The Nature of Covalent Bonds, the Octet Rule, and Lewis Formulas
You're trying to solve a puzzle where each piece has to fit perfectly to complete the picture.
In chemistry, atoms face a similar challenge—they "fit together" by sharing electrons to form stable arrangements.
The Nature of Covalent Bonds
When two non-metal atoms approach each other, their nuclei attract the electrons of the other atom. Simultaneously, repulsive forces between the nuclei and between the electrons arise. At just the right distance, the attractive forces dominate, resulting in a covalent bond.
Covalent bond
A covalent bond is a chemical bond formed by the electrostatic attraction between a shared pair of electrons and the positively charged nuclei of the bonded atoms.
Analogy
Picture two magnets being drawn together. The attraction between their opposite poles is analogous to the forces at play in covalent bonding.
Single, Double, and Triple Bonds
Atoms can share one, two, or three pairs of electrons to form different types of bonds:
- Single bond: One shared pair of electrons (e.g., H₂, where each hydrogen atom contributes one electron).
- Double bond: Two shared pairs of electrons (e.g., O₂).
- Triple bond: Three shared pairs of electrons (e.g., N₂).
Example
- In a water molecule (H₂O), each hydrogen atom shares one electron with oxygen.
- This sharing allows oxygen to complete its octet and hydrogen to achieve a stable configuration with two electrons.
Tip
Triple bonds are shorter and stronger than double bonds, which are, in turn, shorter and stronger than single bonds. Keep this in mind when analyzing molecular stability.
Note
Bonds nature will be covered more in detail in the next section.
The Octet Rule
Atoms often form bonds to achieve a full valence shell, which typically consists of eight electrons. This principle, known as the octet rule, reflects the stability of noble gas electron configurations.
Octet rule
Atoms are driven by their quest for stability, often achieved by attaining a noble gas electron configuration—commonly referred to as the octet rule.
Why the Octet Rule?
The octet rule arises from the stability associated with a filled valence shell.
Example
- Carbon (C) has four valence electrons and can form four covalent bonds to complete its octet, as seen in methane (CH₄).
- Oxygen (O) has six valence electrons and forms two covalent bonds to complete its octet, as in water (H₂O).
Exceptions to the Octet Rule
While the octet rule is a useful guideline, there are notable exceptions:
- Incomplete octets: Elements like boron (B) and beryllium (Be) are stable with fewer than eight electrons. For instance, boron trifluoride (BF₃) has only six electrons around boron.
- Expanded octets: Elements in Period 3 and beyond, such as sulfur (S) and phosphorus (P), can have more than eight electrons due to available d-orbitals. For example, sulfur hexafluoride (SF₆) has 12 electrons around sulfur.
- Odd-electron species: Molecules like nitric oxide (NO) have an odd number of electrons, making it impossible for all atoms to achieve an octet.
Example
- In BF₃, boron forms three single bonds with fluorine atoms, resulting in only six electrons in its valence shell.
- This makes BF₃ an electron-deficient molecule.
Common Mistake
It's a common misconception that all atoms must follow the octet rule. Remember that hydrogen (H) only needs two electrons, and some elements can have incomplete or expanded octets.
Lewis Formulas: Representing Molecules
Lewis formula
The Lewis formula is a visual representation of the valence electrons in a molecule.
It shows how electrons are shared (bonding pairs) or remain unshared (lone pairs).
Steps to Draw Lewis Formulas
- Count total valence electrons: Add up the valence electrons of all atoms in the molecule. For polyatomic ions, adjust for the charge (add electrons for negative charges, subtract for positive charges).
- Determine the central atom: Usually the least electronegative atom (except hydrogen, which is always terminal).
- Connect atoms with single bonds: Each bond represents two electrons.
- Distribute remaining electrons: Place lone pairs around atoms to satisfy their octet (or duet for hydrogen).
- Form multiple bonds if needed: If the central atom lacks an octet, convert lone pairs on surrounding atoms into double or triple bonds.
Example
Drawing the Lewis formula for carbon dioxide (CO₂):
- Carbon has 4 valence electrons, and each oxygen has 6, for a total of 16 electrons.
- Carbon is the central atom.
- Connect carbon to each oxygen with single bonds (4 electrons used).
- Place lone pairs around oxygen atoms to complete their octets (12 electrons used).
- Carbon lacks an octet, so forms double bonds with each oxygen.
The final structure is: O=C=O.
Molecules with Fewer or More Than Eight Electrons
- Fewer than 8 electrons: BF₃ has three single bonds and no lone pairs on boron.
- Expanded octets: In SF₆, sulfur forms six single bonds with fluorine, using 12 valence electrons.
Note
In molecules with expanded octets, the central atom belongs to Period 3 or beyond, where d-orbitals are available for bonding.
Applications of Lewis Formulas
Predicting Molecular Shapes
Lewis formulas provide the foundation for predicting molecular geometry using the Valence Shell Electron Pair Repulsion (VSEPR) theory.
Valence Shell Electron Pair Repulsion (VSEPR) theory
Valence Shell Electron Pair Repulsion (VSEPR) theory states that electron pairs repel each other and arrange themselves as far apart as possible.
Formal Charge
Lewis formulas also help calculate formal charges, which are used to determine the most stable resonance structure. The formal charge of an atom is calculated as:
Reflection
Self review
Draw the Lewis formula for SO₄²⁻ and calculate the formal charges on sulfur and oxygen atoms. Which resonance structure is most stable?
Theory of Knowledge
- How does the octet rule reflect our human tendency to simplify complex systems?
- Are there parallels in other areas of knowledge where exceptions challenge established rules?
