Teaching Chemistry with Electron Density Models

Gwendolyn P. Shusterman, Alan J. Shusterman

Journal of Chemical Education, 74(7), 771-775 (1997)

The entire text of this article, complete with "small" figures can be found at the Journal of Chemical Education Web site.

Larger versions of the figures, along with a brief explanation of each figure's content can be found below (click on thumbnail for full-size figure).

Figure List

Figure 1

(30K) The figure shows how electron density varies in Li and Li⁺. The colored maps (A & D) show how electron density varies in a plane containing the atomic nucleus. The density is coded by color, increasing from extremely low values (RED < 0.0001 a.u.) to the highest values calculated (ORANGE -> YELLOW -> GREEN -> BLUE). If you think of the atom's electron "cloud" as being defined by the size and shape of the latter region (ORANGE-BLUE), then it is obvious that Li has a larger electron cloud (A) then Li⁺ (D).

The mesh and solid surfaces show the same electron density data in a different way. The solid surfaces identify points where the electron density is relatively high (0.02 a.u.). These surfaces are virtually identical in Li and Li⁺, suggesting that removing an electron has a relatively small effect on the size of the "core" of the electron "cloud". The mesh surfaces identify points where the electron density is ten times lowere (0.002 a.u.). The change in size of these surfaces shows that removing an electron from Li shrinks the atom by removing electron density from the outer part of the atom. (back to top)

Figure 2

(34K) The figure uses electron density models to compare the shapes of methane, CH₄ and ammonia, NH₃. The mesh surfaces identify points where the electron density is relatively low (0.002 a.u.). These points more or less define the "edge" of the electron "cloud" in each molecule. You can see that even though the two molecules contain different numbers of atoms, they still have similar shapes. The bulge above the nitrogen atom in ammonia can be attributed to this atom's nonbonding electrons. (back to top)

Figure 3

(76K) The figure uses electron density models to compare the shapes of cyclohexane and 18-crown-6. The mesh surfaces identify points where the electron density is relatively low (0.002 a.u.). The most interesting feature of these two models is the presence or absence of a "cavity" in the center of each. The carbon atoms in small, and even medium-sized, rings are so close together that their electron clouds overlap. A truly large ring like 18-crown-6, on the other hand, keeps its ring atoms so far apart that a large region of "no" electron density appears in the middle of the ring. This "cavity" can be occupied by other atoms and ions that are not actually part of the ring molecule. (back to top)

Figure 4

(36K) The figure uses three types of electron density models to compare the bonding and polarity of simple hydrides: LiH, H₂, and HF. The mesh surfaces identify points where the electron density is relatively low (0.002 a.u.). These points more or less define the "edge" of the electron "cloud" in each molecule. Notice how the size of the electron cloud near H shrinks as its bonding partner changes from Li -> H -> F. Recalling the analysis of Li and Li⁺ given in Figure 1, you can conclude that the amount of electron density belonging to H is greatest in LiH and least in HF. In other words, Li and F do not share bonding electrons equally with H (equal sharing must occur in H₂). Li donates electron density to H, but F "steals" electron density from H.

Chemists describe the ability of an atom to "steal" bonding electrons from its partner as its electronegativity. These figures demonstrate that electronegativity increases in the order: Li -> H -> F.

The colored maps show how each molecule's electrostatic potential varies on the 0.002 isodensity surfaces. The variation in potential is shown by color - RED (lowest) -> ORANGE -> YELLOW -> GREEN -> BLUE (highest) - and indicates whether a particular region is electron-rich (RED) or electron-poor (BLUE). LiH and HF are polar molecules in that the two ends of these molecules are electron-rich and electron-poor respectively. The change in potential around H is also consistent with the previous analysis based on surface size; H is electron-rich in LiH (RED), neutral in H₂ (GREEN), and electron-poor in HF (BLUE).

Finally, the solid surfaces (inside the mesh surfaces) identify points where the electron density is relatively high (0.08 a.u.). Atoms that share electrons (covalently bonded) build up electron density in the region between the two nuclei. The models of H₂ and HF show that these molecules contain covalent bonds - the electron density between the nuclei is 0.08 a.u. or greater. The model of LiH, on the other hand, suggests that this bond is largely ionic - although there are regions of high electron density around each nucleus, electron density is less than 0.08 between the two nuclei. (back to top)

Figure 5

(28K) The figure uses three types of electron density models to analyze bonding in a hydrocarbon containing CC single, double, and triple bonds: HCCCH₂CH=CH₂ (pent-4-en-1-yne).

The mesh surface (A, top) identifies points where the electron density is relatively high (0.08 a.u.). This surface encloses all of the nuclei and all of the internuclear regions and shows that all of the bonds are covalent (see discussion in Figure 4).

The solid surfaces (A, bottom) identify points where the electron density is even higher (0.2 a.u.). Several high electron density domains can be observed: one around each atomic nucleus, one enclosing the internuclear region corresponding to the CC double bond, and one enclosing the internuclear region corresponding to the CC triple bond. The internuclear regions of the CC and CH single bonds are not included, and this shows that multiple bonds involving more electron sharing (higher electron density build up) than single bonds.

The colored map gives the molecule's electrostatic potential on the 0.002 isodensity surface. RED = lowest potential (electron-rich) and BLUE = highest potential (electron-poor). The CC double bond produces two electron-rich regions, one above the nuclear plane and one below (not visible). The CC triple bond produces a single electron-rich region shaped like a ring around the CC bond axis. The map also shows that the H attached to the alkyne carbon is much more electron-poor than the other H atoms. These features are all consistent with 1) standard pi bonding models, and 2) increasing C electronegativity in the order: sp³ < sp² < sp. (back to top)

Figure 6

(34K) The figure shows separate electron density maps of carbonic acid, bicarbonate, and carbonate. Each model identifies points where the electron density is extremely high (0.2 a.u.), and the amount of internuclear electron density (roughly given by the size of the surface in these regions) can be used to identify the degree of CO multiple bonding (see discussion of Figure 5). Every CO bond seems to include some degree of multiple bonding with the sole exception of the C-OH bond in the bicarbonate ion.

The models also indicate the relative degree of CO multiple bonding. For example, the two multiple CO bonds in bicarbonate are equivalent. Also, the three CO bonds in carbonate are equivalent. It is impossible to represent either of these bonding arrangements using a single Lewis structure. Chemists say that these ions are resonance hybrids of two and three Lewis structures respectively, but it is important to realize that these ions do not "switch" or "alternate" between these structures. As the models illustrate, the electron density is constant and does not vary in time. We are forced to use confusing terms like "resonance" because we have not yet found a better (or simpler) language. (back to top)

Figure 7

(36K) The figure shows separate electrostatic potential maps of carbonic acid, bicarbonate, and carbonate. The same type of color coding is used in each model - RED = low potential (ELECTRON-RICH) and BLUE = high potential (ELECTRON-POOR) - but the colors are pegged to different potentials. For example, the RED regions in the carbonic acid model have VERY different potentials (-30 kcal/mol) from the RED regions in the bicarbonate (-160 kcal/mol) and carbonate models (-270 kcal/mol). The same types of differences apply to the other colors as well.

The models clearly show that the negative charge in bicarbonate, and the double negative charge in carbonate, cannot be associated with a particular oxygen atom. That is, the electron-rich (RED) regions are spread over two and three atoms depending on the ion. Any attempt to represent the position of the "extra" electron density using a single Lewis structure is doomed to failure.

The models also show that the most electron-rich regions lie in the plane of the molecule and in between pairs of oxygen atoms. Chemists have observed that metal cations attach themselves preferentially to these regions. (back to top)

Figure 8

(31K) The figure shows two electrostatic potential maps, one of acetic acid and one of sulfuric acid. The molecules have both been positioned so that the "acidic" hydrogen is on the right-hand side of each molecule.

Recalling that electrostatic potential in these models increases RED (electron-rich) -> ORANGE -> YELLOW -> GREEN -> BLUE (electron-poor), one can see that the acidic hydrogen is the most electron-poor (BLUE) hydrogen in each molecule. Also, there is a big difference between the two molecules. The acidic hydrogen in sulfuric acid (RIGHT) is much more electron-poor (more "proton" like) than the corresponding hydrogen in acetic acid (LEFT). This helps explains why sulfuric acid is a much stronger acid than acetic acid. (back to top)

Figure 9

(56K) The figure shows two electrostatic potential maps, one of guanine and one of cytosine. The models have been positioned in the same orientation that is used by theses molecules in the DNA double helix (the rings are closer together in DNA however).

Recalling that electrostatic potential in these models increases RED (electron-rich) -> ORANGE -> YELLOW -> GREEN -> BLUE (electron-poor), one can see that the rings align sites of complementary electrostatic potential. Three electron-poor hydrogens (BLUE) are aligned with electron-rich (RED) oxygen or nitrogen atoms. These pairings are called hydrogen bonds and provides a significant fraction of the attractive force between the two strands of the double helix.

Can you find other hydrogen bonding sites around the perimeter of these molecules? Chemists think that such sites are used by proteins and drugs to control the behavior and reading of DNA. (back to top)

(last updated 6/26/97)