&darn Structural Chemistry Nobel Lecture - Gecember 1954 by Linus Patiling A century ago the structural theory of organic chemistry was developed. Frankland in 1852 suggested that an atom of an element has a definite capacity for combining with atoms of other elements - a definite valence. Six years later Kekule/ and Couper, independently, introduced the idea of valence bonds between atoms, including tht formation of bonds between two c&r-bon atoms; and suggested that carbon is quadri- valent. In 1861 hutlerov, making use for the first time of the term of a compound "chemical structure" stated clearly that the properties/are determined I by its moleoalar structure and reflect the way in which atoms are bonded to one another in the molecules of the compound. The develowent of the structure theory of organic chemistry then progressed rapidly, and this theory has been of inestimable value in aiding organic chemists to interpret their experimental results and to new experiments. r\ -2. A most important early addition to organic structure theory was made by the first Nobel Laureate in Chemistry, van't Hoff, who in 1874 reco$nized that the optical activity of carbon compounds can be ex- plained by the postulate that the four valence bonds of the carbon atom are directed in space toward the corners of a tetrahedron. The structure theory of inorganic chemistry may be said to have been born only fifty years ago, when Werner, Nobel Laureate in Chemistq- in 1913, found that the chemical composition and properties of complex inorganic substances could be explained by assuming thet metal atoms often coordinate about themselves a number of atoms different from their valence, usually four atoms at the corners%of a tetrahedron 62/ or a square coplanar with the central atom, or six atoms at the corners h of an octahedron. His ideas about the geometry of inorganic complexes were completely verified twenty years later, through the application of the technique of x-ray diffraction. After the discovery of the / electron many efforts were made an electronic to develop 6 theory of the chemical bond. A great contribution was made in 1916 Gilbert by ~&&VW& Newton Lewis, who proposed that the chemical bond, such as the single bond between two carbon atoms or a carbon atom and a hydrogen atom represented by a line in the customary structural formula for ethane, consists of a pair of electrons held jointly by the two atoms that are bonded together. * Lewis also suggested that m atoms tend to assume the electronic configuration of a noble gas, through the sharing of electrons with other atoms / or through electron transfer, and that the eight outermost electrons in an atom with a noble-gas electronic structure are arranged tetrahedrelly in pairs P- about the atom. Applica- tions of the theory and additional contributions were mace by many chemists, including Irving Langmuir and Nevil Vincent Sidgwick. &F&e+ After the discovery of quantum mechanics in 1925 / it became evident that the for X the theory of molecular structure. It also soon became evident that these equations, such as the Schrbdinger wave equation, cannot be solved rigorously for any but the simplest molecules. The development of the theory of molecular structure and the nature of the chemical bond during the past twenty-five years has been in considerable part empirical - based upon the facts of chemistry - but with the interpretation of these facts greatly influenced by quantum mechanical. principles and concepts. The solution of the wave equFtion for the hydrogen molecule- ion @ $. Durrau (Det Kgl. Danske Vid. Selsk. Math&@* Meddelel$er, I, 14 (1927)) completely clarified the question of the nature of the one-electron bond in this molecule-ion. &uminating quantum mechanical discussions of the shared-electmn-pair bond in the hydrogen molecules then simultaneously * were eee~ published, on*-- Iy Heitler and London (Z. 3!$ Physik, & 455 (1927)), and the other by E. U. Condon (Proc. Nat. Acad.Sci. U.S., 2, 466 (1927)). In the m me&-e& L8 approximate solution of the wave equation for the hydrogen molecule by Ueitler and London a wave function is used that WM+WWM, requires the ~$?zct,$rb be separated, each being close to one of the two nuclei. The treatment by Condon permits the electrons to be dis- tributed between the two nuclei independently of one another, each occupying a wave function similar to Burrau's function for the hydrogen- molecule ion. Condon's treatment is the prototype of the molecular- orbital treatment that has been extensively applied in the discussion of aromatic and conjugated molecules, and Heitler and London's treatment is the prototype of the valence-bond method. `#hen the effort is made to refine the two treatments / they tend to become identical. @' These early applications of quantum mechanics to the problem of the nature of the chemical bond made it evident that in general a co t valent bond, involving the sharing of a pair of electrons between two atoms, can be formed if two electrons are available (their spins must be opposed, in order that the bond be formed), and if each atom has avail- able a stable electronic orbital for occupancy by the electrons. The equivalence of the four bonds formed by a carbon atom, which had become a part of chemical theory, was not at first easily reconciled with the quant d escription of the carbon ato as having one 2p orbital and three 2p orbitals in its outer shell. The solution to this difficulty was obtained when it was recognized that as a result of the resonance phenomenon of quantum mechanics a tetrahedral arrangement of the f%x four bonds of the w carbon atom is 1 / achieved. The'carbon atom can be described as having four equivalent 1. Linus Pauling,-lThe Shared-Electrop Chemical Bond, Proc. Nat. Acad. sci. U.S., & 359-362 (1928). -- --- tetrahedral /bond orbitals, which are hybrids of the p and p orbitals. Further study of this problem led to the discovery of many sets of hybrid bond orbitals, which could be correlated with bond angles, magnetic moments, and other molecular properties. 2 In particular it was found that sp3,w 0-o-w 2. Linus Pauling, The **ature of 9 Chemical Bond. Application of Results Obtained from the G&ir&nn Mechanics and from a Theory of m- --0 Parsmagnetic Susceptibilitv to the Structure of Molecules, ifn.J. Am. Chem. Sot., WI 53, 1367 -J-400 0931). o----m cdidA&+ dsp2, and d2s$ hybrid orbitals correspond respectively to the tetrahedral, square planar, and octahedral configurttions of u inorganic complexes that W seen as to the r\ utilization of atomic orbit&s in bond formation can be drawn from Ym--+-b experimental values of magnetic moments. +he theory of the dsp2 square complexes of bipositive nickel, palladium, and platinum requires th5.t these substances be diamagnetic. The square complexes of bipositive -7. palladium and platinum had been recognized by Werner and their struc- ture verified w Dickinson (J. Am. Chem. Sot., 4, 2404. (1922)); but the assignment of the square configuration to the com$exes of nickel which are diamagnetic had not been made until the development of the new theory. Further detailed information about the chemical bond #!esulted i from a consideration of the energy of single bonds in relation to relative the/electronegativitj of the bonded atoms.3 It was found that the 3. Linus pauling, m Nature of the Chemical Bond. Iv. The bnergv of Single m and the helative Electronegativity of Atoms, -- J. Am. Chem. Sot., 5!t, 3570-82 (1932 1. -0-00-m elements can be assigned electronegativity values such as to permit the rough prediction of the heats of formation of compounds to which chemical structures involving only single bonds are conventionally assigned, and that many of the properties of substances can be discussed in a simple way with the use oi the electronegativity values of the ele:ments. -8. The idea that the properties of many organic compounds, especially the aromatic corn-pounds, cannot be simply correlated with a a somewhat more corn lex single valence-bond structure, but require J the assignment of an P electronic structure,was developed during the period 1923 to 1926 by m&n- a number of chemists, including Iowrg, Latworth, Robinson, and Ingold in England, Lucas in the United States, and Arndt and Eistert in Germany. It was recognized that the properties of aromatic and conjugated molecules can be described by the use of two or more valence-bond structures, w as reflected in the names, the theory of mesomerism and the theory of intermediate e&gee states, proposed for the new chemical theory. In 1931 Slater, E. Hiickel, and others recognized that these theories can be given a quantum mechanical interpretation: an approximate wave the sum of function for a molecule of this sort can be set up as/a-4&+ee+&&&en corresponding to the individual valence-bond structures. %e-ne+m&-&&e-&-&e The /molecule can then be described as e-&b&& having a structure that is a hybrid of the individual valence-bond structures, or as resonating among these structures, and the theory itself is now usually called the resonance theory of chemical structure. Vermany quantitative calculations, approximate solutions of the wave equation, for aromatic and conjugated molecules have been made, with results that are in general in good agreement with experiment. Perhaps more important than the quantitative calculations is the possibility of prediction by simple chemical arguments. For example, the amide group, an important structural feature of proteins, can be described as resonating between two structures, ?I one with the L double bond between the carbon atom and the oxygen atom, and the other with the aouble bond between the carbon atom and the nitrogen atom: ii (7 1 v \ / fi \ / ti ,* c o\/ c-- PJ // \ / / 0 . \ / *' I -e D c, b' --w-w vk , IiLnus Pauline, Interatomic Distances in Covalent Holecules -.- - & Resonance between Two o]: m Lewis Electmnic Structures, Pmt. Nat. Acad. Sci. U.S., 18, 293-297 (1932). -lo- General arguments about P the stability of alternative structures indi- cate / that the structure with the double bond between carbon and oxygen should contribute somewhat more to the normal state of the amide group acquaintance than the other structure; experience with other substances and/with the results of quantum mechanical calculations suggests the ratio 60$:40$ for the respective contributions of these structures. A 40% contribution of the structure with the double bond between the carbon atom and the nitrogen atom would confer upon this bond the property of planarity of the group of six atoms; the resistance to deformation from the p&anar configuration would be expected to be 40% as great as for a molecule such as ethylene, containing a pure double bond, and it can be calculated 4 3O that rotation of one end by w relative to the other end would in- A traduce a strain energy of 100 ca mole. Y The estimate of 40% double-bond character for the C-N bond is supported by the experimental value of the bond lajbr length, 1.32 it. interpreted with the aid of the empirical relation between double-bond character and interatomic distance. c 0 &. Linus Pauling and L. 0. Brockway, and J. Y* Beach, -ll- The Denendence of Interatomic Distance u S&g&4&4 Single Bond-Double Bond --- Resonance, J. Am. Chem. Sec., 5'& 2705 (1935). Knowledge of the structure of amides, and also of the amino acids, provided by the theory of resonance and verified by extensive careful experimental studies made'by R. B. Corey and his coworkers, has been of much value in the determination of the structure of proteins. In the description of the theory of resonance in cbamistry B i, there has been a perhaps unnecessarily strong emphasis on its ar&itrqy character. &M-I It is true, of course, that a description of the benzene molecule can be given, in quantum mechGnica1 language, without any reference to the two Kekule/structures, in which double bonds and single bonds alternate in the ring. An approximate wave function for the beni;ene molecule may be formulated by adding together two functions, representing the two Kekul< structures, and adding other terms, to make the wave function approximate the true wave function for the molecule more closely, or it mzy be constructed without explicit introduction of the wave functions two / representing the/Kekule structures. It might be possible to develop a simple way of discussing the structure of the amide group, for example, that would have permitted chemists to predict its cropertfes, such as planarity; but in fact no simple way of discussing me other than the way given above, involving resonance between two alternative valence-bond structures, has been discovered, and it seems likely that the discussion of complex molecules in terms of resonance among two or more valence-bond structures will continue k in the future to be useful to chemists, as it has been during the past twenty years. 13 The convenience and usefulness of the concept of resonance in the discussion of chemical problems are so great as to make the dis- advantage of the element of arbitrariness of little significance. Also, occurs it must not be forgotten that the element of arbitrariness &epe in essentially the same way in the simple structure theory of organic chemistry as in the theory of resonance - t?-ere is the same use of resonence idealized, hypothetical structural elements. In the/discussion of the benzene molecule the two Kekule/ structures have to be described as hypothetical: it is not possi.ble to synthesize molecules with w / one or the other of the two dekule structures. In the same way, however, the concept of the carbon-carbon single bond w,.T _, is an idealization. The benzene molecule has its own structure, which cannot be exactly composed of structural elements from other molecules. The propane molecule has its own structure, which cannot be ee~ek~eke A composed of structural elements from other molecules- &he-ee&em ke&e it is not possible to isolate a portion of the propane molecule, involving parts of two carbon atoms and perhaps two electrons in between * them, and say that this na+t o f the propane molecule is the carbon-carbon single bond, identical with a @ of the ethane molecule. The description \ of the propane molecule,/ as involving carbon-carbon single bonds and carbon-hydrogen single bonds is arbitrary; the concepts themselves are ideqlixations, 0 in the same way as the concepts of the Kekule structures that are described as contributing to the normal state of the benzene /molecule. For nearly a century organic chemists have found that thm idealizations of carbon-carbon single bonds, double bonds, etc. are useful, and thet the simple structural theory of organic chemistry is valuable despite its arbitrariness. The resonmce __ -.. _ ---- 5 -@hemists have found that the simple structurea theory of organic chemistry and also the resonance theor>; are valuable, despite their use of idealizations and their arbitrary character. Other extensions of the theory of the chemical bond made in recent years involve the concept of fractional bonds. Twenty-five years a simple theory of ago it was discovered that/complex crystals with largely ionic struc- tures, such as the silicate minerals, can be developed on the basis of the assumption that each cation or metal atom divides its charge or valence equally among the anions that are coordinated about it. 6 For example, in a crystal of topaz, Al,SiO,F,, -0 c ,`-' 6 Linus Pauling, The Princinles Determining - the Structure of Complex Ionic - Crystals, J. Am. Chem. SOC., $l, 1010-1026 (1929). -15. silicon each m atom is surrounded by a tetrahedron of four oxygen atoms, oxygen Lnd each aluminum atom is surrounded by an octshedr P n of four e&&s~ atoms and two fluorine atoms. The valence of silicon, 4, is assumed to be divided among four bonds, which then have the bond number 1 - they are eight- single bonds. The valence of aluminxn, 3, is divided among six bonds, each of which is a half bond. A stable structure results when the atoms are arranged in such a way that each anion, oxygen or fluorine, forms bonds equal to its valence. In topaz each oxygen atom forms one single bond with silicon and two half bonds with aluminum, / whereas each fluorine ato forms only two half bonds with aluminum. The distribution of the valences hence then corresponds to the bivalence of oxygen and the univalence of fluorine. / It was pointed out by W. L. /' Bragg that if the metal atoms are idealized as cations &S$-WH~~- (Si+* and & Al+ ) and the oxygen and fluorine atoms as anions (o-- and F-), this distribution corresponds to having %%RJE the shortest possible lines of force between the cations and the anions - the lines -16- of force need to reach only from a cation to an immediately adjacent anion, which forms part of its coordination polyhedmn. Occasionally ionic crystals are found in which there are small deviations from this requirement, but &e +s m&-4 yr-4 e . mche deviations mt larger than one quarter of a valence unit. Another application of the concept of fractional valence bonds has been made in the field of metals and alloys. In the usual quantum mechanical dis&ussion of metals, initiated by k`. Pauli (2. Physik, &l, 81 (192'7)) and So mmerfeld (Naturuiss., k,z, 825 (1927)), the assumption was made ti that only a small number of electrons contribute significantly to the binding together of the metal atoms. For example, it was cus- significantly tomary to assume that only one electron, occupying a 4~ orbital, is / the involved in the copper-copper bonds in/metaBiz copper. Sixteen years an ago/analysis of the magnetic properties of the transition metals was made that indicated that the number of bonding electrons in the transition metals is much larger, $ of the order of magnitude of six. -r . Linus Pauling, m Nature of the Interatomic Forces & -- Metals, Phys. Rev., 2, 899-904 (1938). Iron, for example, can be described as having six e&-&e valence electrons, which occupy hybrid d3sp2 orbitals. I Thq six bonds, corresponding to kheee these six valence electrons, resonate among the fourteen positions connecting an iron atom with its fourteen nearest neighbors. The bonds to the eight nearest neighbors have bond number approximately 5/8, and 4 those to the six slightly more distant neighbors have bon number l/6. In gamma iron, where each atom is surrounded by twelve enually distant neighbors, the bonds are half bonds. The concept that the ti,, ke&e structure of metals and intermetallic com;?ounds can be described in terms of valence bonds that resonate among alternative positions, aide6 by an extra orbital on most or all of the atoms (the metallic orbital) has been found of value in the discussion of the properties of these *. , sub&aces. 8- @&lT-eewepi;-wee-&w&q& resonating-bond theory of ---m-m Linus Pauling, A Resonating-tind Theory of LieiLLs & Communds, Tracy Roy. Sot. London, 4 a, 343-362 (1949). -w---m metals is supported especially strongly by the consideration of inter- atomic distances in metals and intermetallic compounds. --m-w-- Linus ?auling, Atomic Radii and Interatomic Distances & Metals, J. . Chem. SC., @2, 542 Gm - -18. The iron atom has eight electrons outside of the argon shell of eighteen. Six of these electrons are assumed, in the resonating- valence-bond theory, to be valence electrons, and the remaining two are atomic electrons, occupying tkxz 3d_ orbitals, and contributing two Bohr magnetons to the megnetic moment of the atom. A theory of the ferro- &!I magnetism of iron has recently been developed, in which, as suggested by LenerqxUmx (whys. icev., 8& ,440 (1951)), the interaction producing the Weiss field in the ferromagnetic metal is an interaction of the i i / I spin moments of the atomic electrons and uncouplcdj.spins of some of the valence electrons. It has been found possible to use spectroscopic energy values to predict the s~~+&-e&-~e&kkg number of uncoupled valence electrons, and hence the pcl~ saturation magnetic moment for iron: 0.26 the calculation leads to WO uncoupled valence electrons per atom, and saturation magnetic moment 2.26 Bohr ma?netons, which might be sub- ject to correction by two or three percent because of the contribution of orbital moment. The experimental value is 2.22. A calculated value of the Curie temperature in rough agreement with experiment is also ob- tained. w---m 10, Linus Pauling, 4 Theoq of Ferromagneti_sm_, Proc. Net. Acad. S&o U.S., 2, 551-560 (1953). a---- -19. The valence theory of metals and intermetallic compounds is still in a m unsatisfactory state. It is not yet possible to make pre- A dictions about the composition and properties of intermetallic compounds with even a small fraction of the assurance with which they can be made about organic comyunds and ordinary inorganic compounds. We may, however, hope that~ignif'icant progress m 2,' in the attack on this prob- lem during the next few years* Let us now return to the subject of the structural chemistry of organic substances, especially the complex substances that occur in living organisms, such as prot4ins. Ftecent work in this field 1e has ---em- It Linus 2auling and Robert B. Corey, Do Hydrogen-Bonded Suiral Confipuration& of the Polypeptide Chain, J. Am. Chem. Soc.,x!j& 5349 (19:d3~ ; Linus Pauling, &obert 13. Corey, and H. R. Branson,l, -g, c&e Structtire Z" P%teins, 3 Hydromn-Bonded Helical Configurations of the ---AM ?6lfie$flde Chain, Proc. Nat. Acad. Sci. U.S., 2, 205-511 (1951); Linus Pauling and R. B. Corey, Stable Confi,qura&&E of Polypeptide Chaing, Proc. Roy. SOC. c- -- the use of structural arguments that u go beyond those of the classical The ti interatomic distances and of proteins are precisely known, structure theory of organic chemistry. bond angles in the polypeptide chains i th!!$stances to within about i 0.02 ii :' / ' i `! and the bond angles to within about 2O. It is known thz,t the amide groups ,': London, 2 && 21-33 (1953). ,I +#---- *c 20 must retain their planarity; the atoms are expected not to deviate from Op' the planar configuration by more than perhaps 0.05 A. There is rotational freedom about the single bonds connecting the-e&e3 alpha carbon atom are with the adjacent amide carbon and nitrogen atoms, but there &e-e restrictions on the configurations of the polypeptide chain that can be achieved by rotations about these bonds: atoms of different parts of the chain must not approach one another so closely as to introduce large steric re- h-l+ pulsion, and in general the Y L and 0 atoms of different amide groups must be so located relatige to one another as to permit the formation of hyc!rogen jb&e+f ,)I fJ bonds, with T distance equal to 2.79 2 0.10 1. and with the oxygen rr-/-i atom not far from the MYI axis. A These requirements are stringent ones. Their application to a proposed hydrogen-bonded structure of a polypeptide chain cannot in general be made ~&km&x by x the sin,ple method of drawing a structural formula; instead, extensive numerical calculations must be carried out, or a model must be constructed. For the more complex structures, such as those that are now under consideration for the poly- peptide chains of collagen and gelatin, the analytical treatment is ..- Jhc@.b b Lz 4 so complex as to resist successful w, and only the model 21 method can be used. In order that the principles of modern structural chemistry msy be applied with the power that their reliability justifies, molecular models must be / constructed with great accuracy* For example, molecular models on the scale 2.5 cm = 1 i precision better than 0.01 cm9 believe, anticipate that the chemist of the future who is interested in the structure of proteins, n-&&e nucleic acids, polysaccharides, and otherAsubstances with v .^ high molecular weight will come to rely upon a new structural chemistry, involving precise geometrical relationships k&~een-% among the atoms in the molecules and the rigorous application of new structural principles, and that 4 great progress will be made, through this technique, in the attack, by chemical methods, on the problems of biology and medicine.