1A: NutraSweet: Structural Bonding
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We can learn most of the important principles of organic chemistry through a study of the aspartame molecule. In doing so, we will see the scientific problems that chemists faced as they attempted to synthesize and evaluate this artificial sweetener.
Structure and Bonding
Typically, chemists represent organic structure in an abbreviated, symbolic fashion, using lines and angles to represent bonds and carbon atoms, respectively. Hetero atoms, often defined as not carbon (C) or hydrogen (H), are written with their usual symbols (O, S, N, etc.) and their bond connections are specifically shown, but hydrogens attached to carbon to complete carbon’s valence of four are not usually written. Unshared electron pairs are frequently omitted. See how these conventions can oerate in the depiction of aspartame.
Chemists depend upon the symbolic nature of our language. Chemistry is the study of matter; we look for microscopic understanding at the atomic and molecular level to explain phenomena we see at the macroscopic level. And we seek some symbolic ways to express both those ideas. This need for symbolic representation of the microscopic and macroscopic world is the fundamental source of the complex representation of molecules in organic chemistry.
We present the structure of Nutrasweet in three different ways in Figure 2.
In structure A, all the atoms are shown although we do not show unbonded electron pairs. Though instructive for a novice chemist, such a depiction is cluttered and time consuming to draw.
Typically chemists will represent organic molecules in an abbreviated manner with lines and angles representing bonds and carbon atoms respectively. Hetero atoms, often described as "not carbon, C, or hydrogen, H" are writtenwith their usual symbols (S, N, O, etc.) and their bonded connections are specifically shown. We generally do not show the bonded connections of hydrogen to carbon - the number of hydrogens is implied by assuming any bonds we do not see connected to other atoms, are hydrogens. Structure B uses these conventions to depict the aspartame molecule.
Structure C uses wedges () and dashes () in an attempt to bring some three-dimensionality to the drawing. Wedges denote bonds that project out from the paper toward the viewer while dashes represent bonds projecting away from the viewer. Typically, only some of the bonds in a molecule are shown in this manner. Some of the remaining atoms will be in the plane of the two-dimensional representation while others will be forward or away; they just aren’t being represented as such. Using one of the model kits frequently available to chemistry students to build a model of a compound helps in visualizing the three-dimensional nature of a structure. Computer software programs also can be used to visualize in three dimensions. The actual spacial arrangement of atoms is extremely important in biochemical/physiological functioning, as will be elaborated upon later.
Organic chemistry is frequently defined as the chemistry of carbon compounds, and organic compounds are defined as containing carbon. The importance of organic compounds to us cannot be overstated, for they are the working components of the natural world. The chemistry and biochemistry of organic molecules are the means by which all living systems function and reproduce. The more we understand about the physical and chemical characteristics of such systems, the better we understand the natural world. This raises many philosophical and moral questions, but this ChemCase will focus on somewhat more mundane questions and answers.
Carbon has an atomic number of six, with an electronic structure of 1s2 2s2 2p2. All of carbon’s compounds show it to have a valence of four (i.e., it uses four electrons to form four bonds). It can be readily determined that in methane (CH4) all four C-H bonds are identical. And it is easily determined that carbon can use its four bonds in four different arrays: four single bonds; two single bonds and one double bond; one single and one triple bond; and, rarely, two double bonds. Recall that a bond is a shared electron pair (two electrons). A double bond, then, would be two shared pairs (four electrons total); a triple bond, three electron pairs. See Figure 3.
Note the bonding types used by carbon in NutraSweet.
It is not immediately obvious how carbon’s bonding behavior is explained based on these explanations. For instance, how can all of the C-H bonds in methane be identical when the atomic structure of carbon has no more than two electrons of identical energy?
The explanation used by chemists to account for all of carbon’s bonding is based on the use of hybridized atomic orbitals. When atoms react to form molecules, the process occurs to form a low(est) energy arrangement. That may be a covalent bond, a polar covalent bond, or an ionic bond. As chemists, we attempt to develop theories based on experimental observations that can guide us in understanding what happens. To account for carbon’s use of four identical atomic orbitals in methane’s bonding, we suggest that the 2s and 2p atomic orbitals are spacially scrambled to produce four new orbitals, four new regions in space where there is the highest probability of finding electrons. It makes sense energetically to put these electrons as far away from each other as possible (since electrons repel each other). What is the shape that puts four orbitals, four of anything emanating from a single point (nucleus), as far from each other as possible? A tetrahedron, with mutually dihedral angles of 109o. This fits nicely with the VSEPR model of molecular structure.
Since these hybridized orbitals are not s or p, what do we call them? It depends. In methane, they are called sp3 and there are four such equivalent orbitals. (sp3 Designation arises from the mental construct that one of the paired 2s orbital electrons can be promoted into the empty 2p orbital to give four unpaired electrons: 2s1 2px1 2py1 2pz1. These four orbitals, one s and three p, can hybridize into the new orbitals, the new regions in space, which are called, then, sp3). A consequence of this hybridization is that, when carbon forms four single bonds, the shape around the carbon will always be tetrahedral, though the bond angles may deviate very slightly from 109.5o if the four attached atoms are not identical. As suggested earlier, molecular shape is related directly to biological activity, so understanding of this type of structural detail is significant.
Carbon’s bonding in ethene is explained by suggesting that when carbon forms one double and two single bonds, it hybridizes the remaining 2s orbital and two of the three 2p orbitals to form three sp2 orbitals. These orbitals are directed to the corners of an equilateral triangle (120o), again adhering to VSEPR concepts. The remaining 2p orbital is perpendicular to the three planar sp2 orbitals. In acetylene, carbon forms two sp hybridized orbitals (180o apart) leaving two 2p orbitals mutually perpendicular to each other and the sp orbitals (see Atkins and Jones, Chapter 9). (The central atom in allene is sp hybridized.) The use of these hybridizations extends to all other compounds containing these four types of bonding, so the shapes around these various bonding units can be determined in every organic molecule, though overall shape of the molecule reflects combinations of factors discussed later.
Let’s verbalize the construction of methane and ethene, though this is not how these compounds are synthesized. A C-H bond in methane forms by the overlap of an s atomic orbital from hydrogen with one of the four sp3 atomic orbitals from carbon to form a molecular orbital, a bond, between the two atoms. Four such overlaps occur and the result is the CH4 molecule, with identical C-H bonds (length and strength) and a tetrahedral shape. For ethene, each carbon overlaps two of its sp2 orbitals with an s orbital of hydrogen to form the two C-H bonds at angles of 120o. The carbons bond to each other by overlapping end on end their remaining sp2 orbitals to form one of the two bonds. The second bond forms by less effective overlap of the two remaining, parallel p orbitals. This overlap occurs in two regions of space since the p orbital has two lobes, though only one electron is associated with each orbital. The net result is a molecule in which the six atoms are in the same plane. Single bonds produced by end-on-end overlaps of atomic orbitals are called sigma (s) bonds, while the parallel overlap of p orbitals leads to bonds called pi (p). Pi bonds are weaker than sigma, more easily broken, and the site of many important chemical reactions, though none of interest to us in this study.
Refer to the structure of aspartame (NutraSweet). Using the information just presented, one can predict the proper hybridization of each carbon and, from that hybridization, assign three-dimensional shape to each carbon and its immediate attachments. Hydrogen, with only one electron (1s), has no need to hybridize when it bonds. It is believed that when nitrogen forms three single bonds, each of these hetero atoms (O and N) is sp3 hybridized. When oxygen and nitrogen form multiple bonds, sp2 or sp hybridization is utilized.
The carbons that have been starred are all sp3 hybridized. Each one has four single bonds attached. The shape of the carbon and its four-bonded atoms would be tetrahedral. All of the remaining carbons, including the six in the ring, are sp2 hybridized. Those carbons and their attached three atoms would be in the same plane.