Chapter 1

Working with Models and Molecules

IN THIS CHAPTER

Diagramming Lewis structures

Predicting bond dipoles and dipole moments of molecules

Seeing atom hybridizations and geometries

Discovering orbital diagrams

Organic chemists use models to describe molecules because atoms are tiny creatures with some very unusual behaviors, and models are a convenient way to describe on paper how the atoms in a molecule are bonded to each other, and where the electrons in an atom are located. Models are also convenient for helping you understand how reactions occur.

In this chapter, you use the Lewis structure, the most commonly used model for representing molecules in organic chemistry. You also practice applying the concept of atom hybridizations to construct orbital diagrams of molecules, explaining where electrons are distributed in simple organic structures. Along the way, you see how to determine dipoles for bonds and for molecules - an extremely useful tool for predicting solubility and reactivity of organic molecules.

Constructing Lewis Structures

The Lewis structure is the basic word of the organic chemist; these structures show which atoms in a molecule are bonded to each other and also show how many electrons are shared in each bond. You need to become a whiz at working with these structures so you can begin speaking the language of organic chemistry.

To draw a Lewis structure, follow four basic steps:

1. Determine the connectivity of the atoms in the molecule.

   Figure out how the atoms are attached to each other. Here are some guidelines:

   • In general, the central atom in the molecule is the least electronegative element. (Electronegativity decreases as you go down and to the left on the periodic table.)
   • Hydrogen atoms and halide atoms (such as F, Cl, Br, and I) are almost always peripheral atoms (not the central atom) because these atoms usually form only one bond.

2. Determine the total number of valence electrons (electrons in the outermost shell).

   Add the valence electrons for each of the individual atoms in the molecule to obtain the total number of valence electrons in the molecule. If the molecule is charged, add one electron to this total for each negative charge or subtract one electron for each positive charge.

3. Add the valence electrons to the molecule.

   Follow these guidelines:

   • Start adding the electrons by making a bond between the central atom and each peripheral atom; subtract two valence electrons from your total for each bond you form.
   • Assign the remaining electrons by giving lone pairs of electrons to the peripheral atoms until each peripheral atom has a filled octet of electrons.
   • If electrons are left over after filling the octets of all peripheral atoms, then assign them to the central atom.

4. Attempt to fill each atom's octet.

   If you've completed Step 3 and the central atom doesn't have a full octet of electrons, you can share the electrons from one or more of the peripheral atoms with the central atom by forming double or triple bonds.

   You can't break the octet rule for second-row atoms; in other words, the sum of the bonds plus lone pairs around a second-row atom (like carbon) can't exceed four.

Q. Draw the Lewis structure of CO32-.

A.

Most often, the least electronegative atom is the central atom. In this case, carbon is less electronegative than oxygen, so carbon is the central atom and the connectivity is the following:

Carbon has four valence electrons because it's an atom in the fourth column of the periodic table, and oxygen has six valence electrons because it's in the sixth column. Therefore, the total number of valence electrons in the molecule is 4 + 6(3) + 2 = 24 valence electrons. You add the additional two electrons because the molecule has a charge of -2 (if the molecule were to have a charge of -3, you'd add three electrons; if -4, you'd add four; and so forth).

Start by forming a bond between the central carbon atom and each of the three peripheral oxygen atoms. This accounts for six of the electrons (two per bond). Then assign the remaining 18 electrons to the oxygens as lone pairs until their octets are filled. This gives you the following configuration:

The result of the preceding step leaves all the oxygen atoms happy because they each have a full octet of electrons, but the central carbon atom remains unsatisfied because this atom is still two electrons short of completing its octet. To remedy this situation, you move a lone pair from one of the oxygens toward the carbon to form a carbon-oxygen double bond. Because the oxygens are identical, which oxygen you take the lone pair from doesn't matter. In the final structure, the charge is also shown:

1 Draw the Lewis structure of BF4-.

2 Draw the Lewis structure of H2CO.

3 Draw the Lewis structure of NO2-.

Predicting Bond Types

Bonds can form between a number of different atoms in organic molecules, but chemists like to broadly classify these bonds so they can get a rough feel for the reactivity of that bond. These bond types represent the extremes in bonding.

In chemistry, a bond is typically classified as one of three types:

• Purely covalent: The bonding electrons are shared equally between the two bonding atoms.
• Polar covalent: The electrons are shared between the two bonding atoms, but unequally, with the electrons spending more time around the more electronegative atom.
• Ionic: The electrons aren't shared. Instead, the more electronegative atom of the two bonding atoms selfishly grabs the two electrons for itself, giving this more electronegative atom a formally negative charge and leaving the other atom with a formal positive charge. The bond in an ionic bond is an attraction of opposite charges.

You can often determine whether a bond is ionic or covalent by looking at the difference in electronegativity between the two atoms. The general rules are as follows:

• If the electronegativity difference between the two atoms is 0.0, the bond is purely covalent.
• If the electronegativity difference is between 0.0 and 2.0, the bond is considered polar covalent.
• If the electronegativity difference is greater than 2.0, the bond is considered ionic.

Ionic and covalent bonding are extreme ends of a continuum of possibilities for how much the electrons are shared, so these numbers are just guidelines (some texts even give slightly different numbers as the cutoffs between covalent and ionic). For example, there is not a huge difference in the bonding situation arising between atoms having an electronegativity difference of 1.9 or between atoms having an electronegativity difference of 2.0, even though the first bond would be classified as polar covalent and the second one ionic. The bond with a 1.9 electronegativity difference would just have slightly more shared bonding electrons than the bond with a 2.0 difference, but in both cases the electrons would spend most of their time around the more electronegative element.

Figure 1-1 shows the electronegativity values.

FIGURE 1-1: Electronegativity values for common atoms.

Q. Using the following figure, classify the bonds in potassium amide as purely covalent, polar covalent, or ionic.

A. You classify the N-H bonds as polar covalent and the N-K bond as ionic. To determine the bond type, take the electronegativity difference between the two atoms in each bond. For the nitrogen-potassium (N-K) bond, the electronegativity value is 3.0 for nitrogen and 0.8 for potassium, giving an electronegativity difference of 2.2. Therefore, this bond is considered ionic. For the N-H bonds, the nitrogen has an electronegativity value of 3.0 and hydrogen has an electronegativity value of 2.2, so the electronegativity difference is 0.8. Therefore, the N-H bonds are classified as polar covalent.

4 Classify the bond in NaF as purely covalent, polar covalent, or ionic.

5 Using the following figure, classify the bonds in hexachloroethane as purely covalent, polar covalent, or ionic.

Determining Bond Dipoles

Most bonds in organic molecules are of the polar covalent variety. Consequently, although the electrons in a polar covalent bond are shared, on average they spend more time around the more electronegative atom of the two bonding atoms. This unequal sharing of the bonding electrons creates a separation of charge in the bond called a bond dipole.

Bond dipoles are used all the time to predict and explain the reactivity of organic molecules, so you need to understand what they mean and how to show them on paper. You represent this separation of charge on paper with a funny-looking arrow called the dipole vector. The head of the dipole vector points in the direction of the partially negatively charged atom (the more electronegative atom) and the tail (which looks like a + sign) points toward the partially positive atom of the bond (the less...