Structure and Bonding in Organic Compounds

Keys to the Chapter

Atomic Structure and Properties

Two periodic trends are important to understanding the physical and chemical properties of organic compounds. They are electronegativity and atomic radius.

The electronegativity scale is an index of the attraction of an atom for an electron. It increases from left to right in a period and from bottom to top in a group of the periodic table. The order of electronegativities for the three most common elements in organic molecules, excluding hydrogen, is C < N < O. Their electronegativity values differ by 0.5 between neighboring elements in this part of the second period. There is a more pronounced difference between second and third period elements. Thus, fluorine and chlorine differ by 1.0, as do oxygen and sulfur. The order of the electronegativity values of the halogens is I < Br < Cl.

Ionic and Covalent Bonds

There are two main classes of bonds. Ionic bonds predominate in inorganic compounds, but covalent bonds are much more important in organic chemistry. When positive and negative ions combine to form an ionic compound, the charges of the cations and anions must be balanced to give a neutral compound. For ionic compounds, the cation is named first and then the anion. Thus, ammonium sulfide contains (NH4)2 and S2-. Two ammonium ions are required to balance the charge of one sulfide ion, so the formula of ammonium sulfide is (NH4)2S. Parentheses enclose a polyatomic ion when a formula unit contains two or more of that ion, and the subscript is placed outside the parentheses.

A covalent bond forms when two nuclei are simultaneously attracted to the same pair of electrons. Carbon usually forms covalent bonds to other elements. The stability of Lewis structures is attributed to the octet rule that states that second row elements tend to form associations of atoms with eight electrons (both shared and unshared) in the valence shell of all atoms of the molecule.

One or more pairs of electrons can be shared between carbon atoms. Single, double, and triple bonds are linked one, two, and three pairs of electrons, respectively. In applying the octet rule, the bonding electrons are counted twice. That is, each atom "owns" the bonding electrons, so they count toward the total of eight for each atom.

With the exception of bonds to carbon and to hydrogen, carbon forms polar covalent bonds to other elements. The degree of polarity depends on the difference in the electronegativity values of the bonded atoms. The direction of the bond moment is indicated by an arrow with a cross at the end opposite the arrow head. The symbols d+ and d- indicate the partially positive and partially negative atoms of the bonded atoms.

Strategy for Writing Lewis Structures

When we write a Lewis structure, we first need to know how many electrons are in a molecule based and where they are located.

Consider vinyl chloride, C2H3Cl, which is used to produce polymers for commercial products such as PVC pipes. It contains a total of 18 electrons. Hydrogen forms only one bond in all compounds. Chlorine also forms one bond to carbon. The basic skeleton of the molecule is shown below.

The molecular skeleton accounts for eight electrons; two per single bond. Each carbon atom still needs two more electrons to complete its octet, and the chlorine atom needs six. The six electrons on chlorine form three lone pairs. Each carbon contributes one electron to the single bond. Each carbon has four electrons, and each donates one more to form a double bond.

Formal Charge

We determine formal charges in several steps.

1. Count the total number of valence electrons for each atom in the molecule.

2. Each atom "owns" its nonbonded electron pairs.

3. Electrons in bonds are shared equally between the bonded atoms; in a single bond each atom gets one electron, in a double bond it gets two, and so forth.

4. If an atom has more electrons in the bonded structure than it would have if neutral, it has a formal negative charge; if it has fewer electrons than it would have as a neutral atom, it has a formal positive charge.

A few simple rules make it easy to determine the formal charge in most cases by inspection. For example, if nitrogen has three bonds-regardless of the combination of single, double, or triple bonds-and a pair of electrons, then it has no formal charge. If there are four bonds to nitrogen-regardless of the combination of single, double, or triple bonds-the nitrogen atom has a formal +1 charge. Similarly, if oxygen has two bonds-regardless of the combination of single or double bonds-and two pairs of electrons, then it has no formal charge. If there are three bonds to oxygen-regardless of the combination of single or double bonds-the oxygen atom has a formal +1 charge. The structure shown below contains an oxygen atom with a +1 formal charge; the entire species has a net +1 charge.

Resonance Theory

For most compounds, one Lewis structure describes the distribution of electrons and the types of bonds in a molecule. However, for some species a single Lewis structure does not provide an adequate description of bonding. Resonance structures provide a bookkeeping device to describe the delocalization of electrons, giving structures that cannot be adequately described by a single Lewis structure. Such bonding is described using two or more resonance contributors that differ only in the location of the electrons. The positions of the nuclei are unchanged. The actual structure of a molecule that is pictured by resonance structures has characteristics of all the resonance contributors.

Curved arrows are used to show the movement of electrons to transform one resonance contributor into another. The electrons move from the position indicated by the tail of the arrow toward the position shown by the head.

The degree to which various resonance forms contribute to the actual structure in terms of the properties of the bonds and the location of charge is not the same for all resonance forms. The overriding first rule is that the Lewis octet must be considered as a first priority. After that, the location of charge on atoms of appropriate electronegativity can be considered.

Valence-Shell Electron-Pair Repulsion Theory

Like charges repel each other, so the electron pairs surrounding a central atom in a molecule should repel each other and move as far apart as possible. We use valence-shell electron-pair repulsion (VSEPR) theory to predict the shapes of molecules. VSEPR theory allows us to predict whether the geometry around any given atom is tetrahedral, trigonal planar, or linear.

Using VSEPR theory requires that regions of electron density be considered regardless of how many electrons are contained in the region. Thus, a single-bonded pair or two pairs of electrons in a double bond are considered as "equal." The following rules cover most cases.

1. Two regions containing electrons around a central atom are 180° apart, producing a linear arrangement.

2. Three regions containing electrons around a central atom are 120° apart, producing a trigonal planar arrangement.

3. Four regions containing electrons around a central atom are 109.5° apart, producing a tetrahedral arrangement.

The electron pairs around a central atom may be bonding electrons or nonbonding electrons, and both kinds of valence-shell electron pairs must be considered in determining the shape of a molecule. When all of the electron pairs are arranged to minimize repulsion, we look at the molecule to see how the atoms are arranged in relation to each other. The geometric arrangement of the atoms determines the bond angles.

Consider the structure of an isocyanate group in methylisocyanate.

The nitrogen atom has three regions containing electrons around it. They are a single bond, a double bond, and a nonbonded pair of electrons. So, these features will have a trigonal planar arrangement, and the R-N=C bond angle is 120°. The isocyanate carbon atom has two groups of electrons around it-two double bonds-so they will have a linear arrangement. The N=C=O bond angle is 180°.

Dipole Moments

The polarity of a molecule is given by its dipole moment. The dipole moment depends upon both the polarity of individual bonds and the arrangement of those bonds in the molecule. In some molecules, the dipole moments are pointed in opposite directions so that they cancel one another. As a result, there is no net resultant dipole moment. In other molecules, the dipole moments may reinforce each other or partially cancel, causing a net dipole moment.

Atomic and Molecular orbitals

Atomic orbitals are mathematical equations that describe the discrete, quantized energy levels of atoms. They are described as 1s, 2s, 2p, and so forth. Each atomic orbital can contain a maximum of two electrons with opposite spins. The square of the equation for an atomic orbital gives the probability of finding an electron within a given region of space.

The concepts developed for atomic orbitals can be extended to molecular orbitals that extend across a molecule. Molecular orbitals are linear combinations of atomic orbitals, which represent the distribution of...