Ruth A. Walker
Hunter College City University of New York New York
General Chemistry Exercise Using Atomic and Molecular Orbital Models
Increased emphasis in general chemistry on atomic and molecular orbital theory has produced a need for a laboratory exercise which will familiarize the students with the three-dimensional aspects of these concepts and with their role in determining the fundamental spatial relationships in molecular bonding. I\luch recent work has involved the development of space-filling and projection models as a means of portraying both atomic and molecular orbitals (1-11). However, very few specific suggestions have been made as to the exact use of such models in a laboratory exercise for general chemistry (12-15), partially because the preparation of the number of units needed presents a major hurdle. Martins (16) has described an exercise during which students use solid forms representing s and p orbitals to make compounds such as HzO, BF,, BH,, etc. His solution to the problem is quite different from the one proposed here which includes a study of the ground state of each element before consideration of the valence state and molecular geometry. This exercise was written for a particular general chemistry laboratory manual (17) and utilizes models in which the orbitals are merely outlined with wire rather than being represented by solid forms. These models are simple and inexpensive to make. If necessary, the members of the first laboratory section can be directed to prepare the kit for the use of subsequent sections. The fact that the models are not space-filling means that the students have to vizualize the solid forms from the outlines. This does not in any way interfere with the major objective which is to convey an understanding of the threedimensionality of the electron distribution in the ground state atom and of the effect of bonding on this distribution as manifested in hybridization, bond length, bond angle, and the relationship betweeu the sigma and the pi bond.
by a white halo around the central sphere (Note: the exaggerated size of the nucleus as compared to the orbitals must be pointed out at the start). The p electron orbitals are yellow, red, or green depending on their direction in space. An orbital which contains only one electron is represented by a solid color. The presence of a pair of electrons in a given orbital is shown by a blue stripe (colored wire). The position for inserting the pipe cleaners is determined from a cardboard template which fits around the circumference of the sphere and has points on it that are 90' apart. The experimental directions guide the student through a careful consideration of each structure, pointing out such relationships as Hund's rule of maximum multiplicity. Particular emphasis is placed on the fact that the pipe cleaners represent the surfaces of solids rather than mere halos. The student is constantly referred to the solid shapes at the bottom of the exhibit and is made to realize that these stable arrangements of spheres are what determine the relative positions of the various orbitals. He is directed to make a number of drawings designed to help him relate the models to the diagrams he will find in books. He also prepares a chart giving the spectral notation and the Lewis structures for the first 18 elements. Sigma Bond
The characteristics of the sigma bond are studied with the models shown in Figure 2. These are all in the kit given to each pair of students. They represent the valence state of hydrogen, chlorine, and the Period 2 elements from beryllium through fluorine. Blue pipe cleaners represent sp, sp2, and sp3 hybridization, with shades of blue indicating the extent of p character. The tetrahedral positions for the orbitals are located
Atomic Orbitals
The first part of the exercise deals with the ground state con6gurations of the Period 2 elements. The exhibit panel shown in Figure 1 is placed at the end of each laboratory desk. The construction materials are: 1-in. Styrofoam spheres, colored pipe cleaners, and electric train wire.' The s electrons are represented Presented before the 4th Annual Metropolitan Regional Meeting of the New York-New Jersey Sections of the American Chemical Society, February, 1965. ' Whit,e, red, yellow, and green pipe cleaners and the Styrofoam spheres can be purchased from Star Band Co., Broad and Commerce Streets, Portsmouth, Virginia. Other colored pipe cleaners can be purchased from Rochester-Livingstone, Inc., 81 Charlotte Street, Rochester 7, New York. Tmin wire is available from Apex, Brooklyn, New York, cat. no. 24-100.
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Figure 1. Ground state exhibit panel. per orbital.
Stripe indicates two electrons
Valence rtote of Period 2 elements. Striped bond on CI Figure 2. indicates inner completed octet. N, 0, and F ore 011 hybridized.
from a cardboard temnlate according to directions bv Sanderson (18) using 'a circle withYa diameter 0.f3 the diameter of the sphere. The tetrahedral hybridized shape is employed for the two halogens as well as for the nitrogen and oxygen, but the orbitals of all four are the colors used for the sand p orbitals in the ground state rather than being blue. The orbitals in the chlorine atoms are made from longer pipe cleaners than the others to show the increased atomic size, and a red striped band is placed around the sphere to indicate the completed inner octet. The hydrogen sphere is off-center to show the increased electron density which results between nuclei after bond formation. There are four chlorines and two hydrogens in each kit. These "atoms" can be readily combined into "molecules" by overlapping the appropriate orbitals and clamping them together with a spring paper clip. Bond length and the dependence of the degree of overlap on the nature of the electrons are illustrated by hydrogen, hydrogen chloride, and chlorine as shown in Figure 3. Increased bond length due to increased size can be seen from the models of H F and HC1 (Fig. 3). Bond angles and the role of hybridization are demonstrated by having each student prepare successively the halidesTand/or the hydrides of the elements from beryllium through fluorine. The bond angles become obvious as soon as the atoms are clam~edtoaether and in some instances can be traced and measured (see H 2 0in Fig. 3). The molecules so formed are amazingly stable and sturdy. Also obvious is the fact that BC13 is a planar molecule while that of NC13 is pyramidal (Fig. 4). Again the relationship between the model and the printed page needs to be established, so the student is asked to draw various pictures and to p r e pare a chart containing both the structural line formulas and the electron dot formulas for the halides of Period 2 and Period 3 from families 2-7.
Figure 4.
Pyrornidd structure of NCll compared to the planarity of BCla
Pi Bond
The atoms needed to demonstrate pi bonding are constructed somewhat differently from those above (see Fig. 5 ) . The p orbitals are perpendicular to each other and extend on both sides of the sphere just as in the ground state, except that the rear lobe of the orbital to be used in "bonding" has beeu made very small to indicate the shift that occurs in electron density on bonding. The s electrons are represented by a striped black and white band around the sphere to show the inner spherical orbital without having it interfere with the geometry of the pi bond. Each kit contains one of these oxygens and two nitrogens. When such an oxygen and such a nitrogen are clamped together with the yellow p orbitals overlapping, the red orbital of a single p electron in oxygen automatically parallels a similar p orbital in nitrogen. Additional red pipe cleaners are used to join these two orbitals across the space between the atoms above and below the sigma bond. They represent the interaction between two such parallel orbitals and they serve to emphasize the relationship between a sigma and a pi bond (Fig. 5 ) . Similarly, two nitrogen atoms can be joined with two extra red and green pieces, thus indicating two pi clouds which become like a cylinder of electron density around the sigma bond. ~ i t ~ ~ t , cited ,,~ (1) CAMPBELL, J. A., J. CHEM.EDUC.,25,200 (1948). (2) NOELLER, C. R,, J. CHEM. EDUC,,26, 429 949,. (3) LAMBERT, FRANK L., J. CEIEM.EDUC.,30, 503 (1953); 34, 217 (1957). (4) KAUFFMAN, G. B., J. CHEM.Eouc., 36, 82 (1959).
~~~~4~&,",","O~bJ;~~,"i~~,";~,8i~,"$~,"~$1)~ BENT, HE^^^, J. CAEM. E ~ u c . 40, , 446, 523 (1963). (8) FIESER, LOUISF., J. CHEM.EDUC.,40,457 (1963).
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Figure 3. Variation in bond length, bond mgle, ond in the degree of orbital overlop as determined by the notvre of the electrons involved.
Figure 5.
Unhybridired atornr "red to show the pi cloud in NO and N1.
Volume 42, Number 12, December 1965
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OGRYZLO, E. A., AND PORTER,GERALD B., J . CREM.EDUC., 40, 256 (1963). G. OLAF,J. CHEM.EDUC.,41, 219 (1964). LARSON, J., AND BAUMBRUMLIK,GEORGEC., BARRETT, EDWARD GARTEN, REUBENL., J . CHEM.EDUC.,41. 221 (1964). B., J. CHEW.EDUC.,36,595 (1959). PIERCE,JAMES HART,HAROLD, AND SCAUETZ, ROBERT D., "A Laboratory Manual for a Short Course in Organic Chemistry," Houghton Mifflin Co., Boston, hlilss., 1961, pp. 11-18.
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SANDERSON, R. T., "Teaching Chemistry with Models," D. van Nostrand Co., Princeton, New Jersey, 1962, pp. 89-94. IRWIN, J. CUEM. STONE,A. HARRIS,AND SIEGELMAN, E ~ u c .41, , 395 (1964). J. CHEM.EDUC.,41,658 (1964). MARTINS,GEORGE, JOHNSTON, H., "Experimental General Chemistry," Wm. C. Brown, Dubuque, Iowa.52003, 1965, pp. 64-71. SANDERSON, R. T., op. cit., p. 104.
