2.2 intramolecular forces

At short distances

Strong repulsion between electron clouds and nuclei causes potential energy to rise steeply.

At intermediate distances

Attractive forces (nucleus–electron interactions) dominate, lowering potential energy.

Equilibrium bond length

The separation at which potential energy is minimized — the most stable bond distance.

Bond energy

The energy required to separate the atoms from equilibrium distance to infinite distance.

The lowest point on the potential energy curve represents the most stable arrangement of atoms

defining both bond length and bond strength.

In covalent bonds, the distance between bonded atoms (bond length) and the strength of the bond (bond energy) depend on both

the size of the atoms and the number of shared electron pairs (bond order).

Atomic Size:

Larger atoms have longer bond lengths due to greater distance between nuclei.

Bond order:

• Single bond: Longest and weakest.

• Double bond: Shorter and stronger.

• Triple bond: Shortest and strongest.

Higher bond order → greater electron density between nuclei

→ stronger attraction → larger bond energy.

The strength of ionic interactions (electrostatic attraction between cations and anions) can be understood using Coulomb’s law, which relates

interaction energy to the charges of the ions and the distance between them.

Greater ionic charges →

stronger electrostatic interactions.

Smaller ionic radii →

shorter internuclear distance → stronger interactions.

Lattice Energy

The potential energy required to separate a crystal into ions; larger charges and smaller ions produce higher lattice energies

Electronegativity

the tendency of an atom in a molecule to attract shared electrons toward itself. It is a fundamental factor in determining the type of bond formed between atoms.

Across a Period (left → right): Electronegativity increases because

• Effective nuclear charge increases (more protons pulling on the same shell).

• Atomic radius decreases, so electrons are held more tightly.

Down a Group (top → bottom): Electronegativity decreases because:

• Additional electron shells increase distance between nucleus and valence electrons.

• Shielding effect reduces nuclear attraction for bonding electrons.

Nonpolar Covalent Bond

• Occurs when atoms have similar electronegativity values.

• Electrons are shared nearly equally.

Polar Covalent Bond

• Occurs when atoms have different electronegativity values.

• The more electronegative atom pulls shared electrons closer, creating partial charges.

• Bond dipole points toward the more electronegative atom.

Special Notes

1. The atom with higher electronegativity becomes partially negative; the other becomes partially positive.

2. Greater electronegativity difference → stronger bond dipole.

3. All polar covalent bonds have some ionic character; bonding is a continuum between pure covalent and pure ionic.

Lone pairs repel more because

These electrons are localized on a single atom and are not shared with another atom. They are held only by the pull of one nucleus, allowing their electron cloud to spread out and occupy more space.