Phosphorus trichloride is an inorganic compound with the chemical formula PCl₃. A colorless liquid when pure, it is an important industrial chemical, being used for the manufacture of phosphites and other organophosphorus compounds. It is toxic and reacts readily with water to release hydrogen chloride.
The molecular shape of PCl3 is trigonal pyramidal. In this configuration, the central phosphorus atom is bonded to three chlorine atoms, and an unshared pair of electrons on phosphorus creates a pyramidal shape.
To draw the Lewis structure for PCl3 (phosphorus trichloride), follow these steps:
Count the Valence Electrons:
- Phosphorus (P) has 5 valence electrons.
- Chlorine (Cl) has 7 valence electrons, and since there are three chlorine atoms, that’s 3×7=21 electrons.
- Total valence electrons = 5+21=26.
Determine the Central Atom:
- Phosphorus is less electronegative than chlorine and can expand its octet, making it the central atom.
Place the Atoms:
- Place P in the center and the three Cl atoms around it.
Form Bonds:
- Connect P to each Cl with a single bond. Each bond uses 2 electrons, so 3 bonds use 6 electrons.
Distribute Remaining Electrons:
- You have used 6 electrons for bonding, leaving 26−6=20 electrons.
- Place these electrons as lone pairs around the atoms to satisfy the octet rule. Each Cl needs 6 more electrons (3 pairs), and P will get the remaining electrons:
- Each Cl gets 6 electrons (3 lone pairs), using 3×6=18 electrons.
- This leaves 2 electrons which go on P as a lone pair.
Check Octets:
- Each Cl has an octet (8 electrons: 2 in bond, 6 as lone pairs).
- P has an octet too (8 electrons: 6 in bonds, 2 as a lone pair), though P can exceed the octet, it doesn’t need to here.
Formal Charges Check (optional for this molecule but good practice):
- Cl: 7 valence e- – (6 non-bonding e- + 1 bonding e-) = 0 formal charge.
- P: 5 valence e- – (2 non-bonding e- + 3 bonding e-) = 0 formal charge.
Here’s how the Lewis structure looks:
Each colon (:) represents a pair of electrons, either as a bond or as a lone pair. Phosphorus has one lone pair, and each chlorine has three lone pairs.
This structure shows PCl3 with phosphorus having a lone pair, making the molecule’s shape trigonal pyramidal.
Here’s what this structure tells you:
- Hybridization:
- The phosphorus in PCl3 undergoes sp3 hybridization. This is because it forms 4 bonds (3 sigma bonds with Cl and one lone pair), which requires four hybrid orbitals.
- Molecular Geometry:
- The shape of PCl3 is trigonal pyramidal. Although it starts with a tetrahedral electron pair geometry (due to the four electron pairs), the lone pair on phosphorus pushes the Cl atoms down, resulting in a pyramidal shape.
- Bond Angles:
- Due to the repulsion by the lone pair, the bond angles are slightly less than the typical 109.5° of a perfect tetrahedron, generally around 107° for PCl3.
Remember, the lone pair on phosphorus contributes to the molecule’s shape, making it not just a planar triangle but a pyramid where the P atom is at the apex and the three Cl atoms form the base.