an introduction to environmental organic chemicalsmrh/ese 176 sgi book/sgi chp 2 introduction env...

13

Chapter 2

AN INTRODUCTION TO ENVIRONMENTAL ORGANIC CHEMICALS

2.1 Introduction

2.2 The Makeup of Organic Compounds Elemental Composition, Molecular Formula, and Molar Mass Electron Shells of Elements Present in Organic Compounds Covalent Bonding Bond Energies (Enthalpies) and Bond Lengths. The Concept of Electro-

Oxidation State of the Atoms in an Organic Molecule Illustrative Example 2.1 : Determining the Oxidation States of the Carbon

The Spatial Arrangement of the Atoms in Organic Molecules Delocalized Electrons, Resonance, and Aromaticity

negativity

Atoms Present in Organic Molecules

2.3 Classification, Nomenclature, and Examples of Environmental Organic Chemicals The Carbon Skeleton of Organic Compounds: Saturated,

Organohalogens Oxygen-Containing Functional Groups Nitrogen-Containing Functional Groups Sulfur-Containing Functional Groups Phosphorus-Containing Functional Groups Some Additional Examples of Compounds Exhibiting More Complex

Unsaturated, and Aromatic Hydrocarbons

Structures

2.4 Questions and Problems

Environmental Organic Chemistry, 2nd Edition. Rene P. Schwarzenbach, Philip M. Gschwend and Dieter M. Imboden

Copyright 0 2003 John Wiley &L Sons, Inc.

14 An Introduction to Environmental Organic Chemicals

When confronted with the plethora of natural and man-made organic chemicals released into the environment (Blumer, 1975; Stumm et al., 1983), many of us may feel overwhelmed. How can we ever hope to assess all of the things that happen to each of the substances in this menagerie, encompassing so many compound names, formulas, properties, and reactivities? It is the premise of our discussions in this book that each chemical’s structure, which dictates that compound’s “personality,” provides a systematic basis with which to understand and predict chemical behavior in the environment. Thus in order to quantify the dynamics of organic compounds in the macroscopic world (Parts IV and V), we will need to learn to visualize organic molecules in the microscopic environments in which they exist. However, before we do that in Parts I1 and 111, we first need to refresh our memories with some of the terminology and basic chemical concepts of organic chemistry used throughout this book (Section 2.2). For readers with little background in organic chemistry, it may be useful to consult the introductory chapters of an organic chemistry textbook in addition to this section. On the other hand, professional chemists might want to continue directly with Section 2.3, where we will try to give an overview of some important groups of environmentally relevant organic chemicals.

The Makeup of Organic Compounds

To understand the nature and reactivity of organic molecules, we first consider the pieces of which organic molecules are made. This involves both the various atoms and the chemical bonds linking them. First, we note that most of the millions of known natural and synthetic (man-made) organic compounds are combinations of only a few elements, namely carbon (C), hydrogen (H), oxygen (0), nitrogen (N), sulfur (S), phosphorus (P), and the halogens fluorine (F), chlorine (Cl), bromine (Br), and iodine (I). The chief reason for the almost unlimited number of stable organic molecules that can be built from these few elements is the ability of carbon to form up to four stable carbon-carbon bonds. This permits all kinds of three-dimensional carbon skeletons to be made, even when the carbon atoms are also bound to heteroatoms (i.e., to elements other than carbon and hydrogen). Fortunately, despite the extremely large number of existing organic chemicals, knowledge of a few governing rules about the nature of the elements and chemical bonds present in organic molecules will already enable us to understand important relationships between the structure of a given compound and its properties and reactivities. These properties and reactivities determine the compound’s behavior in the environment.

Elemental Composition, Molecular Formula, and Molar Mass

When describing a compound, we have to specify which elements it contains. This information is given by the elemental composition of the compound. For example, a chlorinated hydrocarbon, as the name implies, consists of chlorine, hydrogen, and carbon. The next question we then have to address is how many atoms of each of these elements are present in one molecule. The answer to that question is given by

The Makeup of Organic Compounds 15

the molecular formula, for example, four carbon atoms, nine hydrogen atoms, and one chlorine atom: C4H9C1. The molecular formula allows us to calculate the molecular mass of the compound, which is the sum of the masses of all atoms present in the molecule. The atomic masses of the elements of interest to us are given in Table 2.1. Note that for many of the elements, there exist naturally occurring stable isotopes. Isotopes are atoms that have the same number of protons and electrons (which determines their chemical nature), but different numbers of neutrons in the nucleus, thus giving rise to different atomic masses. Examples of elements exhibiting isotopes that have a significant natural abundance are carbon (I3C:l2C = 0.01 1: l), sulfur (34S:32S = 0.044:1), chlorine (37C1:35C1 = 0.32:1), and bromine (s1Br:79Br = 0.98: I). Conse- quently, the atomic masses given in Table 2,l represent averaged values of the naturally occurring isotopes of a given element (e.g., average carbon is 1.1% at 13 u + 98.9% at 12 u = 12.011 u). Note that 1 u (unified atomic mass unit) is approximately equal to 1.6605 x kg. Using these atomic mass values we obtain a molecular mass (or molecular weight) of 92.57 u for one single molecule with the molecular formula C4H9C1. If we take the amount of 1 mole of pure substance (that is, the Avogadro’s number NA = 6.02 x loz3 identical units, here molecules), this amount weighs 92.57 grams and is named the molar muss of the (pure) substance. So, 1 mole (abbreviation 1 mol) of any (pure) substance always contains the same amount of molecules and its mass in grams is what each molecule’s mass is in u.

Given the molecular formula, we now have to describe how the different atoms are connected to each other. The description of the exact connection of the various atoms is commonly referred to as the structure of the compound. Depending on the number and types of atoms, there may be many different ways to interconnect a given set of atoms which yield different structures. Such related compounds are referred to as isomers. Furthermore, as we will discuss later, there may be several compounds whose atoms are connected in exactly the same order (i.e., they exhibit the same structure), but their spatial arrangement differs. Such compounds are then called stereoisomers. It should be pointed out, however, that, quite often, and particularly in German-speaking areas, the term structure is also used to denote both the connectivity (i.e., the way the atoms are connected to each other) as well as the spatial arrangement of the atoms. The term constitution of a compound is then sometimes introduced to describe solely the connectivity.

Electron Shells of Elements Present in Organic Compounds

Before we can examine how many different structures exist with a given molecular formula (e.g., C,H9C1), we have to recall some of the rules concerning the number and nature of bonds that each of the various elements present in organic molecules may form. To this end, we first examine the electronic characteristics of the atoms involved.

Both theory and experiment indicate that the electronic structures of the noble gases [helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe), and radon (Rn)] are especially nonreactive; these atoms are said to contain “filled shells” (Table 2.1). Much of the chemistry of the elements present in organic molecules is understandable in terms of a simple model describing the tendencies of the atoms to attain such filled-shell conditions by gaining, losing, or, most importantly, sharing electrons.

Tab

le 2.1

Ato

mic

Mas

s, E

lect

roni

c C

onfi

gura

tion,

and

Typ

ical

Num

ber o

f C

oval

ent B

onds

of t

he M

ost I

mpo

rtan

t Ele

men

ts P

rese

nt in

Org

anic

Mol

ecul

es

Elem

ent'

Num

ber o

f Ele

ctro

ns in

She

ll N

umbe

r of

Cov

alen

t Bon

ds

Mas

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et C

harg

e C

omm

only

Occ

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ng in

N

ame

Sym

bol

Num

ber

(4

K

L

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N

0

ofK

erne

l O

rgan

ic M

olec

ules

Hyd

roge

n H

eliu

m

H He

1 2 1.

008

1 -

2 1+

0

1

Car

bon

Nitr

ogen

Oxy

gen

Fluo

rine

Neo

n

C

N

0

F -

Ne

6 7 8 9 10

12.0

1 1

14.0

07

15.9

99

18.9

98

2 4+

5+

6+

7+

0 2

Phos

phor

us

Sulf

ur

Chl

orin

e

Arg

on

P S c1

-

Ar

15

16

17

18

30.9

74

32.0

6

35.4

53

2 2

5+

6+

7+

0

2 2

Bro

min

e

Kry

pton

Br

-

Kr

35

36

79.9

04

2 8 8

18

18

7 - 8

7+

0

1

2

I -

Xe

53

54

126.

905

2 2 8 8

18

18

18

7 7+

18

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.


The first shell (K-shell) holds only two electrons [helium structure; first row of the periodic system (P.s.) of the elements] ; the second (L-shell; second row of the P.s.) holds eight; the third (M-shell; third row of the P.s.) can ultimately hold 18, but a stable configuration is reached when the shell is filled with eight electrons (argon structure). Thus, among the elements present in organic molecules, hydrogen requires two electrons to fill its outer shell (one it supplies, the other it must get somewhere else), while the other important atoms of organic chemistry require eight, that is, an octet configuration (see Table 2.1). It is important to realize that the number of electrons supplied by a particular atom in its outer shell (the so-called valence electrons; note that the remainder of the atom is referred to as kernel) chiefly determines the chemical nature of an element, although some significant differences between elements exhibiting the same number of outer-shell electrons do exist. The latter is due in large part to the different energetic status of the electrons in the various shells, reflecting the distance of the electrons from the positively charged nucleus. We address the differences between such elements (e.g., nitrogen and phosphorus, oxygen and sulfur) at various stages during our discussions.

Covalent Bonding

The means by which the atoms in organic molecules customarily complete their outer-shell or valence-shell octet is by sharing electrons with other atoms, thus forming so-called covalent bonds. Each single covalent bond is composed of a pair of electrons, in most cases one electron contributed by each of the two bonded atoms. The covalent bond may thus be characterized as a mutual deception in which each atom, though contributing only one electron to the bond, “feels” it has both electrons in its effort to fill its outer shell. Thus we visualize the bonds in an organic compound structure as electron pairs localized between two positive atomic nuclei; the electrostatic attraction of these nuclei to these electrons holds the atoms together. The simple physical law of the attraction of unlike charges and the repulsion of like charges is the most basic force in chemistry, and it will help us to explain many chemical phenomena.

Using the simple concept of electron sharing to complete the valence-shell octet, we can now easily deduce from Table 2.1 that H, F, C1, Br, and I should form one bond (monovalent atoms), 0 and S two (bivalent atoms), N and P three (trivalent atoms), and C four (tetravalent atom) bonds in a neutral organic molecule. With a few exceptions (particularly for S and P), which we will address later in Section 2.3, this concept is valid for the majority of cases that are of interest to us. For our compound with the molecular formula of C,H,Cl, we are now ready to draw all the possible structural isomers by simply applying these valency rules. Fig. 2.1 shows that there are four different possibilities. With this example we also take the opportunity to get acquainted with some of the common conventions used to symbolize molecular structures. Since it is clearer to separate shared and unshared electron pairs, the former (the actual covalent bond) are written as straight lines connecting two atomic symbols, while the unshared valence electrons are represented by pairs of dots (line 1 in Fig. 2.1). This representation clearly shows the nuclei and all of the electrons we must visualize. To simplify the drawing, all lines indicating bonds to hydrogen as well as the dots for unshared (nonbonding) electrons are frequently not shown (line 2). For further convenience we may, in many cases, eliminate all the bond lines without


nonbonding electrons

H :El: H I l l

H H H H H H H I I I I 1 I I . .

I I I I H I H

I I I I H I H

H H H H 1

H H H H I

I I I I . . I I I I -

(1) H-C-C-C- C-Cl: H- C- C- C- C-H H- C-Cc-Cc-Cl: H-C-C- e H

H- F;-H H H :Cl: H

H- C-H bonding electrons

I H

H,C- CH, - CH- CH, H,C- CH- CH, - CI I I CI CH,

(2) H,C-CH, - CH, - CH, - CI

I H

CI I

H3C- C- CH, I CH,

(3) H,CCH,CH,CH,CI H,CCH,CH( CI)CH, H,CCH( CH,)CH,CI (H3C)3CCl

Figure 2.1 Conventions for sym- bolizing the molecular structures of the four butyl chloride (or chlo- robutane) isomers.

loss of clarity, as illustrated in line 3 . Note that branching is indicated by using parentheses in this case. Finally, especially when dealing with compounds exhibiting a large number of carbon atoms, it is very convenient to just sketch the carbon skeleton as depicted in line 4. Each line is thus a skeletal bond and is assumed to have carbons at each end unless another element is shown. Furthermore, no carbon- hydrogen bonds are indicated, but are assumed present as required to make up full bonding (four bonds) at each carbon atom. To distinguish the various carbon-carbon bonds, bond lines are placed (whenever possible), at about 120°, roughly resembling the true physical bond angle (see below).

At this point in our discussion about chemical bonds and structural formulas, we should stress that structural isomers may exhibit very different properties and reactivities. For example, the rates of hydrolysis (reaction with water, see Chapter 13) of the four butyl chlorides shown in Fig. 2.1 are quite different. While the hydrolytic half-life (time required for the concentration to drop by a factor of 2) of the first and third compound is about 1 year at 25”C, it is approximately 1 month for the second compound, and only 30 seconds for the fourth compound. When we compare the two possible structural isomers with the molecular formula C,H60, we can again find distinct differences in that the well-known ethanol (CH,CH,OH) is a liquid at ambient conditions while dimethylether (CH,OCH,) is a gas. These examples should remind us that differences in the arrangement of a single collection of atoms may mean very different environmental behavior; thus we must learn what it is about compound structure that dictates such differences.

So far we have dealt only with single bonding between two atoms. There are, however, many cases in which atoms with more than one “missing” electron in their outer shell form double bonds or, sometimes, even triple bonds; that is, two atoms


?-? CI CI

tetrachloroethene

A acetone

0 cyclohexene

furan

KCEN acrylonitrile

Figure 2.2 Some simple molecules exhibiting double and/or triple bonds. Note that we use notation type 4 in Fig. 2.1.

share either two or even three pairs of electrons to complete their valence shells. A few examples of compounds exhibiting double or triple bonds are given in Fig. 2.2. We note that a double bond is indicated by a double line and, logically, a triple bond by three parallel lines between the corresponding atoms. We also note from Fig. 2.2 that there are compounds with ring structures (that may or may not contain double bonds). Rings are usually composed predominantly of carbon atoms, but they may also contain heteroatoms (i.e., elements other than carbon or hydrogen such as 0, N) in the ring (see Section 2.3).

Bond Energies (Enthalpies) and Bond Lengths. The Concept of Electronegativity

An important aspect of chemical bonding that we need to address is the strength of a chemical bond in organic molecules; that is, we should have a general idea of the energy involved in holding atoms together in a covalent bond. The most convenient measure of bond energy is indicated by the bond dissociation enthalpy, AHIB. For a diatomic molecule, this is defined as the heat change of the gas phase reaction:

A-B+A*+B'; MiB at constant pressure and temperature [e.g., 1.013 bar (= 1 atm) and 25"CI. Here MiB also contains the differences in translational, rotational (only AB), and vibrational (only AB) energies between educt (A-B) andproducts (A*,B*). Unfortunate- ly, it is not possible to directly measure bond dissociation (or formation) enthalpies for each of the different bonds present in a molecule containing more than one bond; they have to be determined indirectly, commonly through thermochemical studies of evolved heat (calorimetric measurements) in reactions such as combustion. These studies yield only enthalpies of overall reactions, where several bonds are broken and formed, respectively. The individual bond dissociation (or formation) enthalpies have then to be deduced from this data in various ways. The results are commonly shown in tables as average strengths for a particular type of bond, valid for gas phase reactions at 25°C and 1.013 bar. Table 2.2 summarizes average bond enthalpies (and bond lengths) of some important covalent bonds. From these data, some general conclusions about covalent bonds can be drawn, and a very useful concept can be derived to evaluate the uneven distribution of the electrons in a chemical bond: the concept of electronegativity.

Electronegativity. When visualizing a chemical bond, it is appropriate to imagine that the "electron cloud" or averaged electron position located between the two nuclei is, in general, distorted toward the atom that has the higher attraction for the electrons, that is, the atom that is more electronegutive. This results in the accumulation of negative charge at one end of the bond (denoted as 6 - ) and a corresponding deficiency at the other end (denoted as 6 + ):

Among the elements present in organic molecules, we intuitively (and correctly) predict that the smaller the atom (hence allowing a closer approach of the bonding


Table 2.2 Average Bond Lengths (A) and Average Bond Enthalpies ( kJ. mol-') of Some Important Covalent Bonds"

Bond LengthEnthalpy Bond LengthEnthalpy Bond LengthEnthalpy

Diatomic Molecules

H-H 0.741436 F-F 1.421155 o=o H-F 0.9215 66 C1-C1 1.991243 NsN

H-Cl 1.271432 Br-Br 2.281193

H-Br 1.411367 1-1 2.671152

H-I 1.601298

Covalent Bonds in Organic Molecules

Single bondsb

H-C 1.1 11415 c-c 1.541348 C-F

H-N 1.001390 C-N 1.471306 c-Cl

H-0 0.961465 c-0 1.411360 C-Br

H-S 1.331348 c-s 1.811275 c-I

Double and triple bonds

c=c 1.3416 12 C=Od 1.201737 c=c C=N 1.281608 C=O" 1.201750 C-N

c=s 1.561536 C=@ 1.161804

1.211498

1.101946

1.381486

1.781339

1.94128 1

2.141216

1.161838

1.161888

" Bond lengthhond enthalpy. Note that 1 8, equals 0.1 nm. which none of the partner atoms is involved in a double or triple bond. In such cases bond lengths are somewhat shorter. In carbon disulfide. In aldehydes. " In ketones. carbon dioxide.

Bond lengths are given for bonds in

electrons to the positively charged nucleus) and the higher the net charge of the kernel (nucleus plus the electrons of the inner, filled shells; see Table 2. l), the greater will be that atom's tendency to attract additional electrons. Hence, as indicated in Table 2.3, within a row in the Periodic Table (e.g., from C to F), electronegativity increases with increasing kernel charge, and within a column (e.g., from F to I), electronegativity decreases with increasing kernel size. The most commonly used quantitative scale to express electronegativity (Table 2.3) has been devised by Pauling (1960). On this scale, a value of 4.0 is arbitrarily assigned to the most electronegative atom, fluorine, and a value of 1 .O to lithium. The difference in electronegativity between two atoms A and B is calculated from the extra bond energy in A-B versus the mean bond energies of A-A and B-B in which the electrons should be equally shared. The reason for deriving relative electronegativities based on bond energies is that we interpret the extra bond strength in such a polarized bond to be due to the attraction of the partial positive and negative charges.

Let us discuss the importance of charge separation in bonds involving atoms of different electronegativity, for example, C and N, 0, or C1. The extent of partial ionic character in suchpolar covalent bonds is a key factor in determining a compound's


Table 2.3 Electronegativities of Atoms According to the Scale Devised by Pauling ( 1960)

ChargeofKernel: + I +4 + 5 + 6 + 7

H Increasing Size 2.2 of Kernel

C N 0 F 2.5 3.0 3.5 4.0

P S C1 2.2 2.5 3.0

Br 2.8

I 2.5

behavior and reactivity in the environment. The polarization in bonds is important in directing the course of chemical reactions in which either these bonds themselves or other bonds in the vicinity are broken. Furthermore, the partial charge separation makes each bond between dissimilar atoms a dipole. The (vector) sum of all bond dipoles in a structure yields the total dipole moment of the molecule, an entity that can be measured. However, it is the dipole moments of individual bonds that are most important with respect to the interactions of a given compound with its molecular surroundings.

From Table 2.3 it can be seen that according to Pauling’s scale, carbon is slightly more electron-attracting than hydrogen. It should be noted, however, that the electron- attracting power of an atom in isolation differs from that attached to electron-attracting or electron-donating substituents in an organic molecule. For example, many experimental observations indicate that carbon in -CH, is significantly less electron-attracting than hydrogen. We may rationalize this by recognizing that each additional hydrogen contributes some electron density to the carbon and successively reduces that central atom’s electronegativity. In conclusion, we should be aware that the electronegativity values in Table 2.3 represent only a rough scale of the relative electron-attracting power of the elements. Hence in bonds between atoms of similar electronegativity, the direction and extent of polarization will also depend on the type of substitution at the two atoms.

Hydrogen Bonding. One special result of the polarization of bonds to hydrogen which we should highlight at this point is the so-called hydrogen bonding. As indicated in Table 2.1 , hydrogen does not possess any inner electrons isolating its nucleus (consisting of just one proton) from the bonding electrons. Thus, in bonds of hydrogen with highly electronegative atoms, the bonding electrons are drawn strongly to the electronegative atom, leaving the proton exposed at the outer end of the covalent bond. This relatively bare proton can now attract another electron-rich


centeryespecially heteroatoms with nonbonding electrons, and form a hydrogen bond as schematically indicated below by the dotted line:

-X&-H&...:Y”- X ,Y=N,O, ...

In organic molecules this is primarily the case if X and Y represent nitrogen or oxygen.

If the electron-rich center forms part of the same molecule, one speaks of an intramolecular hydrogen bond; if the association involves two different molecules, it is referred to as an intermolecular hydrogen bond. Although, compared to covalent bonds, such hydrogen bonds are relatively weak (1 5 to 20 kJ. mol-’), they are of enormous importance with respect to the spatial arrangements and interactions of molecules.

We now return to Table 2.2 to note a few simple generalities about bond lengths and bond strengths in organic molecules. As one can see, bond lengths of first-row elements (C, N, 0, F) with hydrogen are all around 1 A (0.1 nm). Bonds involving larger atoms (S, P, C1, Br, I) are longer and weaker. Finally, double and triple bonds are shorter and stronger than the corresponding single bonds; and we notice that the bond enthalpies of double and triple bonds are often somewhat less than twice and three times, respectively, the values of the single atom bonds (important exception: C=O bonds).

To get an appreciation of the magnitude of bond energies, it is illustrative to compare bond enthalpies to the energy of molecular motion (translational, vibrational, and rotational), which, at room temperature, is typically on the order of a few tens of kilojoules per mole. As can be seen from Table 2.2, most bond energies in organic molecules are much larger than this, and, therefore, organic compounds are, in general, stable to thermal disruption at ambient temperatures. At high temperatures, however, the energy of intramolecular motion increases and can then exceed certain bond energies. This leads to a thermally induced disruption of bonds, a process that is commonly referred to as pyrolysis (heat splitting).

It is important to realize that the persistence of organic compounds in the environment is due to the relatively high energy (of activation) needed to break bonds and not because the atoms in a given molecule are present in their lowest possible energetic state (and, therefore, would not react with other chemical species). Hence, many organic compounds are nonreactive for kinetic, not thermodynamic, reasons. We will discuss the energetics and kinetics of chemical reactions in detail later in this book. Here a simple example helps to illustrate this point. From daily experience we know that heat can be gained from burning natural gas, gasoline, fuel oil, or wood. As we also know, all these fuels are virtually inert under environmental conditions until we light a match, and then provide the necessary initial activation energy to break bonds. Once the reaction has started, enough heat is liberated to keep it going. The amount of heat liberated can be estimated from the bond enthalpies given in Table 2.2. For example, when burning methane gas in a stove, the process that occurs is the reaction of the hydrocarbon, methane, with oxygen to yield CO, and H,O:


In this gas phase reaction we break four C-H and two 0=0 “double” bonds and we make two C=O and four 0-H bonds. Hence, we have to invest (4 x 41 5) + (2 x 498) = +2656 kJ-mol-’, and we gain (2 x 804) + (4 x 465) = -3468 kJ.mol-I. The estimated heat of reaction at 25OC for reaction 2-2 is therefore, -8 12 kJ. mol-’ (the experimental value is -802 kT.mol-’), a quite impressive amount of energy. We recall from basic chemistry that, by convention, we use a minus sign to indicate that the reaction is exothermic; that is, heat is given off to the outside. A positive sign is assigned to the heat of reaction if the reaction consumes heat by taking energy into the product structures; such reactions are then called endothermic.

The example illustrates that enthalpy can be gained when nonpolar bonds, as commonly encountered in organic molecules, are broken and polar bonds, such as those in carbon dioxide and water, are formed. Reactions which involve the transfer of electrons between different chemical species are generally referred to as redox reactions. Such reactions form the basis for the energy production of all organisms. From this point of view we can consider organic compounds as energy sources.

Oxidation State of the Atoms in an Organic Molecule

When dealing with transformation reactions, it is important to know whether electrons have been transferred between the reactants. For evaluating the number of electrons transferred, it is convenient to examine the (formal) oxidation states of all atoms involved in the reaction. Of particular interest to us will be the oxidation state of carbon, nitrogen, and sulfur in a given organic molecule, since these are the elements most frequently involved in organic redox reactions.

The terms oxidation and reduction refer, respectively, to the loss and gain of electrons at an atom or ion. An oxidation state of zero is assigned to the uncharged element; a loss of Z electrons is then an oxidation to an oxidation state of + Z. Similarly, a gain of electrons leads to an oxidation state lower by an amount equal to the number of gained electrons. A simple example is the oxidation of sodium by chlorine, resulting in the formation of sodium chloride:

oxidation NaO - Na+’ + e-

reduction .. :cI.O + e- - : CI: -1

To bridge the gap between full electron transfer in ionic redox reactions (as shown in the example above) and the situation with the shared electrons encountered in covalent bonds, one can formally assign the possession of the electron pair in a covalent bond to the more electronegative atom of the two bonded atoms. By doing so, one can then count the electrons on each atom as one would with simple inorganic ions. For any atom in an organic molecule, the oxidation state may be computed by adding 0 for each bond to an identical atom; - 1 for each bond to a less electronegative atom or for each negative charge on the atom; and + 1 for each bond to a more electronegative atom or for each positive charge. We note that in C-S, C-I, and even C-P bonds the


electrons are attributed to the heteroatom, although the electronegativities of these heteroatoms are very similar to that of carbon. Finally, we should also point out that Roman instead of Arabic numbers are frequently used to express the oxidation state of a covalently bound atom.

Some examples demonstrating how to determine the oxidation states of the carbon atoms present in organic molecules are given in Illustrative Example 2.1. More examples covering the other elements that may assume several different oxidation states (i.e., N, S, P) will be given in Section 2.3 in our discussion of functional groups.

~~ ~~

Illustrative Example 2.1

2-methyl-butane (iso-pentane)

1 /p H,C- C 2

\ 0-H

acetic acid

CI,i 2 ,H ,c =c

CI ‘ CI

trichloroethene

H - 0 8 \ / 4 CH3

5 H H

4-methyl-phenol (pcresol)

Determining the Oxidation States of the Carbon Atoms Present in Organic Molecules

Problem Determine the oxidation state of each carbon present in (a) iso-pentane, (b) acetic acid, (c) trichloroethene, and (d) 4-methylphenol (p-cresol). Note that in organic molecules hydrogen always assumes an oxidation state of +I, chloride of -I, and oxygen (in most cases, see Section 2.3) of -11.

Answer (a) The carbons of the methyl groups (Cl, C4, C,) are bound to three hydrogens and one carbon; hence their oxydation state is 3 (-1) + (0) = -111. The methylene group (C,) is bound to two hydrogens and two carbons, which yields 2(- I) +2(0) = -11. Finally, the methene group (C,) exhibits an oxidation state of (-1) + 3(0) = -I.

Answer (b) As in example (a) the oxidation state of the carbon of the methyl group (C,) is -111, while the oxidation state of the carboxylic carbon (C,) is (+II) + (+I) + (0) = +III. Hence the “average oxidation state” of carbon in acetic acid is 0.

Answer (c) In trichloroethene the oxidation states’ of the two carbons are 2(+I) + 0 = +I1 for (C,) and (-1) + (+I) + 0 = 0 for (C,).

Answer (d) The carbons present in the benzene ring exhibit oxidation states of (+I) + 2(0) = +I for (C,), (-1) + 2(0) =-I for (C,, C,, C5, C,), and 3 (0) = 0 for (C,); respectively. The oxidation state of the methyl carbon, is again -111.


Figure 2.3 Examples of bond angles in some simple molecules (from Hendrickson et al., 1970, and March 1992).

H

H

:S .. Go

H..&H

H 106.7"

CH3 , A - 1 1.7"

CH3

H3C 121"

The Spatial Arrangement of the Atoms in Organic Molecules

To describe the steric arrangement of the atoms in a molecule, in addition to bond lengths we need to know something about the angles between the bonds, the sizes of the atoms, and their freedom to move within the molecule, (e.g., rotations about bonds).

Bond Angles. A simple but very effective rule that we may apply when considering bond angles in molecules is that the electrons accept the closeness to one another because of pairing, but that each pair of electrons, shared (i.e., involved in a chemical bond) or unshared, wants to stay as far as possible from the other pairs of electrons [for details see valence shell electron pair repulsion (VSEPR) theory; Pfennig and Frock, 19991. This means that in the case of a carbon atom with four single bonds, the bonds will generally point toward the comers of a tetrahedron. In the symmetrical case, that is, when a carbon is bound to four identical substituents [i.e., atoms or groups of atoms as -H in CH,, or -Cl in CCl,, or -CH3 in C(CH3),], the bond angles are 109.5". In most cases, however, each carbon atom is bound to different substituents,


which leads to minor variations in the bond angles, as illustrated by some examples given in Fig. 2.3. For saturated carbon atoms, that is, carbon atoms not involved in a double or triple bond, the C-C-C bond angles are typically about 112”, except for ring systems containing less than six ring atoms, where bond angles may be considerably smaller. With respect to the heteroatoms N, 0, P, and S, we see from the examples given in Fig. 2.3 that the nonbonded electron pairs behave as if they point wiz to imaginary substituents, thus giving rise to bent or pyramidal geometry (provided that the heteroatoms are also only single-bonded to other atoms).

Stereoisomerism. We should note that the association of electrons in a single, or sigma (ci), bond allows rotation about the axis of the linkage (Fig. 2.4). Such rotation does not disrupt the bonding electron pair (i.e., it does not break the bond), and

%:X therefore under ambient temperatures the substituents attached to two carbons bonded by a sigma bond are usually not “frozen” in position with respect to one another. Thus the spatial arrangement of groups of atoms connected by such a single bond may change from time to time owing to such rotation, but such geometric distributions of Figure 2.4 Rotation about a D-

bond leading to various spatial ar- the atoms in the structure are usually not separable from one another since intercon- rangements of the atoms in a mole- versions occur during separation. However, as discussed below, even if fast rotations

about a single bond occur, stereoisomerism is possible. Stereoisomers are com- cule.

pounds made up of the same atoms bonded by the same sequence of bonds, but having different three-dimensional stvuctures that are not superimposable.

/

COOH

H- k- cH3

When considering stereoisomerism, one commonly distinguishes two different cases. First, there are molecules that are alike in every respect but one: they are mirror images of each other that are not superimposable. We refer to such molecules as being chiral. Note that, in general, any object for which the image and mirror image are distinguishable (e.g., our left and right hands), is denoted to be chiral. For example, if in a molecule a carbon atom is bound to four different substituents (as is the case in the herbicide mecoprop; Fig. 2.5), two structural isomers are possible. In this context one sometimes refers to such a carbon atom as a center of chirality. Mirror- image isomers are called enantiomers or optical isomers (because they rotate the plane of polarized light in opposite directions). In general, we may say that enantiomers have identical properties in a symmetrical molecular environment, but

H-&CooH that their behavior may differ quite significantly in a chiral environment. Most importantly, they may react at very different rates with other chiral species. This is the reason why many compounds are biologically active, while their enantiomers are not. For example, the “R-form” of mecoprop (see Fig. 2.5) is an active herbicide, whereas the “S-form” is biologically rather inactive (Bosshardt, 1988).

Figure 2.5 The two enantiomeric forms of the herbicide mecoprop, The second type of stereoisomerism encompasses all other cases in which the three- The * indicates the asymmetric dimensional structures of two isomers exhibiting the same connectivity among the carbon center. Ar denotes the are- atoms are not superimposable. Such stereoisomers are referred to as diastereomers. matic substituent.

Diastereomers may arise due to different structural factors. One possibility is the presence of more than one chiral moiety. For example, many natural products contain 2 to 10 asymmetric centers per molecule, and molecules of compound classes such as polysaccharides and proteins contain hundreds. Thus, organisms may build large molecules that exhibit highly stereoselective sites, which are important for many biochemical reactions including the transformation of organic pollutants.

I

CI

rnecoprop

+cH3

HOO

\2H3 ArO i H3C OAr

mirror

(R)-form (S)-form


Another important form of diastereoisomerism results from restricted rotation around bonds such as are encountered with double bonds and/or ring structures. When considering the geometry of a double bond, we imagine a combination of two different types of bonds between two atoms. One of the bonds would be equivalent to a single bond, that is, a bond in which the pair of electrons occupies the region around the axis between the doubly bonded atoms. We can picture the second bond, which is called a n-bond (e.g., carbon-carbon, carbon-oxygen, carbon-nitrogen, carbon-sulfur, nitrogen-oxygen), by imagining the two bonding n-electrons to be present in an "electron cloud" located above and below a plane in which the axes of all other bonds (real or imaginary in the case of nonbonded electron pairs) lay, as in the case of ethylene (Fig. 2.6). The atoms closest to a carbon-carbon double bond are in a plane with bond angles of about 120" (see examples given in Fig. 2.3). Rotation about the axis would mean that we would have to break this bond. In triple- bond compounds, as in the case for acetylene (ethyne), there are two n-bond electron clouds which are orthogonal to each other, thus leading to a linear (bond angles = 180") configuration (Fig. 2.6).

acetylene (ethyne) Let us now consider a compound XHC=CHY in which X,Y f H. In this case, there are two isomers (sometimes also called geometric isomers), which are distinct (and,

Figure 2.6 Simplified picture of a in principle, separable) because we can no longer rotate about the C-C bond (Fig. and We (aceQ- 2.7). To distinguish between the two isomers, one commonly uses the terms cis and

lene) bond, respectively. trans to describe the relative position of two substituents (atoms or groups other than hydrogen). The term cis is used if the two substituents are on the same side of the double bond, the term trans if they are across from one another. As in the other cases of isomerism, we note that such closely related compounds may exhibit quite different properties. For example, the boiling points of cis- and trans-l,2-dichloroethene (X=Y=Cl), are 60 and 48"C, respectively. More pronounced differences in properties between cisltrans isomers are observed when interactions between two substituents (e.g., intramolecular hydrogen bonding) occur in the cis but not in the trans form, as is encountered with maleic and fumaric acid (see margin). These two compounds are so different that they have been given different names. For example, their melting points differ by more than 150°C and their aqueous solubilities by more than a factor of 100.

ethylene (ethene)

o-bond

H CIS

H Y w

trans

H\ r P=C\

C\\ 0.- tt 0

maleic acid

The organization of atoms into a ring containing less than 10 carbons also prevents free rotation. Consequently, cis and trans isomers are also possible in such ring systems; the cis isomer is the one with two substituents on the same side of the ring (i.e., above or below); the trans isomer exhibits a substituent on either side (Fig. 2.8).

In ring systems with more than two substituted carbons, more isomers are possible. An example well known to environmental scientists and engineers is 1,2,3,4,5,6- hexachlorocyclohexane (HCH). Three of the possible 8 isomers (the so-called a-, p-, y-isomers) are particularly important from an environmental point of view:

,c =o

(cis)

0 'C- OH

H\ C=< HO-d \\ H

0

furnaric acid (trans)

Figure 2.7 Cis/truns isomerism at double bonds exhibiting two substituents. a -isomer p-isomer y-isomer


X

cis +: ....... - -...

X

trans

Figure 2.8 Cisltrans isomerism in ring systems, such as in cyclohexane.

m chair

twist

boat

Figure 2.9 Different possible conformations of a six-membered ring (e.g., cyclohexane).

At this point we should reiterate that the relative positions of atoms in many structures are continuously changing. The term different conformations of a molecule is used if two different three-dimensional arrangements of the atoms in a molecule are rapidly interconvertible, as is the case if free rotations about sigma bonds are possible. If rotation is not possible, we speak of different con$gurutions, which represent isomers that can be separated. Obviously, the conformation(s) with the lowest energy [the most stable form(s)] is(are) the one(s) in which a molecule will preferentially exist. In the case of six-membered rings such as cyclohexane, three stable conformations (i.e., the chair, twist, and boat form) exist (Fig. 2.9).

Depending on the type of substituents, usually one of the forms is the most stable one. In the case of the HCH isomers, this is the chair form. If we have a closer look at the chair form, we see that six of the bonds linking substituents to the ring are directed differently from the other six:

a a

The six axial bonds are directed upward or downward from the “plane” of the ring, while the other six equatorial bonds are more within the “plane.” Conversion of one chair form into another converts all axial bonds into equatorial bonds and vice versa. In monosubstituted cyclohexanes, for electronic reasons, the more stable form is usually the one with the substituent in the equatorial position. If there is more than one substituent, the situation is more complicated since we have to consider more combinations of substituents which may interact. Often the more stable form is the one with more substituents in the equatorial positions. For example, in a-1,2,3,4,5,6- hexachlorocyclohexane (see above) four chlorines are equatorial (aaeeee), and in the Pisomer all substituents are equatorial. The structural arrangement of the /I-isomer also greatly inhibits degradation reactions [the steric arrangement of the chlorine atoms is unfavorable for dehydrochlorination (see Chapter 13) or reductive dechlo- rination; see Bachmann et al. 19881.

Delocalized Electrons, Resonance, and Aromaticity

Having gained some insights into the’ spatial orientation of bonds in chemicals and the consequences for the steric arrangement of the atoms in an organic molecule, we can now proceed to discuss special situations in which electrons move throughout a region covering more than two atoms. The resulting bonds are often referred to as “delocalized chemical bonds.” From an energetic point of view, this diminished constraint on the positions of these electrons in the bonds results in their having lower energy and, as a consequence, the molecule exhibits greater stability. For us the most important case of delocalization is encountered in molecules exhibiting multiple n-bonds spaced so they can interact with one another. We refer to such a series of n-bonds as conjugated. To effectively interact, n-bonds must be adjacent to each other and the o-bonds of all atoms involved must lie in oneplane. In such a conjugated system, we can qualitatively visualize the z-electrons to be smeared over


the whole region, as is illustrated for propenal (also known as acrolein) in Fig. 2.10. If we try to consider acrolein’s structure by indicating the extreme possible positions of these conjugated electrons, we may write

acrolein The back-and-fourth arrows are not intended to suggest the three structures are interconvertible, but rather that the location of the four electrons is best thought of

Figure 2*10 Schematic picture Of as something in between situation of such extreme possibilities. This freedom in lein (CH,=CH-CHO). electron positions results in what we call their delocalization, and the visualization

of a given molecule by a set of localized structures is called the resonance method for representing a structure. The relative contributions of the extremes to the overall resonance structure is determined by their relative stabilities. The stabilizing effect of delocalization is most pronounced in so-called aromatic systems. The best-known aromatic system is that of benzene, where we have three conjugated double bonds in a six-membered ring:

x-electron delocalization in acro-

Note again that each of the “static structures” alone does not represent the molecule, but that the molecule is a hybrid of these structures. Thus, as shown in parentheses, the electrons in the conjugated n-bonds of benzene are sometimes denoted with a circle.

In substituted benzenes (i.e., benzenes in which hydrogen is substituted by another atom or group of atoms), depending on the type and position of the substituents, the different resonance forms may exhibit somewhat different stabilities. Therefore, their contributions are different.

A quantitative estimate of the stabilization or resonance energy of benzene (which cannot be directly measured) may be obtained by determining the heat evolved when hydrogen is added to benzene and to cyclohexene to yield cyclohexane:

benzene cyclohexane

cyclohexene cyclohexane

If each of the double bonds in benzene were identical to the one in cyclohexane, the heat of hydrogenation of benzene would be three times the heat evolved during hydrogenation of cyclohexene. The values given above show that there is a large discrepancy between the “expected” (-120.7 x3 = -362 kJ . mol-’) and the measured


Q ..

furan (6)

.. pyridine (6)

\ H

indole (10)

a / /

naphthalene (10)

phenanthrene (14)

(-208.6 W-mol-') @ value. Hence, benzene is about 150 M-mol-' more stable than would be expected if there were no resonance interactions among the n-electrons. Large stabilization energies are observed not only in components containing a benzene ring but, in general, in cyclic n-bond systems with 4n + 2 (i.e., 6 , 10, 14 ...) electrons. In the early days of organic chemistry, it was recognized that the benzene ring is particularly unreactive compared to acyclic (noncyclic) compounds containing conjugated double bonds. The quality that renders such ring systems especially stable was and still is referred to as aromaticity. Some additional examples of aromatic ring systems are given in Fig. 2.11. We note that there are aromatic compounds containing heteroatoms that contribute either one (e.g., pyridine) or two electrons (e.g., furan, indole) to the conjugated n-electron systems. Also note that some polycyclic compounds are referred to as polycyclic aromatic compounds, although they are not aromatic throughout their structure in a strict sense (e.g., pyrene with 16 electrons in its n-bond system, see below; Fig. 2.13).

As already indicated for acrolein and for the five-membered heteroaromatic rings (e.g. furan; Fig. 2.11), resonance may also be important between nonbonded electrons on a single atom and a n-bond system. For example, an unshared electron pair of oxygen greatly contributes to the stabilization of the carboxylate anion:

: 0:- " 0 II & - . + I

R+O - ,Ao:

Similarly, the two unshared electrons of the nitrogen in aniline are in resonance with the aromatic n-electron system:

Figure 2.11 Some additional examples of organic compounds that are aromatic (in parentheses, number of n-electrons).

aniline

As we will see in Chapter 8, the delocalization of the unshared electron pair in aniline has an important impact on the acidhase properties of anilines as compared to aliphatic amino compounds.

In summary, delocalization of electrons enhances stability, and we can visualize delocalized bonding by using the resonance method. In later chapters we will learn more about the effects of resonance on chemical equilibrium and on the kinetics of chemical reactions of organic compounds.

Classification, Nomenclature, and Examples 31

lindane

Classification, Nomenclature, and Examples of Environmental Organic Chemicals

When grouping or classifying environmental organic chemicals, instead of taking a strictly structural approach, it is common practice to use compound categories based on some physical chemical properties, andor based on the source or the use of the chemicals. Terms such as “volatile organic compounds” (VOCs), “hydrophobic compounds,” “surfactants,” “solvents,” plasticizers,” “pesticides” (herbicides, insecticides, fungicides, etc.), organic dyes and pigments, or mineral-oil products are very common in the literature. Using such compound categories undoubtedly has its practical value; however, one has to be aware that each of these groups of chemicals usually encompasses compounds of very different structures. Thus the environmental behavior of chemicals within any such group can vary widely.

If we look at organic compounds from a structural point of view, we have already seen that organic molecules are composed of a skeleton of carbon atoms. This skeleton is sheathed in hydrogens, with heteroatoms (ie., halogens, 0, N, S, P) or groups of heteroatoms inserted in or attached to that skeleton. Such “sites” are called functional groups, since they are commonly the site of reactivity or function. Hence, the functional groups deserve our special attention if we want to understand or assess a compound’s behavior in the environment. For classifying organic chemicals according to structural features, we may then have to base the classification on the type of the carbon skeleton, the type of functional group(s) present, or a combination of both. Consequently, particularly when dealing with a compound exhibiting several functionalities, its assignment to a given compound class may be somewhat arbitrary.

There exists a systematic nomenclature for naming individual organic compounds. For a more detailed treatment of this topic we refer, however, to any organic chemistry textbook. It should be noted, however, that especially in environmental organic chemistry, one frequently uses so-called “common77 or “trivial” names instead of a compound’s systematic name. Furthermore, quite frequently, the systematic nomenclature is applied incorrectly. For a given compound, one may therefore find a whole series of different names in the literature (see, for example, synonym listings in the Merck Index), which, unfortunately, may sometimes lead to a certain confusion. For example, one of the 1,2,3,4,5,6-hexachlorocyclohexane isomers, lindane (see margin), also goes by the following diverse aliases: y-HCH, y-benzene hexachloride, gamma hexachlor, ENT 7796, Aparasin, Aphtiria, y-BHC, Gammalin, Gamene, Gamiso, Gammexane, Gexane, Jacutin, Kwell, Lindafor, Lindatox, Lorexane, Quellada, Streunex, Tri-6, and Viton. In this book, however, we shall consider problems encountered with nomenclature to be only of secondary importance as long as we are always aware of the exact structure of the compound with which we are dealing.

In the following discussion we will try to pursue three goals at the same time. Our first goal is to provide an overview of the most important functional groups present in environmentally relevant organic chemicals. The second goal is to take a first qualitative look at how such functional groups may influence the chemistry of given


CH3-

methyl

CH3- CHz-

ethyl

CH3- CHT CHz-

n -propyl

R- CH, - primary

\ R,

R2 /“”-

secondary

tertiary

6& 5 / 3

4

ortho or 1,2-

5 6& 4 B

meta or 1,3.

14 B

para or 1,4-

organic compounds. Finally, by introducing the nature and origins of specific compounds and compound classes, we will try to familiarize ourselves with the most important “chemical actors” that will be with us throughout the book.

The Carbon Skeleton of Organic Compounds: Saturated, Unsaturated, and Aromatic Hydrocarbons

Let us start out by a few comments about the terms used to describe carbon skeletons encountered in organic molecules. When considering a hydrocarbon (i.e., a compound consisting of only C and H) or a hydrocarbon group (i.e., a hydrocarbon substituent) in a molecule, the only possible “functionalities” are carbon-carbon double and triple bonds. A carbon skeleton is said to be saturated if it has no double or triple bond, and unsaturated if there is at least one such bond present. Hence, in a hydrocarbon, the term “saturated” indicates that the carbon skeleton contains the maxi- mum number of hydrogen atoms compatible with the requirement that carbon always forms four bonds and hydrogen one. A saturated carbon atom is one that is singly bound to four other separate atoms.

Carbon skeletons exhibiting no ring structures (ie., only unbranched or branched chains of carbon atoms) are named aliphatic, those containing one or several rings are alicyclic, or in the presence of an aromatic ring system, aromatic. Of course, a compound may exhibit an aliphatic as well as an alicyclic and/or an aromatic entity. The hierarchy of assignment of the compound to one of these three subclasses is commonly aromatic over alicyclic over aliphatic. A saturated aliphatic hydrocarbon is called an alkane or a paraffin. If considered as a substituent, it is referred to as an alkyl group. Alkyl groups represent the most ubiquitous hydrocarbon substituents present in environmental organic chemicals. Their general formula is -CnH2n+l. Examples of common alkyl groups are methyl, ethyl, propyl, and butyl groups (see margin). The prefix n, which stands for normal, is used to denote an unbranched alkyl chain, is0 means that there are two methyl groups at the end of an otherwise straight chain, and neo is used to denote three methyl groups at the end of the chain. Alkyl groups are further classified according to whether they are primary, secondary, or tertiary. Hence, an alkyl group is referred to as primary if the carbon at the point of attachment is bonded to only one other carbon, as secondary (s-) if bonded to two other carbons, and as tertiary (t-) ifbonded to three other carbons (see margin). Note that R stands for a carbon-centered substituent.

With respect to unsaturated hydrocarbons we should note that compounds exhibiting one or several double bonds are often called alkenes or olejins. Finally, we need to add a brief note concerning the nomenclature in aromatic systems, particularly, in six-numbered rings such as benzene. Here the terms ortho-, meta-, andpara-substitution are often used to express the relative position of two substituents in a given ring system. Identically, we could refer to those isomers as 1,2-(ortho), 1,3-(meta), 1,4-(para) disubstituted compounds (see margin).

Hydrocarbons are ubiquitously in the environment. Natural hydrocarbons range widely in size from methane to p-carotene (Fig. 2.12); many other branched, ole-finic, cyclic, and aromatic hydrocarbons are found in fossil hels, or tend to derive


Figure 2.12 Examples of aliphatic, alicyclic, and olefinic hydrocarbons.

CH4 /\/\/ methane n-hexane n-hexadecane

p-carotene (orange pigment in carrots; converted to vitamin A by human liver)

Y% 0 iso-octane 2,6,10,14-tetramethyl-pentadecane (pristane) cyclohexane

H

(trans-decalin) bicyclo[2.2.1 Iheptane trans-bicyclo[4.4.0]decane - v

1 -hexene 1,3-butadiene

P cis-bicyclo[4.4.0]decane

(cis-decalin)

from synthetic processing of fossil fuels (Figs. 2.12 and 2.13). When we recognize that the global annual production of liquid petroleum products (e.g., gasoline, kerosene, heating oils) is about 3 billion metric tons, it should be no surprise that producing, transporting, processing, storing, using, and disposing of these hydrocarbons poses major problems. Such organic chemicals are released to the atmosphere when we pump gasoline into the tanks of our cars, or are introduced to the surfaces of our streets when these same cars leak crankcase oil. Thus it is not only the spectacular instances such as the wreck of the Exxon Vuldez or blowout of the IXTOC-I1 off- shore oil well through which petroleum hydrocarbons pollute the environment. Fur- thermore, although they share a common source, the various hydrocarbons in the exceedingly complex mixture that is oil certainly do not behave the same in the environment. Some constituents are noted for their tendency to vaporize while others clearly prefer to bind to solids; some oil hydrocarbons are extremely unreactive (“methane is the billiard ball of organic chemistry”) while others interact beautifully with light; some are quite nontoxic while others are renowned for their carcinogenic- ity. Thus petroleum hydrocarbons have motivated a great deal of the research in environmental organic chemistry, and the individual components are a good example of why chemical structures must be visualized to predict fate. We should note at this point that hydrocarbons are not capable of forming hydrogen bonds to polar species such as water. As a consequence, many of these compounds are hydrophobic (“water-hating”).


Figure 2.13 Examples of aromatic hydrocarbons.

Benzene Derivatives

benzene toluene ethyl-benzene o, m, pxylene styrene

biphenyl

naphthalene

pyrene

indene 1,2,3,4-tetrahydro-naphthalene (tetralin)

P olycyck Aromatic Hydrocarbons (PAHs)

anthracene phenanthrene

benzo[a]pyrene perylene

Among the very large number of hydrocarbons present in the environment, we want to specifically mention two groups of compounds that are of particular interest to many environmental chemists and engineers: the so-called BTEX compounds (BTEX stands for benzene, toluene, ethylbenzene, and the three xylene isomers; Fig. 2.13), and thepolycyclic aromatic hydrocarbons (PAHs; Fig. 2.13). The BTEX components are important gasoline constitutents, and they are also widely used as solvents. They are very common soil and groundwater pollutants, and, particularly because of the high toxicity of benzene, they have triggered much research in the field of soil and groundwater remediation. The major sources of PAHs in the environment include the combustion of fossil fuels (gasoline, oil, coal), forest fires, (accidental) direct inputs of mineral oils, and the use of creosotes as wood preservatives. We should note that we produce PAHs when barbecuing steaks or other meat. From a human health perspective, PAHs have drawn considerable interest primarily because some of them (e.g., benzo(a)pyrene, Fig. 2.13) are very potent carcinogens. This is the main reason why PAHs are considered to be among the most important air pollutants. Furthermore, because of their high tendency to bioaccumulate (Chapter lo),


PAHs are also of great ecotoxicological concern. It is, therefore, not surprising that these compounds are among the most intensively investigated organic pollutants. Hence, they will serve us as important model compounds throughout this book.

Organohalogens

Many of the organic chemicals that are of particular environmental concern contain one or several halogen atoms, especially chlorine (Cl), and, to a lesser extent, fluorine (F) and/or bromine (Br). Although the number of known naturally produced organohalogens (e.g., produced by marine algae) has increased drastically over the last two decades, the problems connected with the ubiquitous occurrence of halogenated compounds in the environment are primarily due to anthropogenic inputs. There are several reasons for the still-vast industrial production of halogen-containing chemicals. First, because of their high electronegativity (Table 2.3), particularly fluorine and chlorine form rather strong bonds with carbon (Table 2.2). Hence, in many cases, substitution of carbon-bound hydrogens by halogens enhances the inertness of a molecule (and thus also its persistence in the environment). For example, many agrochemicals capable of enzyme inhibition are fluorine-stabilized analogues of the natural enzyme substrate (e.g., they exhibit a CF,- instead of a CH,-group; see Key et al., 1997). Furthermore, we should note that halogens (including fluorine) present in organic compounds, particularly when bound to aromatic carbon atoms, have a very weak tendency to be engaged in hydrogen bonds with H-donors such as water. Therefore, as we will learn in later chapters, the presence of the larger halogens (i.e., chlorine and bromine) renders a compound more hydrophobic, which increases its tendency to partition into organic phases including organisms. This is one reason why numerous pesticides contain one or several halogens.

From a quantitative point of view, the major halogenated compounds used and released into the environment are the polyhalogenated hydrocarbons, that is, compounds that are made up only of carbon, hydrogen, and halogens. By choosing the appropriate type and number of halogens, inert, nonflammable gases, liquids, and solids exhibiting specific physical-chemical properties that make them suitable for specific purposes can be designed. Prominent examples include the polyhalogenated C,- and C,-compounds that are used in enormous quantities (i.e., in tens to hundreds of thousands of tons per year) as aerosol propellants, refrigerants, blowing agents for plastic forms, anaesthetics, and, last but not least, as solvents for various purposes (see examples given in Table 2.4). Due to their high volatility and persistence in the troposphere, particularly the fluoro- and chlorofluorocarbons (FCs and CFCs, also calledfreons) are of great concern because of their significant stratospheric ozone- depletion and global-warming potential. This is the reason why these compounds are being replaced by hydrofluoro- and hydrochlorofluorocarbons (HFCs and HCFCs, respectively), which exhibit significantly shorter atmospheric lifetimes (Wallington et al., 1994). However, they are still not the ultimate solution to the problem.

The chlorinated solvents, including dichloromethane, tri- and tetrachloroethene, and 1,l , 1-trichloroethane (see Table 2.4), are still among the top groundwater pollutants. Under oxic conditions these compounds are quite persistent and, because they are also quite mobile in the subsurface, they can lead to the contamination of large groundwater areas. As we will discuss in detail in Chapters 14 and 17, in anoxic


Table 2.4 Examples of Important Industrially Produced C, - and C,-Halocarbons

Compound Name( s) Formula Major Use

Dichlorodi fluoromethane (CFC- 12)

Trichlorofluoromethane (CFC- 11)

Chlorodifluoromethane (HCFC -22)

1, 1,1,2-Tetrafluoroethane (HCFC- 134a)

1,1 -Dichloro-1 -fluoroethane

Dichloromethane (Methy lene chloride)

Trichloroethene (Trichloroethylene, TRI)

Tetrachloroethene (Tetrachloroethylene, Perchloroethylene, PER)

1,l ,l -Trichloroethane

(HCFC- 14 1 b)

CC12F2

CC13F

CHClFz

CF3-CHF

CClZF-CH3

CH2C12

CHCkCC12

cc12=cc12

CCI3-CH3

Aerosol propellants, refrigerants (domestic, automobile air conditioning), flowing agents for plastic foams. etc.

solvent

solvent

solvent

solvent

environments (i.e., in the absence of oxygen), such polyhalogenated hydrocarbons may be reduced to less halogenated compounds by a process called reductive deha- logenation. We should note that substitution of a carbon-bound hydrogen (oxidation state +I) by a halogen (-I) increases the oxidation state of the corresponding carbon atom by +II. Hence, for example, in tetrachloromethane (CCI,) the carbon atom is fully oxidized (i.e., its oxidation state is +IV). Therefore, it should not be surprising that some of the polyhalogenated hydroarbons are very effective electron acceptors in abiotic and biological redox reactions.

The C,- and C,-halocarbons are, of course, not the only group of persistent polyhalogenated hydrocarbons that are of great environmental concern. Because of their hydrophobic character and tendency to bioaccumulate (see Chapter lo), a variety of polychlorinated aromatic hydrocarbons have gained a very bad reputation. Recently, some of these classical environmental contaminants have again been subject of considerable debates, because some of the compounds are suspected to exhibit endocrine-disrupting properties (i.e., they may disturb the hormonal balance ir? organisms; see Keith, 1997). Examples of such compounds include the polychlorinated biphenyls (PCBs) and the polychlorinated terphenyls (PCTs) (see Fig. 2.14). These chemicals are commonly used as complex technical mixtures of different congeners; that is isomers and compounds exhibiting different numbers of chlorine atoms but having the same source (i.e., chlorination of a biphenyl skeleton). Note that there are 209 possible PCB congeners (Ballschmiter et al., 1992), and as many


polychlorinated biphenyls polychlorinated terphenyls (PCBs, 209 possible congeners) (PCTs, 8149 possible congeners)

CI H CI p,p'-DDT hexachlorobenzene 1,2,3,4,5,6-hexachlorocyclohexane

(HCB) (HCH, 8 isomers, one of them exists as a pair

of enantiomers)

Figure 2.14 Some examples of classical polychlorinated hydrocarbons that are globally distributed.

as 8149 possible PCT congeners (Wester et al., 1996). To date, more than 1 million metric tons of PCBs and PCTs have been produced and have been used in waxes, printing inks, paints and lacquers and as capacitor dielectric fluids, transformer coolants, hydraulic fluids, heat-transfer fluids, lubricants, plasticizers, and fire retardants. They are lost to the environment during production and storage, and particularly from disposal sites. Although PCBs and PCTs have been banned or at least severely restricted in numerous countries, they are still ubiquitous in the environment. Together with other polychlorinated hydrocarbons (e.g., the classical organochlorine pesticides, p,p'-DDT, HCB, and HCH; see Fig. 2.14), they can be detected everywhere in the world, at the bottom of the ocean as well as in arctic snow, indicating the powerful environmental transport mechanisms acting on such chemicals (Vallack et al., 1998; Kallenborn et al., 1998). A group of compounds that is structurally related to the PCBs, and that cause similar environmental problems, are thepolybrominated biphenyls (PBBs) that are used in large quantities as flame retardants (de Boer et al., 2000).

Finally, we should note that large amounts of known and unknown chlorinated compounds are formed and released to the environment due to the use of chlorine in waste and drinking water disinfection, and in bleaching processes in the pulp and paper industry. Important in this group are, for example, the trihalomethanes (THMs; CHCl,, CHBrCl,, CHBr2C1, and CHBr,).

Oxygen-Containing Functional Groups

Among the heteroatoms present in natural and anthropogenic organic compounds, oxygen plays a unique role because it is part of a large number of important functional groups. Due to its high electronegativity (Table 2.3), oxygen forms polar bonds with hydrogen, carbon, nitrogen, phosphorus, and sulfur. As a consequence, these functionalities have a significant impact on the physical-chemical properties as well as on the reactivity of a given compound. In the following we will discuss some functional groups involving solely oxygen. Additional oxygen-containing functionalities will be addressed below when discussing functional groups exhibiting also other heteroatoms (i.e., N, S, P). Finally, we should note again that, in general,


w

functionalities. Some o f the common uses and/or sources o f the

Figure 2.15 Examples o f important industrially and/or environmentally relevant chemicals exhibiting either alcohol or ether

1 -chloro-2,3-epoxy-propane polychlorinated dibenzo- (epichlorohydrine, pdioxines (PCDDs,

chemical intermediate) 175 possible congeners)

alcohols (R-OH) and phenols (Ar-OH)

H3C- OH OH Ho-OH G O H % / D O H

methanol, ethanol, ethylene glycol CInl

phenol chlorophenols, (chemical inter- (chemical inter- (chemical intermediate, (chemical intermediate) m = 1 - 5.

mediate, solvent) mediate, solvent) solvent, coolant, chemical inter. antifreeze agent) mediates, biocides)

2,2-bis-(4-hydroxy-phenyl)-propane 4-nonylphenol ("bisphenol A," chemical intermediate) (metabolite of nonionic surfactants)

ethers (R,-O-RJ

diethyl ether 1,4-dioxane methyl-t-butyl-ether (solvent) (solvent) (MTBE, gasoline

additive)

xc1 0

2,6di-t-butyl-pcresol (DBPC, antioxidant)

methyl-phenyl-ether (anisole, methoxy-benzene)

polychlorinated dibenzo- furans (PCDFs,

135 possible congeners)

oxygen forms two bonds and that its oxidation state in organic molecules is commonly -11, except for peroxides (e.g., R-0-0-R), where its oxidation state is -I.

Alcohol and Ether Functions. Let us start by looking at the simplest oxygen functions, that is, an oxygen atom that forms single bonds to a carbon and a hydrogen atom, or to two carbon atoms. The former group is called a hydroxyl group or alcohol function (R-OH), while in the latter case we speak of an ether function (R1-O-R2). Some illustrative examples of industrially and/or environmentally relevant chemicals exhibiting such functional groups are given in Fig. 2.15.

As we will see in Part 11, both alcohol and ether functions have a great influence on the physical-chemical properties and the partitioning behavior of an organic chemical, because of the ability of the oxygen atom to participate in hydrogen bonds. We should, however, note that this ability is significantly greater in the case of the hydroxyl group, because R-OH may act as both H-donor and H-acceptor, while Rl-O-R2 is only an H-acceptor. Furthermore, depending on the nature of R, the


Figure 2.16 One example of a group of wide ly used chemicals that, as a consequence of biological waste-water treatment, are converted into persistent objectionable degradation intermediates.

CI c*cl CI CI

pentachlorophenol

/I Hzo

pentachloro-phenolate ion

4-nonylphenol-polyethyleneglycol-ethers (surfactants)

microbial degradation in sewage treatment plants 1

persistent, ecologically objectionable intermediates

R-OH group may dissociate in aqueous solution; that is, it may act as a weak acid. This is primarily the case for aromatic alcohols, including phenols, that are substituted with electron-withdrawing substituents. A prominent example is the biocide pentachlorophenol (see margin), which in aqueous solution at pH 7.0, is present primarily (> 99%) as pentachloro-phenolate anion. As will be discussed in detail in Chapter 8 and in Part 111, such charged species have very different properties and reactivities as compared to their neutral counterparts. Finally, we should note that phenols, particularly those that are substituted with electron-donating substituents (e.g., alkyl groups), may be oxidized in the environment, leading to a variety of products (Chapters 14 and 16). Some phenols that are especially easily oxidized and that form stable radicals are even used as antioxidants (e.g., DBPC, Fig. 2.15) in petroleum products, rubber, plastics, food packages, animal feeds, and so on (Kirk-Othmer, 1992).

Before we turn to other bctional groups a few remarks about some of the compounds listed in Fig. 2.15 are necessary. First, we should note that with an annual global production of between 15 to 20 million metric tons (!), methanol and methyl-t-butyl ether (MTBE) are among the top high-volume industrial chemicals. While methanol does not have to be considered as a major environmental problem, MBTE, which is used as oxygenate in gasoline to improve the combustion process, has become a wide- spread pollutant in the aquatic environment, particularly in groundwater (Squillace et al., 1997). The reason is the significant water solubility of MBTE and its rather high persistence toward biodegradation. In Chapter 7 we will address the issue of the dissolution of chemicals from complex mixtures such as gasoline into water.

Phenols. Phenolic compounds are used in very large quantities for a variety of industrial purposes. They may also be formed in the environment by abiotic (e.g., hydroxylation in the atmosphere; see Chapter 16) or biological processes (Chapter 17). A prominent example for the latter case is the formation of 4-nonylphenol from the microbial


0 II

0 II

aldehyde (R- C-H) and kefo (RI - C-R,) funcfions

0 I /

H

formaldehyde (disinfectant,

chemical intermediate)

JH acrolein

(chemical intermediate for polymer production)

K,, acetic acid

acetaldehyde (chemical intermediate,

solvent)

acetone (chemical intermediate,

solvent)

isobutylaldehyde (chemical intermediate,

solvent, desinfection byproduct in drinking water)

2-butanone (solvent)

0 II

carboxyl (R---OH) and carboxylic acid ester (R, -C-OR,) functions

O Y "

benzaldehyde (chemical intermediate,

solvent)

methyl phenylketone (acetophenone; chemical

intermediate, solvent)

CI

trichloroacetic acid benzoic acid (food (R,S)-2-(4-chloro-2-methyl (herbicide, atmospheric preservative, phenyl) - propionic acid breakdown product of additive, chemical ((R,S)-mecoprop; herbicide) chlorinated solvents) intermediate)

ethylacetate (acetic acid ethyl ester; solvent)

Preventol 0 82

(roof protection agent)

phthalates

plasticizers) (Rl, R,=Cj,toCI,;

Figure 2.17 Examples o f impor- degradation of 4-nonylphenol polyethyleneglycol ethers (Fig. 2.16) that are used as tant and/or environ- nonionic surfactants (Giger et al., 1984; Ahel et al., 1996). We will make some mentally relevant chemicals exhib- general comments on surfactants later when discussing some other types of iting carbon-oxygen functions with a doubly bound oxygen. Some commercially important surface active chemicals. The case of the 4-nonylphenol ofthe common uses and/or S O ~ ~ S polyethyleneglycol ethers is an illustrative example demonstrating that under

certain circumstances, microbial transformation processes may lead to compounds o f the chemicals are given in parentheses.

that are of significantly greater concern than the parent compounds. Particularly 4-nonylphenol has been shown to exhibit some endocrine-disrupting activity. This is


Q 10 1

2,3,7,8-tetrachloro- dibenzo-p-dioxin

polybrominated diphenylethers (PBDEs)

the major reason why in many countries the use of 4-nonylphenol polyethyleneglycol ethers has been significantly restricted.

Polychlorinated Dibenzo-@)-Dioxins and Dibenzo-Furans. Another group of compounds that we need to specifically address are the polychlorinated dibenzo-p-dioxins (PCDDs) and dibenzo-furans (PCDFs) (Fig. 2.15). The PCDDs and PCDFs are not intentionally produced but are released into the environment from various combustion processes and as a result of their occurrence as unwanted byproducts in various chlorinated chemical formulations (e.g., chlorinated phenols, chlorinated phenoxy herbicides; see Alcock and Jones, 1996). Because some of the PCDD and PCDF congeners are very toxic (e.g., 2,3,7,8-tetrachloro dibenzo-p-dioxin, see margin), there have been and still are considerable efforts to assess their sources, distribution, and fate in the environment. Similarly to the PCBs or DDT (see above), the PCDDs and PCDFs are highly hydrophobic and very persistent in the environment. It is therefore not surprising that they have also been detected everywhere on earth (Brzuzy and Hites, 1996; Lohmann and Jones, 1998; Vallack et al., 1998). Finally, we should note that polybrominated diphenylethers (PBDEs, see margin) that, like the PBBs (see above), are used as flame retardants, are of increasing environmental concern (de Boer et al., 2000).

Aldehyde and Keto Functions. There are a number of other carbon-oxygen functions in which one oxygen forms a double bond with the carbon atom. If there is only one oxygen involved, depending on whether the carbon is bound to a hydrogen or not, one refers to the hnction as an aldehyde (R-CHO) or keto (R,-CO-R,) group, respectively (Fig. 2.17). As in the case of the ether oxygens, the aldehyde and keto oxygens are H-acceptors, which makes some simple aldehydes and ketones suitable solvents for many purposes ( e g , acetaldehyde, acetone, 2-butanone; see Fig. 2.17). Furthermore, because these functional groups are quite reactive, many important chemical intermediates used in industry exhibit an aldehyde or keto group. Finally, we should note that a variety of simple aldehydes such as isobutyraldehyde (Fig. 2.17) may be formed during water treatment. Since some of these compounds have an extremely unpleasant odor, their formation (e.g. as disinfection byproducts in drinking water) may cause considerable problems (Froese et al., 1999).

Carboxyl Groups. If we oxidize an aldehyde group, that is, if we replace the hydrogen by hydroxyl, we obtain an acidic function that is referred to as carboxylic acid (R-COOH) group (Fig. 2.17). As 'the name implies, carboxylic acids may dissociate in aqueous solution. Depending on R, their pl(,-values (see Chapter 8) are in the range between 0 and 6. Therefore, compounds exhibiting one or several carboxylic groups are present in natural waters primarily in dissociated (anionic) forms. Furthermore, we should point out that carboxylic acid functions are both strong H-donors and H-acceptors. Hence, the presence of a carboxylic function increases the water solubility of a compound significantly. The presence of such compounds in the environment may be due to direct input (e.g., as herbicides such as mecoprop, Fig. 2.17) or due to abiotic or biological transformation of other compounds (see Part 111). For example, the halogenated acetic acids, including mono-, di-, and trichloroacetic acid, as well as trifluoroacetic acid, have been detected in high concentrations (> 1 pg-L-') in rainwater. These acids are considered to be, at least in


part, atmospheric oxidation products of some of the chlorinated and fluorinated ethanes and ethenes discussed above (see, e.g., Berg et al., 2000). Carboxylic acids in the environment are also formed from hydrolysis reactions of carboxylic acid derivatives such as carboxylic acid esters and amides (see below). We will discuss such reactions in detail in Chapter 13.

Carboxylic Acid Ester Functions. As illustrated by Fig. 2.17, we generally talk of an ester hnction if we replace the OH of an acid by an OR group. Hence, because of the lack of an OH, an ester function (R,-COOR,) can only act as a H-acceptor. Therefore, it has a smaller impact on a compound’s water solubility as compared to the corresponding acid group. Ester functions can be found in many natural (e.g., oils and fats) and anthropogenic chemicals. One important group of man-made chemicals exhibiting carboxylic ester functions are the phthalates (Fig. 2.17), which are diesters of phthalic acid. The R, and R2 in Fig. 2.17 denote hydrocarbon groups, in most cases alkyl groups consisting of between 1 and 10 carbon atoms. The annual global production of phthalates, which are used primarily as plasticizers, exceeds 1 million metric tons. Phthalates are ubiquitous in the environment, and they are among the most notorious laboratory contaminants encountered when analyzing environmental samples. Carboxylic ester functions are also frequently encountered in compounds used as pesticides, or, more generally, as biocides. In some cases, the biologically active compound is actually the “free” carboxylic acid. The ester, which can be designed to have very different physical-chemical properties, is then used to provide the active compound continuously at the target site by (slow) hydrolysis of the ester function. An example is Preventol@ B2 (Fig. 2.17), which is applied to prevent root penetration in bituminous construction materials such as roofing felts, sealants, insulations, and asphalt mixtures. When hydrolyzed, Preventol@ B2 releases both (R,S)-mecoprop enantiomers (see Fig. 2.5 and Section 2.1, discussion on chirality). In a study on the occurrence and behavior of pesticides during artificial groundwater infiltration of roof runoff (Bucheli et al., 1998a), both mecoprop enantiomers were detected in roof runoff in much higher concentrations (up to 500 pg- L-’) than in the corresponding rainwater. It could be shown (Bucheli et al., 1998b) that Preventol@ B2 present in bituminous sealing sheets used on flat roofs was the source of these compounds. A comparison of the (R,S) mecoprop loads from flat roofs and from agricultural applications into surface waters in Switzerland revealed that these loads were in the same order of magnitude!

We conclude this brief discussion of carbon-oxygen functions by noting that, particularly among the pesticides, there are numerous compounds that are carbonic acid derivatives including carbonic acid esters (“carbonates,” R,OCOOR,). In this functional group, the carbon atom is fully oxidized, and therefore, hydrolysis yields CO, and the corresponding alcohols (see Chapter 13). We will encounter other carboxylic and carbonic acid derivatives in the following when introducing additional heteroatoms.

Nitrogen-Containing Functional Groups

Among the heteroatoms present in organic molecules, nitrogen and sulfur (which will be discussed later) are somewhat special cases in that, like carbon, they may assume many different oxidation states. Therefore, there are many nitrogen- and


Table 2.5 Some Important Nitrogen-Containing Functional Groups Present in Anthropogenic Organic Compounds

Name Group (oxidation state of nitrogen)

Name (oxidation state of nitrogen) Group

R2 I +

R4

RI-N-R~ ammonium (-111) R1-NH-NH-Rz hydrazo (-11) I

RI-N \ R3

amino' (-111) (amine)

0 carboxylic acid

R i ANNA. I amide' (-111) R3

, R2 ,N=N

R i azo (-1)

hydroxyl-amine (-I) /OH R-N,

H

0 R-CIN cyano, nitrilo (-111) R- N '/ nitroso (+I)

0

urea (-111)

carbamate (-111)

+,0-

+O R-N

+,0-

"0 R-0-N

nitro (+III)

nitrato (+V) (nitrate)

"Primary if R2 = R3 = H; secondary if R2 = H and R3 # H; tertiary if R2 # H and R3 # H.

sulfur-containing bctional groups exhibiting very different properties and reactivities. Table 2.5 gives an overview of the most important hc t iona l groups containing nitrogen, carbon and/or oxygen, but no other heteroatoms. As is evident from this table, nitrogen forms, in general, three, and in some special cases, four bonds, and it may assume oxidation states between -111 and +V. In the following we will confine our discussions to a few of the functional groups listed in Table 2.5. Some of the others will be addressed in later chapters (e.g., the carboxylic and carbonic acid derivatives including amides, nitriles, ureas, and carbamates will be discussed in the context of hydrolysis reactions in Chapter 13).

Amino Group. Nitrogen-containing groups often have a significant impact on the physical-chemical properties and reactivities of environmentally relevant chemicals in which they occur (Fig. 2.18). A very important group is the amino group that is present in numerous natural (e.g., amino acids, amino sugars) and anthropogenic chemicals. For example, aromatic amino compounds, particularly aminobenzenes (anilines; Fig. 2.18) are important intermediates for the synthesis of numerous different chemicals including dyes, pharmaceuticals, pesticides, antioxidants, and many more (Fishbein, 1984). Furthermore, benzene rings carrying amino groups are present in a variety of other, more complex molecules. In addition, as we will see in Chapters 13 and 14, hydrolysis of various acid derivatives including the ones mentioned above, as well as reductive transformations of aromatic azo- and nitro-


Figure 2.18 Examples of industrially and/or environmentally important chemicals exhibit ing pr imar i ly nitrogen-containing functional groups. Some of the common uses andor sources o f t h e chemicals are given in parentheses.

aniline and substituted anilines (chemical

intermediates)

f vNl

triethylamine (solvent, wetting agent,

corrosion inhibitor, propellant)

"0. NO2 O2

2,4,6-trinitrotoluene (TNT, explosive)

CI I

H H

atrazine (a triazine herbicide)

quaternary ammonium salts (cationic surfactants)

2,4-dinitro-ecresol (DNOC, herbicide)

N,N'-dialkyl/aryl p-phenylendiamines

(antioxidants)

azobenzene (chemical intermediate,

pesticide)

"nitroglycerin" (explosive)

Br

Dispersive Blue 79 1 -nitropyrene (textile dye) (airborne pollutant,

fuel combustion product)

compounds (see below), may lead to the formation of aromatic amino compounds in the environment.

From a chemical point of view, amino groups have several effects on the properties and on the reactivity of a given compound. Amino groups may engage in hydrogen bonding, both as H-acceptors and (to a somewhat lesser extent) as H-donors (only primary and secondary amines; see footnote in Table 2.5). However, in contrast to the weakly acidic OH, amino groups are weak bases. Hence, in aqueous solution they may acquire a proton, thus forming a cationic ammonium species (see Chapter 8). This is particularly important for aliphatic amino groups (e.g., triethylamine, Fig. 2.18) that are significantly stronger bases as compared to aromatic amino groups. Note that for use as surfactants, some amines are alkylated to quaternary ammonium compounds, thus forming stable cations (see example given in Fig. 2.18). We will address some of the special properties of surfactants below when


discussing another functional group that is widely used in this type of chemicals, the sulfonate group.

A second important feature of amino groups, particularly when bound to aromatic systems, is their capability to act as a n-electron donor (see Section 2.1). An interesting example is the azo dye, Disperse Blue 79 (Fig. 2.1 S), where the extreme shift in light absorption (blue color) is achieved by introducing a strong electron-donating substituent (i.e., a dialkylamino group) in para-position to the azo group at one of the benzene rings, and two strong electron-withdrawing substituents (i.e., two nitro groups, see below) in ortho- and para-position at the other benzene ring. We will discuss the effect of such substituents on the light absorption of organic compounds, and we will explain the reason why aromatic azo compounds absorb light quite strongly in the visible wavelength range (hence their use as dyes) in Chapter 15 when treating direct photolysis reactions.

Like some phenolic groups, amino groups bound to an aromatic system (e.g., benzene) may be oxidized in the environment. In some cases such transformations may lead to the formation of products that are of considerable concern (Chapter 14). Finally, some aromatic amines that form stable radicals are actually also used as antioxidants (see example given in Fig. 2.18 and Kirk-Other, 1992).

Nitro Group. The second nitrogen-containing functional group to discuss here is the nitro group. Numerous synthetic high-volume chemicals contain one or several nitro groups that are bound to an aromatic system, in particular to a benzene ring. Aromatic compounds exhibiting nitro groups are pivotal intermediates in various branches of the chemical industry, and they are also present in many end products including explosives ( e g , TNT, Fig. 2.18), agrochemicals (e.g., DNOC, Fig. 2.18), and dyes (e.g., Dispersive Blue 79, Fig. 2.18). In addition, a large number of nitroaromatic compounds are formed in the atmosphere by photochemically induced reactions of aromatic hydrocarbons involving hydroxyl radicals and nitrous oxides (Atkinson, 2000). Furthermore, such compounds have also been shown to be produced during fuel combustion, for example, in car engines (Tremp et al., 1993). Hence, such nitration processes may lead to a significant input of quite toxic compounds in the environment (e.g., nitropyrene, Fig. 2.18; see also Fiedler and Mucke, 1991).

In order to understand how a nitro group affects the properties and reactivity of a compound, we recall that the nitro group has a strong electron-withdrawing character and that it may delocalize z-electrons (see Chapter 8). Thus, particularly aromatic nitro groups have a big influence on the electron distribution in a molecule (or in parts of the molecule). This affects properties such as the acidity constant of an acid or base group bound to an aromatic system (Chapter S), the specific interaction of aromatic compounds with electron donors such as oxygen atoms present at clay mineral surfaces (Chapter 1 l), or the light absorption properties of the compound (Chapter 15). Furthermore, because in a nitro group the nitrogen exhibits an oxidation state of +I11 (Table 2.5), this group may act as an oxidant (i.e., as an electron acceptor). This is most evident in explosives (e.g., TNT, Fig. 2.18) where the nitro group hctions as built-in oxidant, thus ensuring a very fast oxidation of the molecule leading to the liberation of a large amount of energy in a very short time period. This effect can also


Table 2.6 Some Important Sulfur-Containing Functional groups Present in Anthropogenic Organic Compounds

Group Name

(oxidation state of sulfur) Group Name

(oxidation state of sulfur)

0 I I

II 0

0 I I

I I

R-SH thiol, mercaptan (-11) R-S-OH sulfonic acid (+IV)

R1- S- R2 thioether, sulfide (-11) RI-S-O-R~ sulfonic acid ester (+IV) 0

Ri- S-S-Rp

0 It

0 I I

RI-S-Rz II 0

thiocarbonyl (-11)

disulfide (-1)

sulfoxide (0)

sulfone (+II)

0 R2 II / sulfonic acid amide,

R1-S-N sulfonamide (+N) 1 ‘R,

0 It

R i -0 - S- 0- R2

I I 0

sulfuric acid ester, sulfate (+VI)

be achieved with nitrate groups (e.g., nitroglycerin, Fig. 2.18). Finally, in the environment, under reducing conditions, nitro groups may be reductively trans- formed into nitroso-, hydroxylamine-, and ultimately, amino groups (see Table 2.5), yielding products that may be of similar or even greater concern than the parent nitro compound (Chapter 14). This is particularly important when dealing with the assessment of ammunition waste sites, which, in many countries, cause severe problems with respect to soil and groundwater contaminations (Haderlein et al., 2000).

Sulfur-Containing Functional Groups

Sulfur in organic molecules may assume various different oxidation states (i.e., -11 to +VI, see examples given in Table 2.6). Because sulfur is a third-row element with d-orbitals, it is capable of valence-shell expansion to accommodate more electrons than are allowed by the simple octet rule. Hence, as can be seen from Table 2.6, sulfur may be surrounded by 10 ( e g , sulfoxide) or even 12 (e.g., sulfone, sulfonic acid and its derivatives) electrons, and this capacity allows for (redox) reactions that are inaccessible to the corresponding oxygen analogues. Other important differences from oxygen include its significantly smaller electronegativity (Table 2.3) and its much weaker tendency to be engaged in H-bonding. Furthermore, as compared to oxygen, sulfur forms weaker bonds with carbon and hydrogen (Table 2.3), causing mercapto groups to be more acidic (Chapter 8) and more nucleophilic


I? R1&,

I

I

- 0 I +

R, /%~

0 t

R2 R, /Sx

sulfoxides

R R1- 8-b

0

I -

0 I

R, -S++-R2 I 0-

sulfone

Figure 2.19 Examples o f “special” double bonds between sulfur and oxygen.

Figure 2.20 Examples o f industri- a l l y and/or environmentally important chemicals exhibit ing a sulfur-containing functional group. Some (not all) of the common uses and/or sources of the chemicals are given in parentheses.

(Chapters 13 and 14) as compared to their oxygen analogues. Finally, we should point out that sulfur may not only form regular n-double bonds (as in a thiocarbonyl, Table 2.6), but may also undergo a special type of double bond in which the sulfur’s d-orbitals participate (indicated by an arrow between S and 0 in Fig. 2.19). Such double bonds exist in functional groups in which sulfur is engaged in more than two bonds. In such cases we may represent the functional group by two resonance structures, but the bond is nevertheless localized. We should note that this type of double bond does not significantly change the geometry at the atoms involved (in contrast to the regular n-bond). Hence, the geometry at the sulfur atom is pyramidal. In fact, unsymmetrical sulfoxides (R, # R, in Fig. 2.19) are chiral (like an asymmetrically substituted carbon atom, see Section 2.2) in that the oxygen and the lone electron pair play the role of the other two substituents.

Figure 2.20 gives some examples of chemicals in which sulfur-containing functional groups govern the compound’s properties and/or reactivity. Among the compounds exhibiting reduced sulfur atoms (-1, -11), the low-molecular-weight mercaptans,

f i SH

ethylmercaptan (ethanethiol)

(odorant, chemical intermediate)

/ks/

/s\

CI

dimethylsulfide thiophenol chlorthiamide (solvent, important (chemical (herbicide) biogenic sources) intermediate)

dimethyldisulfide dimethylsulfoxide p-tOb?nesu~fOnate 2-amino-naphthalene-4,8-disulfonate (important biogenic (solvent) (hardener, stabilizer) (azo dye component)

source)

linear alkylbenzene sulfonates (US)

(surfactants)

4-4-bis (2-sulfostyryl)biphenyl (DSBP, fluorescent whitening agent)

\

COOH

0

R = C 11 -C 17

fatty alcohol sulfates (FAS) sulfadiazine sulfomethuron (surfactants) (a therapeutical drug) (a “sulfonylurea” herbicide)


/O fNTO CI

Alachlor

‘ I

SO,H

Sulfonic acid metabolite of

alachlor

Figure 2.21 Gluthathione S-trans- ferase enzyme-mediated reaction ultimately yielding a sulfonated metabolite (from Field and Thur- man, 1996).

dialkylsulfides, and dialkyldisulfides have anthropogenic and very important biogenic sources. For example, dimethyl sulfide is one of the major volatile sulfur compounds in the global sulfur cycle. Furthermore, thioether (R,-S-R,) or thioester (R,-CO-SR,) groups can be found in many pesticides. Finally, again in the world of pesticides, a C=S (e.g., chlorthiamide, Fig. 2.20) and, particularly, a P=S function (see below) is sometimes introduced instead of a C=O or P=O function, respectively. This change is made primarily with the goal of lowering the general toxicity of the compound. In some cases, the sulfur is then enzymatically replaced by an oxygen in the target organism(s), thus increasing the toxic effect of the compound at the target site.

With respect to the groups exhibiting oxidized sulfur atoms, aromatic sulfonic acid groups, in particular, benzene- and naphthalene sulfonic acids, are the most important ones (see in Fig. 2.20). Because sulfonic acid groups have very low pK, values (< 1, Chapter S), they dissociate completely in aqueous solution, thus forming the corres- sponding negatively charged sulfonates. Hence, such groups significantly increase the hydrophilicity (and thus the water solubility) of a compound. This is the major reason why aromatic sulfonate groups are present in a variety of commercially important chemicals including surfactants (e.g., linear alkylbenzene sulfonates, LAS, Fig. 2.20), anionic azo dyes, fluorescent whitening agents (Fig. 2.20), construction chemicals, and many more. In addition, introduction of a sulfonic acid group with the goal of forming a highly water-soluble metabolite is a detoxification strategy of a variety of organisms (see example given in Fig. 2.2 1). It is therefore not too surprising that numerous compounds exhibiting aromatic sulfonate groups have been detected in waste waters and leachates, as well as in natural waters (Altenbach and Giger, 1995; Field and Thurman, 1996; Stoll and Giger, 1998; Suter et al., 1999).

Among the compounds containing sulfonate groups the linear alkylbenzene sulfonates (LAS) have drawn particular interest, because they are produced and used in very large quantities (i.e., > 2 million metric tons per year; Ainsworth, 1996). They have been widely detected in aquatic (Matthis et al., 1999) as well as in terres- trial environments (where they are introduced primarily as a result of sewage sludge amendment; Jensen, 1999). We should note that LAS as well as other commercially available surfactants [e.g., alkylphenol polyethylene glycol ethers (Fig. 2.16), quar- ternary ammonium salts (Fig. 2.18), or fatty alcohol sulfates (Fig. 2.20)] are not single substances, but are mixtures of compounds of different carbon chain lengths. Owing to their amphiphilic character (partly hydrophilic and partly hydrophobic), the surfactants have special properties that render them unique among environmental chemicals. In aqueous solutions they distribute in such a manner that their concentration at the interfaces of water with gases or solids is higher than in the innex regions of the solution. This results in a change of system properties, for example, a lowering of the interfacial tension between water and an adjacent nonaqueous phase, and in a change of wetting properties. Furthermore, inside the solution, on exceeding certain concentrations surfactants form aggregates, called micelles. Hence, surfactants may keep otherwise-insoluble compounds in the aqueous phase, and they form an important part of any kind of detergent. Surfactants are also widely used as wetting agents, dispersing agents, and emulsifiers in all kinds of consumer products and industrial applications (Pion, 1987).


Figure 2.22 Examples of important anthropogenic chemicals exhibiting a phosphorus-containing functional group. Some of the common uses of the chemicals are given in parentheses.

i 0 t

R ,@P-OR, I OR2

phosphoric acid triester

Figure 2.23 Examples of a “special” double bond between phosphorus and oxygen indicated by the arrow between 0 and P.

CI

trichlorfon (insecticide)

triphenylphosphate (plasticizer, fire retardant)

phosphonates

B PH:: H O - P - F - r O H

OH CH,OH

0 II

I

1 -hydroxyethyC sarin 1 ,I -diphosphonic acid

(HEDP, complexing agent) (nerve poison)

phosphates and thiophosphates

parathion (paraoxon) disulfoton (insecticide, acaricide) (insecticide, acaricide)

We conclude this section by noting that, as in the case of other acids, there are various sulfonic and sulhric acid derivatives including esters and amides. Fig. 2.20 shows two examples of sulfonic acid amides that are representatives of important groups of drugs (sulfodiazine) and herbicides (sulfometuron), respectively.

Phosphorus-Containing Functional Groups

Although in principle, phosphorus may, like nitrogen, assume a variety of oxidation states (i.e., -111 to +V), in environmentally relevant compounds, it is present primarily in the more oxidized forms [i.e., +I11 (e.g., phosphonic acid derivatives) and, +V (e.g., phosphoric and thiophosphoric acid derivatives); see Fig. 2.221. In these oxidation states, P forms in most cases three single bonds and one double bond, the latter commonly with either an oxygen or a sulfur atom. As in the case of the other third-row element sulfur (Fig. 2.19), these double bonds are “special” double bonds (Fig. 2.23) that do not change the geometry at the atoms involved. From the examples given in Fig. 2.22 it can be seen that particularly esters and thioesters of phosphonic-, phosphoric-, and thiophosphoric acid are used for a variety of purposes including as plasticizers and/or flame retardants (Carlsson et al., 1997), and as pesticides, mostly as insecticides and acaricides (Tomlin, 1994; Hornsby et al., 1996). We will address these types of compounds in more detail in Chapter 13 when discussing hydrolysis reactions. Finally, we note that phosphonates (exhibiting one P-C bond) that contain several phosphonic acid groups are increasingly used as chelating agents (Nowack, 1998). Some phosphonic acid derivatives are extremely toxic nerve gases and, therefore, used as chemical weapons (e.g., sarin, Fig. 2.22).

Some Additional Examples of Compounds Exhibiting More Complex Structures

To round out our tour through the jungle of environmental organic chemicals, we shall have a brief look at some additional examples of compounds that have more


Figure 2.24 Some additional examples of pesticides exhibit ing a somewhat more complex structure.

oxadiazon (herbicide)

aldicarb (insecticide, acaricide, nematicide)

F3c9J30 CI

haloxyfop (herbicide)

cyperrnethrin (a pyrethroid insecticide;

mixtures of diastereomers)

methidathion (insecticide, acaricide)

metobenzuron (herbicide)

complex structures, both with respect to the carbon skeleton as well as with respect to the type and number of functional groups (Figs. 2.24 and 2.25). Obviously, assessing the environmental behavior of such compounds poses a particular challenge, since the mutual electronic and steric interactions of functional groups may influence the properties and reactivities of a compound in a way that cannot be a priori predicted from looking at the various fkctionalities in isolation. Nevertheless, with the general knowledge that we will acquire in Parts I1 and 111 on how structural moieties influence the environmental behavior of organic compounds, we should be able to put ourselves into the position to also tackle the assessment of such “complex” chemicals in the environment.

Some examples of compounds with structures that are somewhat more complex than the structures of the majority of the chemicals that we will encounter throughout the book include pesticides and pharmaceuticals. Such compounds are designed to be biologically highly active. Pesticides and pharmaceuticals are used by human society to exert a desired biological effect at a given target site. Obviously, from an environmental point of view, such compounds are of particular concern when they are present at locations where they are not supposed to be. For pesticides, this has been recognized for a long time, and therefore, there is quite a lot of information available on the environmental behavior of such compounds (e.g., Tomkin, 1994; Homsby et al., 1996; Montgomery, 1997). In contrast, for pharmaceutical chemicals, and in particular, for drugs and hormones used in human and veterinary applications,

Questions and Problems 51

HO

17-ethynyl estradiol (hormone, birth control pill)

H

O T N P >

0- 0

&OH 0

penicillin V (0-lactam antibiotic)

1 7a-acetoxy-progesterone (hormone)

levothyroxine (hormone)

ciprofloxacin (fluoroquinolone antibiotic)

clarithromycin (macrolide antibiotic)

Figure 2.25 Some examples of environmental aspects are only an emerging issue (Halling-Ssrensen et al., 1998; Pharmaceutical ing a somewhat more complex structure. Note that such compounds have been detected in waste waters and in surface waters (Ternes and Wilken, 1999).

exhibit- Ternes and Wilken, 1999; Sedlak et al., 2000; Kolpin et al., 2002).

Questions and Problems

Questions

Q 2.1

Which are the most common elements encountered in organic chemicals? What is a heteroatom in an organic molecule?

Q 2.2

Explain the simplest model used to describe the tendency of the various elements present in organic chemicals to undergo covalent bonding. For which elements is this simple model not strictly applicable?


Q 2.3

What does the structure of a given compound describe? What are structural isomers?

Q 2.4

What types of covalent bonds exist between the atoms present in organic molecules? What factors determine the strength of covalent bonds? Give some examples of very strong and very weak bonds.

Q 2.5

Which atoms present in organic molecules may exhibit different oxidation states? Explain in words how you assign the oxidation states to the different atoms present in a given molecule.

Q 2.6

What is stereoisomerism? What type of stereoismerism exists? Give some examples of different types of stereoisomers.

Q 2.7

In what context do you use the terms cis and trans? In what context do you use ortho, meta, and para?

Q 2.8

Explain the terms delocalized electrons, resonance, and aromaticity. Give examples of compounds for which these terms apply.

Q 2.9

What are the major sources of the BTEX compounds in the environment? Why are these compounds considered to be a problem?

Q 2.10

What are the major sources of PAHs in the environment? What are the major problems connected with this class of compounds?

Q 2.11

Why do so many high-volume production chemicals contain several halogen atoms? Why are polyhalogenated hydrocarbons rather persistent in the environment?

Q 2.12

Give at least five examples of polyhalogenated hydrocarbons and explain what kind of environmental problems they cause.

Q 2.13

How may an OH-group influence the behavior of an organic compound in aqueous solution?

Questions and Problems 53

Q 2.14

What is an ester function? Give some examples of environmentally relevant chemicals exhibiting carboxylic acid ester functions.

Q 2.15

How do amino groups and nitro groups affect the properties of (a) aliphatic, and (b) aromatic compounds?

Q 2.16

Name some important differences in the chemical nature of oxygen and sulfur. How is this reflected in sulfur-containing functional groups?

Q 2.17

For which major purpose(s) are sulfonic acid groups introduced into man-made organic chemicals?

Q 2.18

What are the characteristics and the major uses of surfactants? Give some examples of important commercially available surfactants.

Q 2.19

What is special about double bonds between oxygen and sulfur or phosphorus?

Q 2.20

What are the most common phosphorus-containing functional groups encountered in environmentally relevant compounds? What are the major uses of such compounds?

Problems

P2.1 Determining the Oxidation State of the Atoms in an Organic Molecule

Determine the oxidation states of the numbered atoms in the following organic molecules:

1

1 4 0

(b)

1

54 Questions and Problems

P2.2 Assessing the Number of Stereiosorners

Write down all possible stereoisomers of 1,2,3,5- and 1,2,4,5-tetrachlorocyclohexane. Which of them are chiral, that is, which ones exist as pair of enantiomers?

CI

1,2,3,5-tetrachlorocyclohexane 1,2,4,5-tetrachlorocyclohexane

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