Molecular Compound Bonds: Single, Double & Triple Bonds

Molecular compounds consist of two or more different elements chemically combined.
Bond formation between atoms can occur in three distinct ways:
- Single bond
- Double bond
- Triple bond
Driving force behind bond formation: atoms seek a stable, noble-gas electron configuration (the octet rule).
- Octet rule: atoms gain, lose, or share electrons until they are surrounded by $8$ valence electrons (except for a few small-atom exceptions like $\text{H}$ ).
Visual representation: Lewis (electron-dot) structures show valence electrons as dots around an element’s symbol.

The rule is a conceptual guideline, not an absolute law; it best predicts bonding for main-group elements.
Practical importance:
- Explains why non-metals frequently share electrons (covalent bonding) instead of transferring them like metals/non-metals do in ionic compounds.
- Provides a straightforward way to predict molecular formulas and bond multiplicity.
Limitations and extensions (mentioned implicitly):
- Expanded octets possible for elements in Period $\ge 3$ (not covered in this transcript but conceptually related).

Definition: One pair (two electrons) shared between two atoms.
Lewis structure shorthand: one line (–) connecting element symbols.
Example discussed: Hydrogen chloride (HCl).
- $\text{Cl}$ has $7$ valence electrons ( $\text{ns}^2\text{np}^5$ configuration).
- Needs 1 more electron to reach an octet ➔ shares one electron with $\text{H}$ .
- Resulting electron pair constitutes a single covalent bond.
Key points:
- Most common bond type.
- Allows each atom to count the shared pair toward its octet.
- Line–angle formula: $\text{H–Cl}$ .

Definition: Two pairs (four electrons) shared between two atoms.
Lewis structure shorthand: two parallel lines (=).
Example discussed: Carbon dioxide ( $\text{CO}_2$ ).
- $\text{O}$ has $6$ valence electrons; needs 2 to reach an octet.
- $\text{C}$ needs 4 more electrons.
- Each O forms a double bond with C, supplying the needed electrons.
- Lewis depiction: $\text{O}=\text{C}=\text{O}$ .
Conceptual highlights:
- Arises when a single bond would leave atoms short of an octet.
- Generally shorter and stronger than single bonds (energy not explicitly given here but implied by multiplicity).

Definition: Three pairs (six electrons) shared between two atoms.
Lewis structure shorthand: three parallel lines (≡).
Example discussed: Nitrogen gas ( $\text{N}_2$ ).
- Each $\text{N}$ has $5$ valence electrons; needs 3 to reach an octet.
- By sharing three electron pairs, both N atoms obtain octets.
- Lewis depiction: $\text{N} \equiv \text{N}$ .
Important features:
- Shortest and strongest of the three bond types (bond energy highest, though numerical values are not provided here).
- Lower bond length ➔ higher bond strength, significant in biochemical stability of atmospheric $\text{N}_2$ .

Dots represent electrons; lines represent shared electron pairs.
- One line ➔ single bond ➔ $1$ shared pair.
- Two lines ➔ double bond ➔ $2$ shared pairs.
- Three lines ➔ triple bond ➔ $3$ shared pairs.
Formalism helps rapidly convey both the qualitative (types of bonds) and quantitative (number of electrons shared) aspects of molecules.

Predicting molecular shape (via VSEPR) starts with recognizing bond multiplicity—though geometry is beyond this transcript, bond counts feed directly into shape determination.
Industrial significance:
- Understanding bond strength (e.g., triple bond in $\text{N}_2$ ) explains the energy demand of the Haber process for ammonia synthesis.
Biological relevance:
- Carbon-oxygen double bonds feature prominently in carbonyl groups (proteins, sugars, lipids), influencing reactivity.
Environmental context:
- Single vs. double bonds in greenhouse gases affect infrared absorption (vibrational modes differ by bond order).

Bond Order vs. Shared Electrons:
- Single: $1$ pair
- Double: $2$ pairs
- Triple: $3$ pairs
Octet Fulfilment Examples:
- $\text{H–Cl}$ : each atom reaches $2$ (H) or $8$ (Cl) valence electrons.
- $\text{O}=\text{C}=\text{O}$ : C attains $8$ via two double bonds; each O attains $8$ .
- $\text{N} \equiv \text{N}$ : each N attains $8$ via one triple bond.

Always count electrons: confirm each atom achieves its required octet (or duet for H).
Remember line notation equivalence:
- $\text{line} = 2$ electrons.
Sketch before naming: Lewis structures clarify formula writing & nomenclature.
Bond order hierarchy: Triple > Double > Single for strength & bond energy; Single > Double > Triple for length.
Use these template examples (HCl, CO $<em>2$ , N $</em>2$ ) as benchmarks for recognizing similar bonding patterns in less familiar molecules.