LEWIS FORMULAS, STRUCTURAL ISOMERISM, AND RESONANCE STRUCTURES LEARNING OBJECTIVES: To understand the uses and limitations of Lewis formulas, to introduce structural isomerIsm, and to learn the basic concept of resonance structures. CHARACTERISTICS OF LEWIS FORMULAS: Lewis formulas are structures that show the connectivity, or bonding sequence of the atoms, indicating single, double, or triple bonds. They should also show any formal charges and unshared electrons that might be present in the molecule. Additional examples of Lewis formulas follow. H H O H O H H H O H C C C H H C C C Cl H C C C Cl H Cl H H H H C3H5ClO C3H5ClO C3H5ClO These examples were deliberately chosen because all three molecules shown have the same molecular formula, but different connectivities, or bonding sequences. Such substances are called structural isomers, or sometimes constitutional isomers. Notice that only the first structure shows the unshared electrons of chlorine. In Lewis formulas of organic compounds, it is customary to omit the lone electron pairs on the halogens unless there is a reason to show them explicitly. Lewis formulas are mostly used for covalent substances, but occasionally they also show ionic bonds that might be present in certain compounds. H The bond between nitrogen and H N H Cl chlorine is ionic. All others are covalent. H H O The bond between oxygen and H C C O Na sodium is ionic. All others are covalent. H COMMON BONDING PATTERNS FOR FIRST AND SECOND ROW ELEMENTS: Once we write enough Lewis formulas containing the elements of interest in organic chemistry, which are mostly the second row elements, we find that certain bonding patterns occur over and over. Learning these patterns is useful when trying to write Lewis formulas because they provide a convenient starting point. For example, in several of the structures given in the previous section, we find that the carbon bonded to three hydrogens is a unit that occurs quite frequently. It is called the methyl group, represented by CH3. It is so common that it is valid to write it as such in Lewis formulas, even though it is in fact an abbreviated form, because everybody knows what it stands for. H O O It is equally valid to represent the acetate H C C O or CH3 C O ion by either of these formulas. H Other common bonding patterns are shown below. HYDROGEN: Usually forms only one bond. H O H F H H H C H Cl CARBON: Forms four bonds when neutral, but it can also have only three bonds by bearing a positive or a negative charge. When it bears a negative charge it should also carry a pair of unshared electrons. H H H H C H C C H C C H Neutral carbon Cl H H CH3 A carbocation has a central carbon with C an incomplete octet and a formal +1 charge. H3C CH3 CH3 A carbanion has a central carbon with an C unshared electron pair and a formal -1 charge. H3C CH3 NITROGEN: Forms three bonds and carries a lone pair of electrons when neutral. It can also form four bonds by bearing a positive charge, in which case it carries no unshared electrons. Finally, it can also form two bonds as it carries two unshared electron pairs and a negative charge. H H3C N C N H C N Neutral nitrogen H H H3C H H H3C H H N H C N Positively charged nitrogen H H3C H H N H Negatively charged nitrogen OXYGEN: Forms two bonds and carries two lone pairs when neutral. It can form three bonds with a positive charge, or one bond with a negative charge. In each case it must carry the appropriate number of unshared electron pairs to complete the octet. H O O O H H H H H water hydronium ion hydroxide ion HALOGENS: Form one bond and carry three electron pairs when neutral. Can carry a negative charge with no bonds. They are rarely seen with positive charges. H F Cl THIRD ROW ELEMENTS: They behave like their second-row counterparts, except that they can expand their valence shells if needed. O Electron pairs on oxygen S H O S O H H H are not shown for clarity. O Cl Cl Cl Br P Br P Cl Cl Br ELECTRON DEFICIENCY IN SECOND ROW ELEMENTS: One thing worth noting is that, in the second row, only boron and carbon can form relatively stable species in which they bond with an incomplete octet. Examples have already been discussed. Boron has no choice but to be electron deficient. Carbon can bond with a complete octet or with an incomplete octet. Obviously bonding with a complete octet provides higher stability. CH F 3 C B H C CH F F 3 3 Boron has no choice but An electron deficient to have an incomplete octet carbon in a carbocation It is however very rare to observe species where nitrogen or oxygen bond with incomplete octets. Their high electronegativity renders such situation high energy, and therefore very unstable. For all intents and purposes, avoid writing formulas where oxygen or nitrogen are shown with incomplete octets, even if they carry a positive charge. CH3 CH3 O CH3 O CH3 To write this structure without the lone pair This structure is unacceptable of electrons on oxygen is unacceptable. and indeed it looks quite awkward H This species might exitst in the high energy environment N of a mass spectrometer, but it is not frequently observed H H in common organic reactions. RESONANCE STRUCTURES AND SOME LIMITATIONS OF LEWIS FORMULAS : Lewis formulas are misleading in the sense that atoms and electrons are shown as being static. By being essentially two-dimensional representations they also fail to give an accurate idea of the three-dimensional features of the molecule, such as actual bond angles and topography of the molecular frame. Furthermore, a given compound can have several valid Lewis formulas. For example CH3CNO can be represented by at least three different but valid Lewis structures called resonance forms, or resonance structures, shown below. H H H H C C N O H C C N O H C C N O H H H I II III However, a stable compound such as the above does not exist in multiple states represented by structures I, or II, or III. The compound exists in a single state called a hybrid of all three structures. That is, it contains contributions of all three resonance forms, much like a person might have physical features inherited from each parent to varying degrees. In the resonance forms shown above the atoms remain in one place. The basic bonding pattern, or connectivity, is the same in all structures, but some electrons have changed locations. This means that there are certain rules for electron mobility that enable us to “push” electrons around to arrive from one resonance structure to another. These rules will be examined in detail in a later paper. ALL RESONANCE STRUCTURES MUST BE VALID LEWIS FORMULAS: By convention, we use double-headed arrows to indicate that several resonance structures contribute to the same hybrid. Continuing with the example we’ve been using, the resonance structures for CH3CNO should be written in this way if we want to emphasize that they represent the same hybrid. C N O CH3 C N O CH3 CH3 C N O I II III Do not confuse double-headed arrows with double arrows. A double arrow indicates that two or more species are in equilibrium with each other and therefore have a separate existence. Double-headed arrows indicate resonance structures that do not exist by themselves. They simply represent features that the actual molecule, the hybrid, possesses to one extent or another. When writing resonance structures keep in mind that THEY ALL MUST BE VALID LEWIS FORMULAS. The factors that make up valid Lewis formulas are as follows. 1. Observe the rules of covalent bonding, including common patterns as discussed previously. Make sure to show all single, double, and triple bonds. 2. Account for the total number of valence electrons being shared (from all the elements), including bonding and nonbonding electrons. Make sure to show these nonbonding electrons. 3. Account for the net charge of the molecule or species, showing formal charges where they belong. 4.Observe the octet rule as much as possible, but also understand that there are instances where some atoms may not fulfill this rule. 5. Avoid having unpaired electrons (single electrons with no partners) unless the total number of valence electrons for all elements is an odd number. This is not a very frequent occurrence, but the following example shows a species that could exist as a reaction intermediate in some high energy environments. H C H H The total number of valence electrons being shared for all atoms is 4 from carbon and 3 from the three hydrogens, for a total of 7. Because it is an odd number, it is impossible to have all these electrons paired. Therefore the presence of a single electron cannot be avoided. Notice that there is no formal charge on carbon, since it has no surplus or deficit of valence electrons. RELATIVE ENERGIES OF RESONANCE STRUCTURES. From the examples given so far it can be seen that some resonance forms are structurally equivalent and others are not. The potential energy associated with equivalent Lewis structures is the same. If the Lewis structures are not equivalent, then the potential energy associated with them is most likely different.