V. V. Tronov.

On the Question of the Significance of the Electrochemical Character of Elements and Radicals in Organic Chemistry

In chemistry the notion of the electropositive and electronegative character of elements and radicals has long been established. An electropositive element (or radical) is one whose atom (or, respectively, group of atoms), upon electrolytic dissociation, forms a positive ion; and, conversely, an element forming a negative ion is called electronegative.

However, by far not all compounds are capable of decomposing into ions. Meanwhile the concept of electrochemical properties has proved so useful that it is desirable to extend it to all chemistry. Therefore, for determining (or at least for approximately estimating) the electrochemical character, indirect methods are used. One takes some type of compounds—alkalis or acids—that dissociate rather strongly, substitutes various atoms or radicals into them, and observes how the acidic or basic properties change. If the basic properties have increased, this means that the radical or atom taken is more electropositive than the one that was there previously. Thus, by replacing in ammonium hydroxide, \(NH_4OH\)—a fairly weak alkali—the hydrogen atoms by methyl groups, we have a series of transitions to tetramethylammonium hydroxide \((CH_3)_4NOH\), comparable in its properties with caustic soda (ratio of dissociation constants \(= 0.92 : 1\)). Hence methyl is more electropositive than hydrogen. If phenyl is put in place of hydrogen, a weaker alkali is obtained; triphenylamine \((C_6H_5)_3N\), where phenyl groups stand in place of all three hydrogens of ammonia, no longer combines either with water or even with strong acids; consequently, phenyl is more electronegative than hydrogen.

The same result is also given by comparing the strength of the corresponding acids. If in acetic acid \(CH_3 \cdot COOH\) hydrogen is replaced by methyl, the resulting propionic acid \(CH_3 \cdot CH_2 \cdot COOH\) is weaker than acetic acid; trimethylacetic acid \((CH_3)_3C \cdot COOH\) is still weaker, while phenylacetic acid \(C_6H_5 \cdot CH_2 \cdot COOH\) is stronger.

Using these criteria, it was possible to arrange the elements and radicals in a certain series. For organic radicals and inorganic radicals often encountered in organic chemistry, the series is approximately as follows:

\[ -NO_2,\ CN,\ COOH,\ Cl,\ Br,\ J,\ OH,\ldots\ldots C_6H_5,\ldots\ldots H,\ldots\ldots C_3H_7,\ C_2H_5,\ CH_3 + \]

The existence of a continuous series, without a sharp division into two parts, proves that in essence no element can be considered definitely electropositive or electronegative, but one can only say that one is more positive or more negative than another.

If into the molecule of some acid one of the members of the above-written series enters, and we replace it by another standing higher, then a stronger acid will be obtained; a radical standing lower weakens acidic properties. Thus, for example, the dissociation constants of acetic acid and some of its derivatives are:

It must be borne in mind, however, that the electrochemical action of a radical depends not only on its nature, but also on its position in the molecule. The cited-

…mentioned above, expresses, so to speak, an average distribution suitable for most cases. Sometimes rather large deviations are observed. For example, \(NO_2\) and \(CN\) prove to be more electronegative than halogens only when sufficiently removed from the acidic hydrogen; \(H—O—NO_2\) is a very strong acid, whereas \(H—O—Cl\) is a weak one, but HCl is much stronger than nitrous acid \(HNO_2\). Likewise hydroxyl is in general considerably more electronegative than hydrogen; however, para-oxybenzoic acid (dissociation constant 0.00286) is twice as weak as benzoic acid (0.0060).

The electrochemical character of atoms and atomic groups must, of course, influence not only the ability of a substance to undergo electrolytic dissociation, but also its other properties.

On it must depend, for example, the rate of breakdown of molecules in all sorts of reactions.

Putting in order the material bearing on the question of the formation of a double bond by the splitting off of certain groups,¹ I noticed a clear dependence between the direction and ease of cleavage and the electrochemical character of the parts of the molecule.

It has long been known that if, in some reaction, a compound is obtained in which two hydroxyls are attached to one carbon, then almost always water is immediately eliminated and a carbonyl group is formed. Thus, the hydrogen of one hydroxyl leaves together with the other hydroxyl. In exactly the same way, hydrogen chloride is split off from a hydroxyl and chlorine. If a hydroxyl and chlorine stand side by side, then it is not water but hydrogen chloride that leaves, i.e. chlorine is torn off more readily than hydroxyl:

\[ CH_3—C\begin{matrix} OH\\[-2pt] OH\\[-2pt] Cl \end{matrix} \;\longrightarrow\; CH_2—C\begin{matrix} OH\\[-2pt] O \end{matrix} + HCl . \]

Having examined many such reactions,² one may note that it is always the most electronegative atom or radical that leaves. For the time being I have compiled the following series (still very approximate and incomplete):

\[ Cl,\ O.CO.CH_3,\ NH_3,\ Cl,\ OH,\ OC_2H_5,\ CCl_3,\ COOH,\ NH_2,\ C\!\begin{matrix}O\\[-2pt] \| \\[-2pt] O\end{matrix}\!—C—,\ (HO)_3C_6H_2—,\ C_6H_5,\ CH_3.COOK,\ H,\ COONa,\ CH_3,\ OCaO,\ ONa. \]

The sequence is very close to that which we found above in determining the electrochemical character from the dissociation constant.

If, at one carbon atom, there are a hydroxyl and three further atoms or radicals of some kind, then the one of them that stands higher in this series is split off more readily.

The upper members of the series split off even at low temperature; compounds with chlorine and hydroxyl at one carbon cannot exist at all under ordinary conditions; compounds with two hydroxyls, on the other hand, are sometimes rather stable; methyl is split off (the preparation of methane from sodium acetate) at high temperature, almost on ignition.

Alongside the principal reaction, accompanying and side reactions also proceed, since ordinarily other atoms or groups may also be split off (depending on the difference in electrochemical character). If two reactions are possible and both are reversible, then a definite equilibrium is established. A typical example is the formation and saponification of esters:

\[ R—C\begin{matrix} O\\[-2pt] OH \end{matrix} + R'OH \;\rightleftarrows\; R—C\begin{matrix} OR'\\[-2pt] OH\\[-2pt] OH \end{matrix} \;\rightleftarrows\; R—C\begin{matrix} OR'\\[-2pt] O \end{matrix} + HOH . \]

¹ See the preceding article, “On reactions of cleavage and formation of a double bond.”
² Some of them are given in the preceding article.

The more positive the radical, the more firmly the group \(OR'\) is held, and the more the equilibrium is shifted toward ether formation. Methyl, ethyl, and other residues of fatty hydrocarbons are more electropositive than hydrogen; therefore, in the interaction of equal molecular quantities of acetic acid and methyl alcohol, 69.5% of ether and water is formed, and with ethyl alcohol 66.6%; with phenols the reverse reaction predominates, since phenyl and the like are more negative than hydrogen; here the limit is reached at 1.5–2% of ether.

If we go further in the direction of increasing the electropositive character and, in the group \(OR'\), place some typical metal in place of the hydrocarbon radical, then this whole group must be held still more firmly.

And indeed, for example,

\[ \mathrm{CH_3 - C\begin{matrix} \mathrm{ONa}\\[-2pt] \mathrm{OH}\\[-2pt] \mathrm{OH} \end{matrix}} \]

when heated gives off almost exclusively water; hydrolysis, i.e. decomposition into

\[ \mathrm{CH_3 - C\!\left(=O\right)OH} \]

and \(\mathrm{NaOH}\), occurs only to an insignificant extent. Likewise from

\[ \mathrm{CH_3 - C\begin{matrix} \mathrm{ONa}\\[-2pt] \mathrm{OH}\\[-2pt] \mathrm{OCH_3} \end{matrix}} \]

(the product of addition of caustic soda to acetic-methyl ester), \(\mathrm{CH_3OH}\) is practically reliably split off (saponification of esters by alkalis).

An interesting conclusion is obtained by comparing not the methyl of sodium (which, of course, is more electropositive), but the group \(ONa\). The compound

\[ \mathrm{CH_3 - C\begin{matrix} \mathrm{ONa}\\[-2pt] \mathrm{ONa}\\[-2pt] \mathrm{OH} \end{matrix}} \]

on heating splits off \(\mathrm{CH_4}\), i.e. \(ONa\) proves to be, as it were, positive \(\mathrm{CH_3}\). True, this cannot be regarded as fully proven, since precisely this reaction may go to completion simply because one of the products is removed (gas), whereas the products of the other possible reaction, \(\mathrm{CH_3COONa}\) and \(\mathrm{NaOH}\), remain together and continue to take part in the process. But in any case the \(ONa\) radical proves comparable in its electrochemical properties with methyl.

Thus, the study of cleavage processes accompanying the formation of a carbonyl bond gives us a certain means of judging the electrochemical character of a large series of the most diverse elements and radicals. For the present one can speak only of an approximate estimate, not of an exact determination, but it seems to me that this method is capable of further development. It is necessary, using suitable examples, to carry out a series of measurements of reaction rates. Such are the works of N. A. Menshutkin on the esterification of alcohols, but the question must be posed more broadly, and observations must be extended to the most diverse radicals. Then it will be possible to derive quantitative dependencies, to connect them with what is given by the study of electrolytic dissociation, etc.

Knowing these dependencies, having compiled a series several times more complete than the one I have sketched, and supplying it with figures that express, in one way or another, the strength of the bonds, one may hope, on the one hand, to penetrate more deeply into the structure of molecules (since the strength of the bond depends on the arrangement of the electrons and atomic nuclei), and, on the other hand, to apply this in practice, to laboratory work.

The outlined general regularity embraces an enormous field, precisely in organic chemistry, where crude empiricism still prevails. It has been found, for example, by experiment, that in order to split off carboxyl from acids it is best to heat them with alkalis or with lime; it is known that esters are saponified by alkalis, and amides by alkalis and acids, etc. These facts have partly found an explanation for themselves, but some of them have remained unclear; for example, all

Divivita’s reaction, discovered at one time by V. N. Ipatiev, of the catalytic oxidation of alcohols to aldehydes.

Now we have the following:

The decomposition of a carboxylic acid into a hydrocarbon and \(CO_2\) consists in the fact that the hydrogen of the hydroxyl is split off and the bond between the hydrocarbon residue and the carboxyl carbon is broken:

\[ \boxed{R{-}C\left(\begin{array}{c} O\\[-2mm] \|\\[-2mm] O{-}H \end{array}\right)} \;\longrightarrow\; RH + C\left(\begin{array}{c} O\\[-2mm] \|\\[-2mm] O \end{array}\right). \]

With a sufficiently electronegative radical this is achieved by simple heating.

Therefore, for example, gallic acid

\[ \begin{array}{c} \mathrm{HO}\\[-1mm] \mathrm{HO}\!\! \end{array} \left( \begin{array}{c} \text{benzene ring} \end{array} \right) {-}\mathrm{COOH} \quad \begin{array}{c} \\[-7mm] \mathrm{HO} \end{array} \]

loses \(CO_2\) rather readily.

But benzoic acid can hardly be decomposed in this way. For the decomposition of this acid, and also of acetic acid, advantage is taken of the fact that one can facilitate the formation of a carbonyl bond by destroying an already existing such bond (a double bond produces a certain strain in the molecule, and therefore, if it exists, a second one at the same atom is formed with difficulty). This can be achieved by the preliminary addition of some third molecule. If water is added, nothing comes of it.

\[ \mathrm{CH_3{-}C(OH)_3} \]

will of course be formed, but it is so unstable that it cannot even be isolated. In doing so it loses water back again, since the hydroxyl is more electronegative than methyl. To split off methane, one must leave one hydroxyl and replace the other two by groups standing higher in the series written above, i.e. by \(ONa\) groups. Then, on heating, methane escapes and sodium carbonate remains.

In general, in order to break the bond between the carbonyl carbon and the rest of the carbon chain of the molecule, the carbonyl bond is destroyed, replacing the oxygen by two univalent groups. Of these, one must contain an element that is easily split off from oxygen and at the same time attaches itself firmly to carbon, i.e. hydrogen; while the other itself must remain attached to the carbon before some end of the chain. These requirements are satisfied by the addition of alkalis.

Therefore acids decompose when heated with alkalis; \(\beta\)-diketones are likewise cleaved by alkalis.

\[ \mathrm{CH_3{-}C(=O){-}CH_2{-}C(=O){-}CH_3 + KOH} \;\longrightarrow\; \mathrm{CH_3{-}C\left(\begin{array}{c} OK\\ OH \end{array}\right){-}CH_2{-}C(=O){-}CH_3} \;\longrightarrow \]

\[ \longrightarrow\; \mathrm{CH_3{-}C\left(\begin{array}{c} OK\\ \|\\ O \end{array}\right) + CH_3{-}C(=O){-}CH_3}. \]

Chloral readily adds water, but on simple heating gives it back again, since the hydroxyl is more electronegative than trichloromethyl:

\[ \mathrm{CCl_3{-}C\left(\begin{array}{c} OH\\ OH\\ H \end{array}\right)} \;\longrightarrow\; \mathrm{CCl_3{-}C\left(\begin{array}{c} O\\ \|\\ H \end{array}\right) + H_2O}. \]

If, however, in one hydroxyl the hydrogen is replaced by a metal—usually calcium is taken—then the splitting-off of chloroform very readily occurs:

\[ \mathrm{CCl_3{-}C\left(\begin{array}{c} OH\\ OH\\ H \end{array}\right) + H.O.CaOH} \;\longrightarrow\; \mathrm{CCl_3{-}C\left(\begin{array}{c} OH\\ O.CaOH\\ H \end{array}\right) + H_2O,} \]

and then:

\[ \mathrm{CCl_3{-}C\!\begin{matrix} /\,OH\\ -\,O\cdot Ca\cdot OH\\ \backslash\,H \end{matrix}} \;\longrightarrow\; \mathrm{CCl_3H + C\!\begin{matrix} /\,O\\ -\,O\cdot Ca\cdot OH\\ \backslash\,H \end{matrix}} \]

Likewise, the preparation of iodoform:

\[ \mathrm{CI_3{-}C\!\begin{matrix} /\,OH\\ -\,OH\\ \backslash\,H \end{matrix} + KOH \;\longrightarrow\; CI_3{-}C\!\begin{matrix} /\,OH\\ -\,OK\\ \backslash\,H \end{matrix} + H_2O \;\longrightarrow\; CI_3H + C\!\begin{matrix} /\,O\\ -\,OK\\ \backslash\,H \end{matrix} + H_2O} \]

We have already analyzed the saponification of esters and clarified why it is accelerated so readily by alkalis. The hydrolysis of amides, however, proceeds more rapidly both with alkalis and with acids.

Here the matter is as follows:

The ester molecule, e.g.

\[ \mathrm{CH_3{-}C\!\begin{matrix} /\,OC_2H_5\\ \backslash\,O \end{matrix}} \]

can be affected almost not at all by an acid—say hydrochloric acid. If it adds to the double bond, then the resulting product

\[ \mathrm{CH_3{-}C\!\begin{matrix} /\,OC_2H_5\\ -\,Cl\\ \backslash\,OH \end{matrix}} \]

again splits off the same HCl with extraordinary ease.\(^1\)

With amides still another change is possible. If acetamide simply adds water,

\[ \mathrm{CH_3{-}C\!\begin{matrix} /\,O\\ \backslash\,NH_2 \end{matrix} + HOH \;\longrightarrow\; CH_3{-}C\!\begin{matrix} /\,OH\\ -\,OH\\ \backslash\,NH_2 \end{matrix}} \]

then here, attached to the central carbon, there are methyl (very difficult to tear away), an amino group, and two hydroxyls. The latter are split off most easily, since OH is more negative than NH\(_2\). In order to liberate ammonia, one can make one hydroxyl positive by replacing its hydrogen with sodium. Then—apart from that hydroxyl which gives up its hydrogen—the most electronegative residue is the amino group. It therefore leaves:

\[ \mathrm{CH_3{-}C\!\begin{matrix} /\,O\\ \backslash\,NH_2 \end{matrix} + NaOH \;\longrightarrow\; CH_3{-}C\!\begin{matrix} /\,ONa\\ -\,OH\\ \backslash\,NH_2 \end{matrix} \;\longrightarrow\; CH_3{-}C\!\begin{matrix} /\,ONa\\ \backslash\,O \end{matrix} + NH_3} \]

Everything is just as with esters.

But in amides we have still another means of splitting off the group NH\(_2\)—even when two hydroxyls exist alongside it. We make this group more negative than hydroxyl by adding HCl:

\[ \mathrm{CH_3{-}C\!\begin{matrix} /\,OH\\ -\,OH\\ \backslash\,NH_2 \end{matrix} + HCl \;\longrightarrow\; CH_3{-}C\!\begin{matrix} /\,NH_3Cl\\ -\,OH_2\\ \backslash\,OH \end{matrix} \;\longrightarrow\; CH_3{-}C\!\begin{matrix} /\,OH\\ \backslash\,O \end{matrix} + NH_4Cl} \]

Ammonium chloride is split off.

The Ipat’ev reaction is explained as follows.

In an alcohol—let us take the most ordinary, ethyl alcohol—with the carbon there are, besides hydroxyl, two hydrogens and methyl. Methyl is more electropositive than hydrogen, and therefore, if the reaction is carried out under such conditions that for some reason not the whole hydroxyl, but only its hydrogen, must be split off, then together with this hydrogen atom not methyl leaves, but one of the hydrogens of the alcohol group:

\[ \mathrm{CH_3{-}C\!\begin{matrix} /\,OH\\ -\,H\\ \backslash\,H \end{matrix} \;\longrightarrow\; CH_3{-}C\!\begin{matrix} /\,H\\ \backslash\,O \end{matrix} + H_2} \]

\(^1\) Some acceleration is possible because HCl can add complexly to the oxygen of the group OR and thereby make this group more electronegative. But in any case such addition takes place considerably less readily than in amides, so that acids accelerate the reaction less here.

Here, finally, is one more example illustrating the influence, on the decomposition of a molecule, of the electrochemical character of its parts. This is the different behavior of oxyacids when boiled with dilute sulfuric acid.

\(\alpha\)-oxyacids, for example lactic acid, give an aldehyde and formic acid:

\[ \mathrm{CH_3 - C\!\left(\begin{array}{c} \mathrm{H}\\[-2pt] \mathrm{OH} \end{array}\right) - COOH} \;\longrightarrow\; \mathrm{CH_3 - C\!\left(\begin{array}{c} \mathrm{H}\\[-2pt] \|\\[-2pt] \mathrm{O} \end{array}\right)} + \mathrm{H \cdot COOH}. \]

\(\beta\)-oxyacids split off water, passing into an unsaturated acid:

\[ \mathrm{CH_2OH - CH_2COOH} \;\longrightarrow\; \mathrm{CH_2 = CH - COOH + H_2O}. \]

Let us see how the reaction must proceed in the one and in the other case. The hydroxyl group may depart either as a whole or only by its hydrogen. If the hydrogen departs, then together with it there must be split off the most negative of the radicals attached to the same carbon; moreover, the more electronegative this radical is, the more readily it is torn away. In lactic acid this is \(\mathrm{COOH}\), while in \(\beta\)-oxypropionic acid it is the group \(\mathrm{CH_2 \cdot COOH}\). It is clear that carboxyl is the more negative (this is evident at least from the fact that if, in acetic acid, we replace hydrogen by carboxyl, malonic acid, \(\mathrm{COOH \cdot CH_2 \cdot COOH}\), is obtained, with a dissociation constant of \(0.158\); whereas if we put in its place the group \(\mathrm{CH_2 \cdot COOH}\), the resulting glutaric acid has a constant of only \(0.0065\)); and consequently, in this direction the \(\alpha\)-acid will react more readily. In addition, what is negative (i.e. not having a free negative charge, but only possessing affinity for electrons) is not the ion \((-\mathrm{COO})'\), but precisely carboxyl—\(\mathrm{COOH}\), so that the acid should split more readily not in the form of a salt (strongly ionized), but in the form of undissociated molecules of the free acid—these being present to the greatest extent in acid solution, i.e. in the presence of some strong mineral acid.

When the entire hydroxyl is torn away, then together with it there usually departs hydrogen from the neighboring carbon. It is known that, in general, the hydrogen atoms of a secondary link are more mobile than those of a primary one. In lactic acid the hydrogen of the methyl group would have to take part in the reaction; in \(\beta\)-oxypropionic acid, the reacting group is \(\mathrm{CH_2}\), and one standing next to carboxyl at that, so that here the second reaction must proceed considerably faster. And again sulfuric acid accelerates the decomposition: hydroxyl can be esterified, i.e. replaced by the group \(\mathrm{O \cdot SO_2OH}\), and this group is in general torn away much more readily; \(\mathrm{CH_3 \cdot CH_2 \cdot SO_2 \cdot OH}\) gives ethylene and sulfuric acid already at a temperature only a little above \(100^\circ\), whereas \(\mathrm{CH_3—CH_2OH}\) does not decompose into ethylene and water (without catalysts) even at \(350^\circ\).

The present work, together with the preceding one, represents an attempt to introduce some order into the study of a vast and very scattered group of reactions, for the present chiefly of organic chemistry—reactions sometimes, at first glance, very strange, and often not finding for themselves any explanation.

Here I am giving, in fact, not a theoretical explanation, but only a classification. For me it has followed from consideration of the electronic structures of atoms and molecules—structures based on certain hypothetical assumptions. These hypotheses may in some of their parts prove to be incorrect, but the method of considering reactions derived from them may remain in the form of a scheme (for I have not said a word about electrodes), just as the electron-ionic theory is not entirely correct, yet the classification of oxidation–reduction reactions based on it nevertheless embraces a very large number of facts very well, and awaits only a more correct theoretical justification.