Organic Covalent Bonding

Ionic bonds hold atoms together through electrostatic forces. Covalent bonds operate through an entirely different means: the sharing of electrons. By sharing electrons, two atoms can mutually complete their valence shells to become more stable. A molecule is a collection of atoms held together by covalent bonds. For example, below, two hydrogen atoms, each with a single electron, can share their electrons to form a covalent bond and create the diatomic hydrogen molecule. In this molecular state, both individual hydrogen atoms attain the noble gas configuration of Helium. Note: the simplest way to represent molecules is to use a Lewis Dot structure, which is what you see in the diagram below. We will explain how to draw Lewis structures later on in this section. The pair of electrons that form a bond between two atoms are called bonding electrons; as can be seen in the diagram, bonding electrons are commonly drawn as a line between the two atoms.

Since hydrogen is the simplest of atoms, having only one electron, diatomic hydrogen is the simplest of molecules, having only a single covalent bond. When more complex atoms form covalent bonds, the molecules they form are also more complex, involving numerous covalent bonds. In some instances, an atom will have valence electrons that are not involved in bonding. These valence electrons are known as lone pairs. The Lewis structures of some common atoms are shown below. Notice that each structure satisfies the octet rule for all its atoms.

When two atoms share a single pair of electrons, the bond is referred to as a single bond. Atoms can also share two or three pairs of electrons in the aptly named double and triple bonds. The first bond between two atoms is called the σ (sigma) bond. All subsequent bonds are referred to as π (pi) bonds. In Lewis structures, multiple bonds are depicted by two or three lines between the bonded atoms. The bond order of a covalent interaction between two atoms is the number of electron pairs that are shared between them. Single bonds have a bond order of 1, double bonds 2, and triple bonds 3. Bond order is directly related to bond strength and bond length. Higher order bonds are stronger and shorter, while lower order bonds are weaker and longer. /PARAGAPH The Lewis structures for some common molecules involving multiple covalent bonds can be found below.

Not all covalent bonds are fit for Sesame Street: some covalent bonds are shared unequally. Some atoms have a greater ability to attract electrons to themselves than do others. The tendency for an atom to attract electrons is its electronegativity. Don't confuse electronegativity with electron affinity. While both are periodic properties that exhibit similar trends, electron affinity is a measure of energy whereas electronegativity is simply a measure of attraction based on an arbitrary scale. Fluorine, at the upper right hand corner of the periodic table, is the most electronegative element and is assigned an electronegativity of 4.0 while other elements are regarded relative to fluorine. Electronegativity increases from left to right across the periodic table, and decreases as you move down a group.

When covalent bonds are formed between atoms of different electronegativities the result is that the shared electrons skew more toward one atom than the other. The resulting molecule is a dipole: the more electronegative atom in the bond gains a partial negative charge while the less electronegative atom becomes partially positive. The resulting bond is called a polar covalent bond. In Lewis structures, a cross-ended arrow is used to represent such polar bonds, with the arrow pointing to the more electronegative element. Partial charge symbols δ⁺ and δ^- are used to represent polarity.

Molecules with polar covalent bonds can result in molecules with overall polar attributes. The measure of the overall polarity of a molecule is called the dipole moment. As the value of the dipole moment increases, so does the polarity of the molecule.

It is possible to get an estimate of where charges in a molecule are likely to lie by inspecting its Lewis structure and assigning formal charges to specific atoms. To obtain the formal charge of an atom:

1. Take the atom's group number.
2. Subtract the number of lone pairs on the atom.
3. Subtract half the number of bonding electrons.

For example, let's examine the molecule methane, seen in . Carbon is group IV; the carbon atom in methane has zero lone pairs and eight electrons involved in bonds around it. Following the steps listed above: 4 - 0 - 1/2(8) = 0. In this Lewis structure Carbon has no Formal charge.

Remember that formal charges are just a bookkeeping tool and do not necessarily represent actual charges. However, the sum of formal charges on a molecule must equal its net charge.

As you have seen throughout this section, the simplest way to represent and describe molecules is to use a Lewis structure. The Lewis structure model generally follows the octet rule and provides a framework to understand covalent bonding. Lewis structures represent valence electrons as dots and bonding electrons as lines. Lewis structures do not represent inner electrons; only valence electrons are shown.

Here we give a step-by-step procedure for writing valid Lewis structures for any given molecular formula:

1. Count the total number of valence electrons by summing the group numbers of all the atoms. If there is a net positive charge, subtract that number from the total electron count. If there is a net negative charge, add that number to the total electron count.
2. Draw single bonds to form the desired connectivity.
3. Add lone pairs and multiple bonds, keeping the octet rule in mind.
4. Add formal charges as needed.

An example: Notice that it is possible to write two separate valid Lewis structures for the given connectivity. This phenomenon is discussed in the next section, Resonance.

You have seen that carbon tends to form four bonds, nitrogen three, oxygen two, and hydrogen/halogens one (remember also: as the number of bonds of an atom decreases, the number of its lone pairs increases). The number of bonds that a neutral atom forms is called its valence. Hence carbon is tetravalent, nitrogen is trivalent, oxygen is divalent, and so on. However, a carbon atom, for example, can be tetravalent in a number of different ways. The following chart shows a number of common bonding motifs for carbon, nitrogen, oxygen, and hydrogen. The majority of motifs in the table above obey the octet rule, with one exception. Carbon can form a trivalent species with a positive charge. This phenomenon is known as a carbocation. Such a carbon atom is extremely unstable due to its lack of an octet, but its reactivity makes it a source of a great deal of fascinating chemistry that we will later discuss.

While atoms can occasionally be short of a full octet, elements in the first two rows of the periodic table can never exceed the octet. Students often make the dreaded mistake of drawing pentavalent carbons. Never do this! (However, if you do, rest assured that all organic students make this mistake at sometime or other.) Elements in row 3 and above can exceed the octet by using d-orbitals.