Lewis structures provide a simple method of estimating molecular shapes. The geometry about any atom covalently bonded to two or more other atoms is found by counting the number of electron groups around the atom. Each unshared pair counts as one groups and each bond, weather single or multiple counts as one group. The number of electron groups around and atom is therefore equal to the sum of the number of electron pairs on the atom and the number of other atoms bonded to it. The geometry is linear if the number of electron groups is two, trigonal if its 3, and tetrahedral if its 4.
The rule is based on teh electron pair repulsion model, which postulates that because electron pairs repel each other, they will try to stay as far apart as possible. In trigonal and tetrahedral geometries, the shape will be exactly trigonal (120° bond angles), or exactly tetrahedral (109.5° bond angles) if the electron groups are all equivalent, as for example in BH3, or CH3+ (trigonal), or in CH4, or NH4+ (tetrahedral).
If the groups are not all equivalent, the angles will deviate from the ideal values. Thus in NH, (four electron groups, three in N-H bonds, one an unshared pair), the unshared pair, being attracted only by the nitrogen nucleus, will be closer to the nitrogen on the average than will the bonding pairs, which are also attracted by a hydrogen nucleus. Therefore the repulsion between the unshared pair and a bonding pair is greater than between two bonding pairs, and the bonding pairs will be pushed closer to each other. The H-N-H angle should therefore be less than 109.5°. It is found experimentally to be 107°. Similarly, in H2O (four electron groups, two unshared pairs, and two 0-H bonds), the angle is 104.5°.
Ambiguity may arise when more than one structure contributes. Then unshared pairs in one structure may become multiple bonds in another, so that the number of electron groups around a given atom is not the same in both structures. An example is methyl azide (19). The central nitrogen is clearly linear (two electron groups), but the nitrogen bonded to CH3, has three electron groups in
19a and four in 19b. In such a situation, the number of electron groups is determined from the structure with the larger number of honds. Thus the nitrogen in question in 19 is trigonal, not tetrahedral.
The Lewis structure notation is useful because it conveys the essential qualitative information about properties of chemical compounds. The main features of the chemical properties of the groups that make up organic molecules,
and so forth, are to a first approximation constant from molecule to molecule, and one can therefore tell immediately from the Lewis structure of a substance that one has never encountered before roughly what the chemical properties will be.
There is a class of structures, however, for which the properties are not those expected from the Lewis structure. A familiar example is benzene, for which the heat of hydrogenation (Equation 1.1) is less exothermic by about 37 kcal/mole than one would have expected from Lewis structure 1 on the basis of the measured
heat of hydrogenation of ethylene. The thermochemical properties of various
types of bonds are in most instances transferable with good accuracy from molecule to molecule; a discrepancy of this magnitude therefore requires a fundamental modification of the bonding model. The difficulty with model 1 for benzene is that there is another Lewis structure, 2, which is identical to 1 except for the placement of the double bonds.
Whenever there are 2 alternative structures for a single compound, and any one of the strucutre becomes an inaccurate representation for the molecular structure. The actual structure of the molecule is actually a hybrid of these 2 strucutres. It is like a “superposition” of these 2 strucutres. The superposition of two more more Lewis structures into a composite picture of the compound is called resonance.
This terminology is well established, but unfortunate, because the term resonance when applied to a pair of pictures tends to convey the idea of a changing back and forth with time. It is therefore difficult to avoid the pitfall of thinking of the benzene molecule as a structure with three conventional double bonds, of the ethylene type, jumping rapidly back and forth from one location to another. This idea is incorrect. The electrons in the molecule move in a field of force created by the six carbon and six hydrogen nuclei arranged around a regular hexagon.
Each of the six sides of the hexagon is entirely equivalent to the otehr side, which is why electrons should, even momentarily, seek out three sides and make them different from the other three, as the two alternative pictures 3 seem to imply that they do. The symmetry of the ring of nuclei (4) is called a sixfold symmetry because rotating the picture by one-sixth of a circle will give the identical picture again. This sixfold symmetry must be reflected in the electron distribution.
A less misleading picture would be the one above, in which the circle in the middle of the ring implies a distribution of the six double bond electrons of the same symmetry as the arrangement of nuclei. We shall nevertheless usually continue to use the notation 3, as it has certain advantages for thinking about reactions.
The most important features of the structures for which the resonance is needed are, first, that the molecule is more stable (of lower energy) than one would expect from looking at one of the individual strucutres. and second that the actual distribution of the electrons in the molecule is different form what one would expect from either of the resonance strucutres. Since the composite picture shows that certain electrons are free to move to a larger area, of the molecule than a single one of the structure implies, resonance is often referred to as delocalization. We shall have more to say about delocalization later in connection with molecular orbitals.
While the benezene ring is the most fimiliar example of the necessity for modifying the Lewis strucutre language by addition of resonance concept, there are many others. The carboxylic acids, for example are much stronger acids than the alcohols; this difference must be largely due to greater stability of the carboxylate ion over the alkoxide iiion. It is the possiblity of writing two equivalent Lewis structures for the carboxylate ions that alertts us to this difference.
Another example is the allylic system. The ally1 cation (8), anion (9), and
radical (10), are all more stable than their saturated counterparts. Again, there is for each an alternativestructure :
In all the examples we have considered so far, the alternative structures have been equivalent. This will not always be the case, as the following examples illustrate :
Whenever there are nonequivalent strucutres, each will contribute to the composite picture to a different extent. The structure would represent the most stable (lowest-energy) molecule, where such a molecule actually to exist, contributes the most to composite, and the others successively less as they represent the higher energy molecules.
It is because the lowest-energy structures are most important that we specifiedin the rules for writing Lewis structures that the number of bonds should be maximum and the valence-shell occupancy not less than 8 whenever possible. Structures that violate these stipulations, such as 11 and 12, represent high-energy forms and hence do not contribute significantly to the structural pictures, which
The followinp; rules are useful in using resonance notatinn:
- All nuclei must be in the same location in every structure. Structures with nuclei in different locations, for example 15 and 16, are chemically distinct substances, and interconversions between them are actual chemical changes, always designated by
.
-
Structures with fewer bonds or with greater seperation of formal charge are less stable than those with more bonds or less charge seperation. thus 11 and 12 are higer energy respectively than 12 and 14.
- When 2 structures with formal charge have the same number of bonds and approximately same charge seperation, the structure with charge on the more electronegative atom will usually be somewhat in the lower energy state, but the difference will ordinarily be small enough that both structures can be included in the composite picture.
- All four groups attached to a pair of atoms joined by a double bond in any structure must lie in the same plane. For example, structure 18b cannot contribute because the bridged ring prevents the carbon 6 and 7 from lying in the same place as carbon 3, and the hydrogen on carbon 2. The impossiblity of strucures with double bonds at bridgeheads of small bridged rings is known as Bredt’s rule. Double bonds can occur at bridgehead if the rings are sufficiently large.
Because the covalent bond is of central importance to organic chemistry, we begin with a review of bonding theory. Later, in Chapter 10, we shall return to develop certain aspects of the theory further in preparation for the discussion of pericyclic reactions.
MODELS OF CHEMICAL BONDING
Understanding and progress in natural science rest largely on models. A little reflection will make it clear that much of chemical thinking is in terms of models, and that the models useful in chemistry are of many kinds. Although we cannot see atoms, we have many excellent reasons for believing in them, and when we think about them we think in terms of models. For some purposes a very simple
model suffices. Understanding stoichiometry, for example, requires only the idea of atoms as small lumps of matter that combine with each other in definite proportions and that have definite weights. The mechanism by which the atoms are held together in compounds is not of central importance for this purpose. When thinking about stereochemistry, we are likely to use an actual physical model consisting of small balls of wood or plastic held together by springs or sticks. Now the relative weights of atoms are immaterial, and we do not bother to reproduce them in the model; instead we try to have the holes drilled carefully so that the model will show the geometrical properties of the molecules. Still other models are entirely mathematical. We think of chemical rate processes in terms of sets of differential equations, and the details of chemical bonding require still more abstract mathematical manipulations. The point to understand is that there may be many ways of building a model for a given phenomenon, none of which is complete but each of which serves its special purpose in helping us understand some
aspect of the physical reality.
The Electron Pair Bond-Lewis Structures
The familiar Lewis structure is the simplest bonding model in common use in organic chemistry. It is based on the idea that, at the simplest level, the ionic bonding force arises from the electrostatic attraction between ions of opposite charge, and the covalent bonding force arises from sharing of electron pairs between atoms. The starting point for the Lewis structure is a notation for an atom and its valence electrons. The element symbol represents the core, that is, the nucleus and all the inner-shell electrons. The core carries a number of positive charge equal to the number of valence electrons. This positive charge is called corecharge. Valence electrons are shown explicitly. For elements in the third and later rows ofthe periodic table, the d electrons in atoms of Main Groups 111, IV, V, VI, and VII are counted as part of the core. Thus :
Ions are obtained by adding or removing electrons. The charge on an ion is given by
charge = core charge – number of electrons shown exvlicidy
An ionic compound is indicated by writing the Lewis structures for the two ions. A covalent bond model is constructed by allowing atoms to share pairs of electrons. Ordinarily, a shared pair is designated by a line:
H-H
All valence electrons of all atoms in the structure must be shown explicitly. Those electrons not in shared covalent bonds are indicated as dots, for example:
If an ion contains two or more atoms covalently bonded to each other, the total charge on the ion must equal the total core charge less the total number of electrons, shared and unshared:
In order to write-correct Lewis structures, two more concepts are needed. First, consider the total number of electrons in the immediate neighborhood of each atom. This number is called the valence-shell occupancy of the atom, and to find it, all unshared electrons around the atom and all electrons in bonds leading to the atom must be counted. The valence-shell occupancy must not exceed 2 for hydrogen and must not exceed 8 for atoms of the first row of the periodic table. For elements of the second and later rows, the valence-shell occupancy may exceed 8. The structures
are acceptable.
The second idea is that of formal charge. For purposes of determining formal charge, partition all the electrons into groups as follows: Assign to each atom all of its unshared pair elec_tronsa nd half of all electrons in bonds leading it. Call the number of electrons assigned to the atom by this process its electron ownership. The formal charge of each atom is given by
formal charge = core charge – electron ownership
To illustrate formal charge, consider the hydroxide ion, OH-. The electron ownership of H is 1, its core charge is + 1, and its formal charge is therefore zero. The electron ownership of oxygen is 7, and the core charge is +6; therefore the formal charge is – 1. All nonzero formal charges must be shown explicitly in the
structure. The reader should verify the formal charges shown in the following examples
The algebraic sum of all formal charges in a structure is equal to the total charge. Formal charge is primarily useful as a bookkeeping device for electrons, but it also gives a rough guide to the charge distribution within a molecule. In writing Lewis structures, the following procedure is to be followed:
- Count the total number of valence electrons contributed by the electrically neutral atoms. If the species being considered is an ion, add one electron to the total for each negative charge; subtract one for each positive charge.
- Write the core symbols for the atoms and fill in the number of electrons determined in Step 1. The electrons should be added so as to make the valence shell occupancy of hydrogen 2 and the valence-shell occupancy of other atoms not less than 8 wherever possible.
- Valence-shell occupancy must not exceed 2 for hydrogen and 8 for a first-row atom; for a second-row atom it may be 10 or 12.
- Maximize the number of bonds, and minimize the number of unpaired erectrons, always taking care not to violate Rule 3.
- Find the formal charge on each atom.
We shall illustrate the procedure with two examples.
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Today, I am going to write on a very basic type of reaction. The reaction is famously called aldol condensation. Most of the chemists of the world would be aware of this reaction and its mechanism. Anyhow, I am getting used to this writing, and would appreciate you people to bear with my new found interest.
Aldol condensation is a very useful reaction in the field of chemistry. It is a “bond forming” reaction and thus has wide applications in the field of synthetic chemistry. Independently discovered by Wurtz and Borodin, it involves a neucleophilic addition of a ketone moiety to an aldehyde molecule. This finally results into the synthesis of a ?-hydroxy ketone. Sometimes the ?-hydroxy ketone loses a molecule of water to give ?,?-unsaturated ketone. This reaction is called aldol condensation.
Aldehydes posessing atleast 1 Hydrogen atom is partially converted to its enolate anion by bases such as hydroxide and alkoxide ions.
Because the pKa’s of the aldehyde and water are similar, the solution contains signifi-
cant quantities of both the aldehyde and its enolate. Moreover, their reactivities are com-
plementary. The aldehyde is capable of undergoing nucleophilic addition to its carbonyl
group, and the enolate is a nucleophile capable of adding to a carbonyl group. And as
shown in Figure 18.4, this is exactly what happens. The product of this step is an alkox-
ide, which abstracts a proton from the solvent (usually water or ethanol) to yield a
[3-hydroxy aidehyde. A compound of this type is known as an aldol because it contains
both an aldehyde function and a hydroxyl group (ald + ol = aldol). The reaction is
called aldol addition.
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