Draw the Lewis structure. B) Predict the electron group geometry. C) Predict the molecular shape. D) Predict the IDEAL bond angle(s). E) Predict the polarity (circle one): Polar / Non-polar. PO3 3- A) Draw the Lewis structure. B) Predict the electron group geometry. C) Predict the molecular shape. D) Predict the IDEAL bond angle(s). E) Predict the polarity (circle one): Polar / Non-polar.
The Correct Answer and Explanation is:
Compound 1: SCl₅F
A) Lewis Structure:
- Sulfur (S) is the central atom.
- It forms six single bonds: five to Cl atoms and one to F.
- Sulfur has 6 valence electrons, and it can expand its octet.
- Each Cl and F has 7 valence electrons, forming single bonds and completing octets with lone pairs.
B) Electron Group Geometry:
- Octahedral, as there are six bonding groups around sulfur.
C) Molecular Shape:
- Octahedral, since all positions are occupied by atoms (no lone pairs).
D) Ideal Bond Angles:
- 90° and 180° (typical for octahedral geometry).
E) Polarity:
- Polar, due to the presence of one fluorine atom (much more electronegative than Cl), causing an uneven charge distribution.
Compound 2: PO₃³⁻ (Phosphite ion)
A) Lewis Structure:
- Phosphorus (P) is the central atom.
- It forms three single bonds with three oxygen atoms.
- One of the oxygens carries a negative charge.
- The other two O atoms each carry a negative charge.
- P has one lone pair.
B) Electron Group Geometry:
- Tetrahedral, due to three bonding pairs and one lone pair on phosphorus.
C) Molecular Shape:
- Trigonal pyramidal, due to one lone pair on phosphorus.
D) Ideal Bond Angles:
- <109.5°, slightly less due to lone pair repulsion.
E) Polarity:
- Polar, because of the lone pair on phosphorus and the asymmetric shape, leading to a dipole moment.
Explanation
Understanding the geometry and polarity of molecules involves analyzing their Lewis structures, electron arrangements, and electronegativity differences. In SCl₅F, sulfur is the central atom surrounded by six atoms—five chlorines and one fluorine—forming six sigma bonds. This results in an octahedral electron geometry and molecular shape, which is highly symmetrical in a perfect case. However, because fluorine is significantly more electronegative than chlorine, it disrupts the symmetry, pulling electron density unevenly. This leads to a polar molecule, despite the symmetric shape.
In contrast, PO₃³⁻ (phosphite ion) exhibits tetrahedral electron geometry because phosphorus is surrounded by three oxygen atoms and has one lone pair. Lone pairs occupy more space than bonding pairs, so the molecular shape becomes trigonal pyramidal—similar to ammonia (NH₃). The presence of a lone pair and three oxygen atoms of varying charge distribution makes the molecule polar. Bond angles are slightly less than 109.5° due to increased lone pair repulsion.
In both molecules, shape and polarity are influenced by electron domain geometry and the nature of surrounding atoms. Molecular polarity is not just about individual bonds but how those bonds are arranged in 3D space. For SCl₅F, despite a symmetric octahedral shape, the differing atom (F) causes polarity. For PO₃³⁻, the lone pair and asymmetric oxygen arrangement lead to a polar structure. Understanding these principles helps explain reactivity, intermolecular interactions, and physical properties.
