Properties of Chemical Bonds and VSEPR Theory

1. Covalent Radius: This is half the internuclear distance between two similar atoms that are covalently bonded in a diatomic element.
2. Van der Waals Radius: This is half the internuclear distance between two similar, non-bonded atoms of adjacent molecules.

Key Comparison: The covalent radius of an element is always shorter than its Van der Waals radius because covalent bonding involves the overlapping of atomic orbitals, pulling the nuclei closer together.

Bond Angle

The bond angle is the angle (measured in degrees) between the orbitals containing bonding electron pairs around the central atom in a molecule. For example, the $\text{H–O–H}$ bond angle in a water molecule is $104.5^\circ$.

Bond Enthalpy and Bond Order

Bond enthalpy is the amount of energy required to break one mole of a specific type of bond between two atoms in a gaseous state. For instance, the $\text{H–H}$ bond enthalpy in a hydrogen molecule is $435.8 \text{ kJ mol}^{-1}$.

The bond order refers to the number of bonds between two atoms in a molecule.

• $\text{H}_2$ has a bond order of 1.
• $\text{O}_2$ has a bond order of 2.
• $\text{N}_2$ has a bond order of 3.

There is a direct relationship between these properties: as the bond order increases, the bond enthalpy increases (the bond becomes stronger) and the bond length decreases (the bond becomes shorter).

VSEPR Theory: Predicting Molecular Shape

The Valence Shell Electron Pair Repulsion (VSEPR) theory is used to predict the geometry of molecules based on the behavior of electron pairs.

The Postulates of VSEPR

The theory is built upon several core principles regarding how electrons interact:

1. The geometry and shape of a molecule depend on the number of valence shell electron pairs (both bonded and non-bonded/lone pairs) around the central atom.
2. Electron pairs surrounding the central atom repel each other.
3. To minimize this repulsion, electron pairs lie as far apart as possible. The valence shell is treated as a sphere where electron pairs localize on the surface at maximum distance from one another.
4. Repulsion Hierarchy: Not all electron pairs repel equally. The strength of repulsion follows this order: $\text{Lone pair (lp)–Lone pair (lp)} > \text{Lone pair (lp)–Bond pair (bp)} > \text{Bond pair (bp)–Bond pair (bp)}$.
5. If there are zero lone pairs on the central atom, the geometry and the shape are identical. If lone pairs are present, the geometry and shape will differ.
6. Only the lone pairs on the central atom are considered when determining shape.

Molecular Geometries without Lone Pairs

When the central atom has no lone pairs, the shape is determined strictly by the number of bonding pairs:

Molecular Shapes with Lone Pairs

When lone pairs are present, they occupy more space and push the bonding pairs closer together, changing the observed shape of the molecule.

Exam Tip: When predicting shape, always count the total number of electron pairs (bonding + lone pairs) to find the geometry first, then look at the lone pairs to determine the final shape.