Electron density maps for acetate, formate and trifluoroacetate anions, displayed in order of increasing stability.



general formula for the class of a-amino acids

The subfragment RCO of carboxylic acids is called an acyl group (the R of RCO can be any carbon group, Figure 35).

R-^OH carboxylic acid o rax carboxylic acid derivative

Figure 35. Formation of carboxylic acid derivatives. The acyl group is highlighted with the yellow box.

There are many reactions of carboxylic acids that form compounds of structure RCOX which are called carboxylic acid derivatives. Some examples are shown in Figure 36.

r och3 methyl ester

O fl r sch3 methyl thiolester o r ci acyl chloride

x r nh2 amide a-amino acid o oxalic acid malonic acid oh

3=< oh o o^j oh oh citric acid oxalic acid, the acidic component of rhubarb malonic acid, a building block for the synthesis of fats in vivo citric acid, the acidic component of lemons, oranges and other fruits.

Figure 36. Examples of simple carboxylic acid derivatives.

Figure 37b. Polyfunctional carboxylic acids.

Sulfur and Phosphorus Acids and their Derivatives

These compounds can all be made by standard reactions from carboxylic acids (RCOOH).

Sulfuric acid is a strong acid that dissolves in water to form the hydrated negative ions (anions) bisulfate and sulfate.

Polyfunctional Carboxylic Acids

The essentially infinite diversity of organic compounds is made possible in part because the various functional groups can occur together in all possible numbers and combinations. Several important specific examples are shown in Figures 37a and 37b.



the simplest a-amino acid and one of the building blocks of proteins

H3N^iloe glycine

Figure 37a. Polyfunctional carboxylic acids.

sulfuric acid h'%

m e -s-o bisulfate ion

Equation 5. Dissociation of sulfuric acid in water

Similarly, sulfonic acids, such as methane-sulfonic acid, ionize completely in water and form salts with bases.

methanesulfonlc acid i e©

-S-ONa sodium methanesulfonate

Equation 6. Neutralization of methanesulfonic acid with aqueous sodium hydroxide to form sodium methanesulfonate.

Long chain sulfonate salts such as C18H37S03"Na+ (sodium octadecane sulfonate, see Figure 38) are important detergents (e.g., Tide) that interact by van der Waals forces to attract greasy deposits and carry them into a water wash because of the strong aqueous solvation of the sulfonate ion.

Phosphorylation of hydroxyl groups in proteins is important in living organisms as a key reaction in chemical signaling.

Benzene, Structure and Stabilization by n-Delocalization

Figure 38. Line drawing and space-filling representation of sodium octadecane sulfonate (Ci8H37S03Na).

The negative charge in bisulfate, sulfate and sulfonate ions is spread out over oxygens in delocalized and highly stabilized molecular orbitals.

Benzene is a remarkably stable hydrocarbon of formula C6H6. The six carbons are held together in a planar hexagonal arrangement as shown in Figure 41. That geometry corresponds to sp2 hybridization of the carbons in the ring.

7t-bonds localized rc-bonds delocalized

7t-bonds localized

rc-bonds delocalized

Figure 41. Structure of benzene.

space-filling model

Bisulfate ion

Sulfate ion Methanesulfonate ion

Figure 39. Computer-generated electron density maps of bisulfate, sulfate and methanesulfonate Ions.

The behavior of phosphoric acid is similar. Phosphoric acid reacts with sodium hydroxide in water to form soluble mono-, di- or trisodium salts.

phosphoric acid m e©


monosodium phosphate


eONa disodium phosphate


eONa trisodium phosphate

Equation 7. Neutralization of phosphoric acid with sodium hydroxide to form phosphate salts.

Phosphate esters can be formed from the attachment of phosphate to a hydroxyl group in an organic molecule (e.g., alcohols), as shown in the examples that follow.

Although analogy with ethylene (see page 8) might suggest structure A for benzene, it has been shown experimentally that the six CC bonds in benzene are equal in length and that it is a delocalized jr-electron structure. Quantum chemical calculations show that the six electrons not used for C-H and C-C o-bonding lie in three bonding orbitals formed by the overlap of the p AOs on each carbon. A MO energy diagram is shown in Figure 42.


Epz AO

diethylphosphate phosphoglyceric acid Figure 40. Simple esters of phosphoric acid.

Figure 42. I: Energy levels of the six 7t-MOs of benzene that result from the combination of the six 2pz atomic orbitals of the six carbons in the ring. II: An MO picture of the 7r-electron clouds above and below the benzene ring. Ill: A picture of the 7t-electron clouds above and below the benzene ring

Each carbon contributes a p-AO to form six rc-MOs. Three of these will be of lower energy than the original p2 atomic orbitals and three will be higher (see diagram I. in Figure 42). There are six electrons available for it-bonding. Since these will be accommodated in the lower energy orbitals MO (1) and MO (2) and MO (2'), the resulting 7r-delocalized structure will be stabilized. In fact, simple calculations show that the delocalized structure B is more stable than the localized three C=C structure A (see Figure 41) by about 35 kcal/mol.

Despite the fact that structure A is less realistic than structure B in Figure 41, it is commonplace to draw the structure of benzene and related compounds with alternating double bonds. The structures of benzenoid compounds in the later sections of this book are drawn with the traditional double/single bond notation.

Chemical Consequences of the Stability of the Benzene Ring

The derealization of the six rc-electrons of benzene over the whole ring in low-energy orbitals bestows special properties and huge importance to this structural unit. Typically the C=C unit tends to undergo reactions in which groups add to the two carbons of the double bond. However, benzene generally reacts with these same reagents differently. For instance, 1-pentene reacts with chlorine gas (Cb) to form 1,2-dichloropentane, but benzene reacts to form chlorobenzene and HCI as shown in Figure 43.




1,2-dichloro-pentane chlorobenzene Figure 43. Reaction of 1-pentene and benzene with chlorine gas (ci2). In the case of benzene one of the hydrogen atoms (H) is replaced with a chlorine (CI) atom - a process which is known as substitution.

Most reactions of benzene are of the substitution type (replacement of atoms) with retention of the very stable it-system, whereas a double bond in a typical non-benzenoid hydrocarbon undergoes addition reactions.

There are innumerable chemical reactions that replace the hydrogens of benzene by some other groups, and countless thousands of substituted benzene compounds can be made (synthesized). The six carbons of the benzene ring allow different positioning of groups. For instance, there are three distinctly different dichlorobenzenes, as indicated in Figure 44.







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