Question
BF₃, NF₃, and ClF₃ all have exactly three bond pairs. Yet BF₃ is trigonal planar, NF₃ is trigonal pyramidal, and ClF₃ is T-shaped. Explain the hybridization and shape of each molecule, and why three bonds lead to three different geometries.
(JEE Main 2023, similar pattern)
Solution — Step by Step
Boron in BF₃: valence electrons = 3, bonds = 3, lone pairs = 0.
Steric number = 3 + 0 = 3 → sp² hybridization
Shape: Trigonal planar with bond angles of exactly 120°. All three B-F bonds are equivalent and lie in one plane. No lone pair means the electron geometry equals the molecular geometry.
Nitrogen in NF₃: valence electrons = 5, bonds = 3, lone pairs = 1.
Steric number = 3 + 1 = 4 → sp³ hybridization
The 4 electron pairs arrange tetrahedrally, but one position is occupied by a lone pair. The molecular shape (considering only atoms) is trigonal pyramidal with bond angles of about 102° — less than the tetrahedral 109.5° because the lone pair repels more strongly than bond pairs, compressing the F-N-F angles.
Chlorine in ClF₃: valence electrons = 7, bonds = 3, lone pairs = 2.
Steric number = 3 + 2 = 5 → sp³d hybridization
The 5 electron pairs arrange in a trigonal bipyramidal pattern. The 2 lone pairs occupy equatorial positions (less 90° repulsion). The 3 bond pairs go to the remaining 3 positions (2 axial + 1 equatorial). This gives a T-shaped molecular geometry with bond angles slightly less than 90° between axial and equatorial F atoms.
| Molecule | Bond pairs | Lone pairs | Hybridization | Shape | Bond angle |
|---|---|---|---|---|---|
| BF₃ | 3 | 0 | sp² | Trigonal planar | 120° |
| NF₃ | 3 | 1 | sp³ | Trigonal pyramidal | ~102° |
| ClF₃ | 3 | 2 | sp³d | T-shaped | ~87° |
The number of bonds alone doesn’t determine shape — it’s the total number of electron pairs (bonds + lone pairs) that dictates the electron geometry, and the placement of lone pairs determines the final molecular shape.
Why This Works
VSEPR theory says electron pairs (both bonding and non-bonding) repel each other and arrange to maximise distance. Lone pairs occupy more space than bond pairs because they’re held closer to the central atom and spread out more.
The decreasing bond angle across BF₃ → NF₃ → ClF₃ reflects increasing lone pair influence. In BF₃ (0 lone pairs), the geometry is ideal. In NF₃ (1 lone pair), the geometry compresses slightly. In ClF₃ (2 lone pairs), the geometry distorts significantly.
This series beautifully demonstrates that molecular geometry depends on lone pairs as much as on bond pairs — a point examiners love to test.
Alternative Method
A quick shape prediction method: count total electron pairs around the central atom, determine the electron geometry, then “remove” the lone pairs to see the molecular shape:
- 3 pairs → trigonal planar → no lone pairs = trigonal planar (BF₃)
- 4 pairs → tetrahedral → 1 lone pair = pyramidal (NF₃)
- 5 pairs → trigonal bipyramidal → 2 lone pairs = T-shaped (ClF₃)
In trigonal bipyramidal arrangements, lone pairs always go to equatorial positions. This minimises 90° repulsions (equatorial positions have only 2 neighbours at 90°, while axial positions have 3). This rule directly explains why ClF₃ is T-shaped and not Y-shaped.
Common Mistake
Students predict that BF₃ and NF₃ have the same shape because “both have 3 bonds.” This ignores lone pairs entirely. The correct approach always starts with counting ALL electron pairs (bonding + lone). BF₃ has 3 electron domains total → trigonal planar. NF₃ has 4 electron domains total → tetrahedral electron geometry → pyramidal molecular geometry. Never skip the lone pair count.