To understand the ion PCl6+, let’s first draw its Lewis structure. Phosphorus (P) is in the center surrounded by six chlorine (Cl) atoms. In total, phosphorus has 5 valence electrons, and each chlorine brings in 7 valence electrons, making 42 valence electrons from the chlorine atoms plus 5 from phosphorus. Since this is a cation (PCl6+), we subtract one electron, giving us a total of 46 valence electrons to distribute.
1. Start by placing phosphorus in the center and surrounding it with six chlorine atoms.
2. Each P-Cl bond counts as two electrons. So for six bonds, that accounts for 12 electrons, leaving 34 electrons to distribute.
3. Each chlorine needs three lone pairs of electrons to complete its octet, consuming 18 electrons (6 chlorines = 18 electrons), leaving us with 16 additional electrons.
4. These extra electrons are used to form expanded octets around phosphorus, resulting in an octahedral structure.
Now, let’s sketch the three-dimensional shape. The geometry of PCl6+ is octahedral because it has six bonding pairs and no lone pairs on the phosphorus atom. The 3D representation shows the phosphorus in the center with the chlorine atoms at the corners of an octahedron.
As for the geometries:
- Electronic Geometry: Octahedral (due to six regions of electron density around the phosphorus).
- Ionic Geometry: Also octahedral, as the arrangement remains the same for the ions.
In summary, the PCl6+ ion exhibits an octahedral shape with electron and ionic geometries that are consistent with its Lewis structure. This configuration allows for optimal spacing among the chlorine atoms while satisfying the octet rule for the surrounding atoms.