VALENCE and MOLECULAR STRUCTURE JOHN DEVRIES Knox College, Galesburp, Illinois
This article deals with the general problem of modernieing the course i n elementary chemistry. A valence bond is defined as being composed of two electrons whose spins are paired. The different types of possible walence bonds are d e h d and the applications and advantages of ap-
proaching our problems i n teaching valence and molecular structure from this viewpoint are offered. A modification of the definition of valence offered by Flood i n the JOURNAL OF CHEMICAL EDUCATION, 12, 132 (1935) to include the different types of valence bonds is suggested.
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HE question of what to teach the student of chem- is in agreement with the Pauli exclusion principle, which istry concerning valence and the structure of mole- states that there cannot be two electrons in the atom cules has interested the author for some time. An which have the same values for all the four quantum examination of most elementary texts reveals a brief numbers. This gives us the possibility of two electrons discussion of such topics as radioactivity and radium having the same principal quantum number, n, the disintegration, positive-ray analysis, and usually a dis- same amount of angular orbital momentum, I, and the cussiou of the Bohr theory with its electron orbits, a same magnetic quantum number, ml, but different theory which prohibits the formation of even such an spins. For example, in helium where the values of n , elementary substance as molecular hydrogen. One I, and ml are the same for both electrons we bave a pairfinds but little concerning the nature of the valence ing of the electron spins which prohibits compound bond, and frequently the discussions attempt to carry formation. Approximately 456,000 calories per gram over into the field of inorganic chemistry the valence atom are required to uupair the electron spins and raise theories of the organic chemist without modifying them. the principal quantum number to two. This explains An examination of the average structural formulas in why helium does not ordinarily form compounds. Two most elementary texts substantiates this statement. unpaired electrons in different atoms, however, may be Flood (1) in a recent article errs likewise when he de- readily coupled together to form a valence bond and the fines valence on the basis of typical C-C- or C-H-like electrons then oscillate b e t ~ e ~ t atoms h e and appear bonds and defines the valence of positive or negative to be part of both. rn the case df atoms having an odd ions as a number representing the positive or negative number of electrons each atom must contain a t least unit charges associated with the ion. Although the one free magneton, and this explains why nitrogen latter part of his definition is correct, strict adherence and chlorine, for example, exhibit odd-numbered vato the first part leads him into difficulty. His state- lences more frequently than do even-numbered atoms, ments that chlorine in the perchldrate ion has a valence which usually have an even-numbered valence. Lewis of four and that sulfur in the sulfa@ ion also has a refers to this pairing of electrons to form a valence bond valence of four are illustrations of the inorganic chemist as the rule of two. Although it is not advisable to applying the valence theories of the organic chemist and teach the fundamental concept of the spinning electron failing to distinguish between the ditferent types of to an elementary student, the concept that a valence bonds in the molecule. Everyone knows that when the bond is composed of two electrons should not be diffiorganic chemist states that carbon has a valence of four cult to accept. The student -who is intelligent enough he means that we can attach other atoms to carbon by to wonder how, on the basis of Coulomb's law, two electrons can exert an attraction for each other rather than means of four d e h i t e valence bonds. G . N. Lewis and his students bave given us much of a repulsion should also be intelligent enough to underthe material which could well be used as a basis for stand an explanation of magnetic attraction in the electeaching the structure of molecules and for unifying the trons. The physics course in which he learned Couvalence theories. Lewis (2) defines a valence bond as a lomb's law will also have taught him Oersted's law that bond composed of a pair of electrons whose magnetic a flow of electricity (movement of electrons) causes moments neutralize each other or, in terms of the newer magnetic phenomenon. theories, the bond is composed of two electrons whose The rule of four, which states that we can attach four spins are paired. By a spinning electron is meant one definite valence bonds or groups to the central atom, which spins on its own axis with a great enough velocity and the rule of eight, which states that in any stable, to give it a magnetic moment equal to the Bohr magne- closed molecule each atom may be surrounded by a ton. The concept of the pairing of two electron spins maximum of eight electrons, can be readily presented. 320
For a complete discussion of these rules and exceptions the reader is referred to Lewis (2) and Rodebusb ( 3 ) . Another concept which is vital for the student to know is the varying attraction which different atoms have for electrons. We usually define an electropositive element as one which tends to give up its electrons and an electronegative element as one which exerts a strong attraction for electrons. Rodebush (3) has suggested that this varying attraction of the atoms for electrons be expressed in terms of the field about the atom. He points out that the terms electropositive and electronegative arose before the discovery of the electron and that these terms have been used in so many ways in connection with the properties of molecules that they are ambiguous. The field in the neigh-. borhood of all atoms, with the exception of the inert gases, is positive, that is, the atoms all attract electrons, but in a varying degree. While i t is not possible to compare quantitatively the atoms in different groups, an approximate measurement of the "positivity" of the atoms in a single group may be found by dividing the number of valence electrons by the number of shells. The atom will be more positive as the value of the effective nuclear charge increases. Thus chlorine is less positive than fluorine, but in comparison with sodium it would be considered strongly electropositive. To sum up his argument, Rodebush states, "This is contrary to the current usage in inorganic chemistry where chlorine is usually termed negative or electronegative, but the current usage arose before the discovej of the electron and there seems to be no good reason for adhering to an illogical terminology." The objection may be raised that this will lead to confusion since we have alwavs . regarded elements such as chlorine, for example, negative, since they migrate to the positive pole, and other elements positive, since they migrate to the negative pole. However, we are then dealing with ions and it can readily be pointed out that the element has either gained or lost an electron and hence bas formed a charged ion. The author has found this concept to he rather teachable and not confusing after the student has mastered this approach. q If we consider the valence bond to he composed of two electrons whose spins are paired, rather than a rigid, unyielding link, we have the possibility of a number of different types of bonds. The first type is sometimes called a polar bond and occurs when an atom takes an electron completely away from another atom and forms an ion. Rodebush (3) has pointed out that this appears to be a misnomer and that the term polar molecule is more suitable in this case. Langmuir, in developing his electronic theory of chemical combination, called this type of bond or union "electrovalency" and it seems desirable to speak of this type of combination in the latter way. A second type of combination results when each element contributes one electron in the formation of the bond and the two electrons then oscillate between the atoms, appearing to belong to both. In some cases the electrons shift toward one atom in order to give i t a
resultant negative charge and the other a resultant positive charge, although it is not correct to speak of any atom having an integral number of units of charge unless one of the atoms takes the electrons completely away from the other and forms an ion. This type of combination has been termed "covalency" by Langmuir and is also designated occasionally as a non-polar bond. In general, according to the Fajans theory, ions of high valency, whether positive or negative; favor the formation of covalent rather than electrovalent linkages. When the cation is small and the anion large an electrovalency tends to pass over into a covalency. Beryllium chloride, for example, exhibits mainly a covalent linkage in the fused state, as indicated by its resistivity, while in water solution, although somewhat hydrolyzed, it becomes a fairly good conductor, indicating the presence of an electrovalency. The metals in the A sub-group of Group I in the periodic system tend to form electrovalent compounds readily, since they need to lose only one electron to form an outside shell of eight electrons. A third type of combination, which Sugden has termed the semipolar bond, occurs when a molecule or ion contributes both of the electrons in forming the bond. This type of bond is a special type of covalency which Sidgwick has termed a coordinate covalence. There are two mechanisms proposed to illustrate the formation of this bond. The fist suggests that the two electrons are simply donated by the one molecule or ion and the electrons are then shared between the two, thus:
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Here A is termed the donw &d B the acceptor molecule or ion. A cannot be an atom, since if it were it would he an inert gas. Such a molecule, AB, is expected to have a certain amount of polarity; hence the term semipolar bond. Another viewpoint suggests that an electron is first transferred from A to B to yield-A+ and B-, thus:
The two particles now unite by means of a covalent bond. Looking a t it from this point of view the semipolar hond would be composed of one electrovalent and one covalent bond and would appear to he a type of double bond. The bond is usually represented either A+-B- or A + B, the arrow indicating the direction of the transferred electrons. The question of the proper name for this type of bond is still a matter of debate. As pointed out before, Sidgwick suggests the name "coordinate covalency," and recently W. A. Noyes has suggested the term "semi-ionic" valence. The author prefers the term semipolar hond, since Sidgwick's term is really an application of Werner's coordination valency. Werner deals with the union of two substances capable of sepa-
rate existence, and represents the bond by an arrow, while Sidgwick applies the term to all semi-polar linkages. I t would be less confusing if we restricted the term coordinate covalence to Werner's original meaning, since the normal semipolar bond is not readily broken while in the coordinate type it can be broken more easily, since the two molecules are capable of separate existence. The definition sueeested bv . Noves also implies a concept of polarity and may be substituted for the term semipol& bond if so desired. ~t may he added, parenthetically, that a semipolar bond behaves ~hvsicallvas thoueh " it were a sin& " bond, although it is permissible to count it numericdly as twd linkages when computing classical valency. ~b~ average value of the mean &toring forces fo; various linkages as calculated from molec&r spectra data indicates that there is really no fundamental difference, stereochemically a t least, between an ordinary covalency and a semipolar or coordinate covalency. Hence, in writing structural formulas for inorganic compounds we should not merely present the usual formula. For ex-
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