Which one is more stable, [PCl4]+ or [PCl6]- ? Why?
I'm thinking it has something to do with the fact that the former is tetrahedral and the latter is octahedral, but I'm not sure. Does that affect stability?
[PCl4]^+ is sp3 which is a tetragedral shape; however, note that the electrons fit exactly; i.e., P = 5 and 4*7 = 28 for Cl for total of 33 for PCl4 or 32 for [PCl4]^+.
For PCl6, it is octahedral which is d2sp3; however, adding an electron to make it [PCl6]^- adds an extra electron to an already filled d2sp3 and it really has no place to go. So I think the latter is less stable.
See this site.
https://www.google.com/search?q=lewis+stgructure+PCl4%5E%7C&ie=utf-8&oe=utf-8&client=firefox-b-1
Yes, you are correct that the structural differences between [PCl4]+ and [PCl6]- can impact their stability. The stability of a species can be influenced by several factors, including electronic configuration and geometry.
In the case of [PCl4]+ (tetrahedral geometry), the central phosphorus atom is surrounded by four chlorine atoms, forming a tetrahedral electron pair geometry. This arrangement allows for optimal electron repulsion, resulting in a more stable structure.
On the other hand, [PCl6]- (octahedral geometry) consists of a central phosphorus atom surrounded by six chlorine atoms, forming an octahedral electron pair geometry. The presence of six chlorine atoms in an octahedral arrangement leads to greater electron-electron repulsion compared to the tetrahedral arrangement. This increased repulsion can make [PCl6]- less stable compared to [PCl4]+.
Additionally, stability can also be influenced by other factors such as bond lengths, bond angles, and the overall energy of the molecule. However, in this case, the difference in geometry plays a significant role in determining the relative stability between [PCl4]+ and [PCl6]-.
The stability of ions can be understood by considering various factors, such as the electronic configuration, molecular geometry, and charge distribution. In the case of [PCl4]+ and [PCl6]- ions, the difference in stability can indeed be attributed to their molecular geometries.
[PCl4]+ consists of a central phosphorus (P) atom bonded to four chlorine (Cl) atoms, resulting in a tetrahedral geometry. On the other hand, [PCl6]- has a central P atom bonded to six Cl atoms, giving rise to an octahedral geometry.
To determine stability, one key factor to consider is electron-electron repulsion. In an octahedral geometry, the six Cl atoms surrounding the central P atom will have increased electron-electron repulsion due to their close proximity. This repulsion is higher compared to a tetrahedral geometry, where only four Cl atoms are present.
Additionally, in an octahedral geometry, the P atom experiences a higher ligand field stabilization energy (LFSE) compared to a tetrahedral arrangement. LFSE refers to the stabilization of the d orbitals of the central atom when surrounded by ligands. In an octahedral geometry, LFSE is greater because there are more ligand-donor orbitals pointing towards the central P atom, resulting in a stronger interaction and thus increased stability.
Therefore, considering the higher electron-electron repulsion and stronger LFSE in the octahedral geometry, the [PCl4]+ ion is generally more stable than the [PCl6]- ion.