Intramolecular Forces and Potential Energy

Intramolecular forces

• A force that holds atoms together within a molecule; these are chemical bonds: ionic, covalent, metallic

• Stronger than intermolecular forces

• Example: The O–H bond in water

Potential energy in bonds

• The energy stored in chemical bonds due to the positions of atoms

• When bonds form → PE decreases

• When bonds break → PE increases

Bond formation releases energy

• When atoms form a bond, they become more stable → release energy → potential energy decreases

• Exothermic process

• Think: “Lower energy = more stable”

Bond breaking requires energy

• To break a bond, you must add energy → potential energy increases

• Endothermic process

• Like pulling magnets apart

Relationship between bond length and bond energy

• As bond length decreases, bond energy increases

• Triple bonds are shortest and strongest

• Single bonds are longest and weakest

Stable bond has minimum potential energy

• Atoms naturally move toward the bond length where potential energy is lowest

• Too close → repulsion ↑ → PE ↑

• Too far → no attraction → PE high

• Just right → lowest PE → most stable

Potential energy diagram

• A graph showing how potential energy changes as two atoms approach each other

• X-axis: Distance between nuclei

• Y-axis: Potential energy

• Bottom of curve = bond length

• Depth of curve = bond energy

• Potential Energy

  ^

  | Atoms far apart

  | High PE

  |

  |

  |

  | • ← Lowest PE = bond length

  | / \

  | / \

  +--------------------------------> Distance

  Bond forms Bonds break

Why atoms don’t collapse

• Because at very short distances, nuclei repel each other (positive vs. positive)

• This repulsion balances the attraction between electrons and nuclei → creates stable bond length

Multiple bonds have lower potential energy

• Double and triple bonds have:

• Shorter bond lengths

• Lower potential energy

• Greater stability

• Triple bond < Double bond < Single bond

  in terms of potential energy (lowest to highest)

Energy change during reaction

• In any reaction:

• Energy released = Σ(bond energies broken) – Σ(bond energies formed)

• If more energy released than absorbed → exothermic

• If more absorbed → endothermic

Bond energy calculations

• Estimate ΔH using bond energies:

• ΔH ≈ Σ (bonds broken) – Σ (bonds formed)

• Break bonds → absorb energy (+)

• Form bonds → release energy (–)

Why multiple bonds are stronger

More shared electrons → greater attraction between nuclei and electron cloud → shorter, stronger bond → harder to break

Nonpolar vs. polar covalent

• Pure Covalent vs. Polar Covalent

• Nonpolar covalent: Equal sharing (ΔEN < 0.5)

• Polar covalent: Unequal sharing (ΔEN 0.5–1.7) → dipole moment

Resonance lowers potential energy

• Molecules with resonance (like benzene, CO₃²⁻) have lower potential energy than any single Lewis structure suggests

• Delocalized electrons = extra stability

• Called resonance stabilization energy

Formal charge helps find most stable structure

The Lewis structure with formal charges closest to zero (and negative charge on most electronegative atom) has the lowest potential energy → most stable