3 Chemical Bonding and Lewis Structures

Why atoms bond

Chemical bonds arise from electrostatic attractions between nuclei and electrons. Forming a bond generally lowers a system’s energy, making the bonded arrangement more stable than separated atoms. Ionic and covalent bonding are useful models for describing how atoms interact, but many bonds have characteristics of both.

Ionic and covalent bonding

In Ionic bonding, oppositely charged ions attract. For example, a sodium atom can lose an electron to become $\mathrm{Na^+}$, while a chlorine atom can gain one to become $\mathrm{Cl^-}$. Their attraction acts in all directions, producing an extended crystal lattice rather than separate $\mathrm{NaCl}$ molecules. The formula $\mathrm{NaCl}$ represents the simplest whole-number ratio of ions in the solid.

In Covalent bonding, atoms share one or more pairs of electrons. Two hydrogen atoms share one pair in $\mathrm{H_2}$; two oxygen atoms share two pairs in $\mathrm{O_2}$. A line in a structure represents a shared pair: one line for a single bond, two for a double bond, and three for a triple bond.

If bonded atoms attract the shared electrons unequally, the bond is polar covalent. The atom that attracts electrons more strongly has a partial negative charge, $\delta^-$, while the other has a partial positive charge, $\delta^+$. In $\mathrm{H-Cl}$, chlorine is $\delta^-$ and hydrogen is $\delta^+$. Ionic and covalent bonding are best treated as ends of a continuum rather than as perfectly separate categories.

Takeaway: Ionic bonding describes attraction between ions; covalent bonding describes shared electron pairs. Unequal sharing gives a covalent bond polarity.

Drawing Lewis structures

A Lewis structure represents valence electrons: lines show shared pairs in bonds, and dots show lone pairs. Use this sequence to construct one:

1. Count the total valence electrons. Add one electron for each negative charge and subtract one for each positive charge.

2. Choose an atom skeleton. Usually place the least electronegative atom, other than hydrogen, in the center, then connect atoms with single bonds.

3. Complete the octets of terminal atoms, except hydrogen, which needs two electrons.

4. Place any remaining electrons on the central atom.

5. If the central atom lacks an octet, form multiple bonds by sharing neighboring lone pairs where appropriate.

6. Check the total electron count and confirm that the formal charges add up to the overall charge.

For $\mathrm{CO_2}$, the total is $16$ valence electrons: $4$ from carbon and $6$ from each oxygen. Begin with the skeleton $\mathrm{O-C-O}$, then form a double bond to each oxygen. The resulting structure is $\mathrm{O=C=O}$; each oxygen has two lone pairs, and carbon has none.

The octet rule is a useful guide, especially for second-row elements such as carbon, nitrogen, oxygen, and fluorine, but it is not universal. Hydrogen follows a duet rule. Some molecules have an odd number of electrons, while some central atoms have fewer or more than eight electrons.

Takeaway: Count first, build the skeleton, complete terminal atoms, and then check both the electron total and overall charge.

Checking formal charge

Formal charge is an electron-bookkeeping value: it assumes that bonding electrons are shared equally, even when a bond is polar. Calculate it using:

Equivalently, assign each atom all of its lone-pair electrons and half of the electrons in each bond. For $\mathrm{O=C=O}$, carbon and each oxygen have formal charge $0$. The formal charges must sum to $0$ for a neutral molecule, or to the stated charge for an ion.

When multiple valid Lewis structures are possible, prefer structures that minimize formal charges and avoid placing a negative formal charge on a less electronegative atom when a reasonable alternative exists. These are guidelines for comparing structures; they do not replace checking electron counts and octets.

Resonance and delocalized electrons

Resonance structures have the same arrangement of atoms but differ in where electrons are placed. Use a double-headed resonance arrow, $\leftrightarrow$, between them—not an equilibrium arrow. The actual species is a resonance hybrid with delocalized electrons; it is not rapidly switching between separate drawings.

For nitrate, $\mathrm{NO_3^-}$, a Lewis structure can show one $\mathrm{N=O}$ double bond and two $\mathrm{N-O}$ single bonds. Moving the double bond among the three oxygen atoms gives three equivalent resonance structures. In the actual ion, all three $\mathrm{N-O}$ bonds are equivalent and intermediate in character between single and double bonds, and the negative charge is distributed over the oxygen atoms.

Takeaway: Resonance drawings are alternative ways to represent electron placement; the real bonding is delocalized.

Bond order, length, and energy

Bond order describes the number of shared electron pairs between two atoms: a single bond has order $1$, a double bond order $2$, and a triple bond order $3$. For the same pair of atoms, higher bond order generally corresponds to a shorter and stronger bond. A carbon–carbon triple bond is generally shorter and stronger than a carbon–carbon double bond, which is generally shorter and stronger than a carbon–carbon single bond.

A bond length is the average distance between the nuclei of bonded atoms. A bond energy is the energy required to break a specified bond in gaseous molecules; values given for a bond type across different compounds are averages. In ionic solids, the collective attraction among ions is described by lattice energy, rather than by treating the solid as separate ion pairs.

Takeaway: For bonds between the same atoms, greater bond order generally means a shorter, stronger bond. Ionic solids are described by the collective attractions in their lattice.