Acidic Halides: Lewis Acid Strength of Group 15 Halides Quiz

The electronegativity of the halogen attached to the Group 15 element significantly influences Lewis acidity. As the electronegativity of the halogen increases, it withdraws more electron density from the central atom. This makes the central atom more electron-deficient and enhances its ability to accept an external electron pair from a base.

Phosphorus trifluoride acts as a stronger electron pair acceptor compared to its heavier analogs. The high electronegativity of fluorine creates a significant partial positive charge on the phosphorus atom. This electron deficiency facilitates the acceptance of a lone pair from a donor molecule, following the fundamental principles of periodic table trends.

This is incorrect because Lewis acidity usually decreases as the central atom increases in size. A larger central atom has a more diffuse orbital, which leads to weaker overlap with the donor's lone pair. Additionally, lower electronegativity on the central atom reduces the driving force to attract and accommodate incoming electron density.

Back-bonding involves the donation of electron density from the halogen lone pairs into the empty orbitals of the central atom. This internal stabilization reduces the electron deficiency of the central atom. Consequently, the atom becomes less eager to accept an external electron pair, thereby lowering the overall strength of the molecule.

Antimony pentafluoride is recognized as one of the strongest known Lewis acids. Its ability to accept a fluoride ion to form the octahedral hexafluoroantimonate ion is a classic example of p-block reactivity. This property is extensively utilized in the creation of superacids, where the halide acts as a potent electron acceptor.

Phosphorus pentachloride can expand its octet by utilizing its vacant 3d-orbitals to accept additional electron pairs. The +5 oxidation state also makes the phosphorus atom highly electropositive. These features combined allow it to react with halide ions to form complex anionic species, demonstrating typical behavior for heavier p-block elements.

Arsenic trichloride is generally a stronger Lewis acid than antimony trichloride. This is due to the smaller atomic size of arsenic, which results in a higher charge density on the central atom. A more concentrated positive center is more effective at attracting and bonding with the lone pair of a Lewis base.

This statement is false. While nitrogen is electronegative, the lack of d-orbitals prevents it from expanding its coordination number. Furthermore, the lone pair on nitrogen often makes the molecule act as a Lewis base rather than an acid. This highlights the unique chemical limitations of second-period elements in the p-block.

Iodine is the least electronegative and largest of the common halogens. When it bonds with Group 15 elements, it provides the least amount of inductive electron withdrawal from the central atom. The resulting molecule is less electron-deficient, making it a much weaker electron pair acceptor compared to fluorinated or chlorinated versions.

As we move down the group, the central atom becomes larger and less electronegative, which reduces its ability to attract external electrons. The increase in metallic character also shifts the bonding nature. These combined physical changes explain why the tendency to act as an electron pair acceptor diminishes as we go from phosphorus to bismuth.

When a trihalide like phosphorus trifluoride accepts an electron pair from a fluoride ion, it typically adopts a trigonal bipyramidal geometry. This change in molecular shape is a direct result of the central atom accommodating an additional bond and its lone pair, reflecting the versatile bonding capabilities found in the p-block.

This is false because Group 13 halides like boron trifluoride are better Lewis acids specifically because they lack a lone pair and have an incomplete octet. Group 15 trihalides already have a lone pair, which can electronically repel incoming bases, often making them weaker acids or even potential Lewis bases.

Bismuth is the most metallic element in Group 15. Its trihalides exhibit more ionic character and the bismuth atom has a very low tendency to accept electron pairs covalently. This shift from covalent Lewis acidity to ionic behavior is a hallmark of moving toward the bottom of the p-block.

Increasing the oxidation state to +5 makes the central phosphorus atom much more electron-deficient. With a higher positive charge, the atom becomes a significantly stronger Lewis acid. It can more easily attract and coordinate with electron-rich species, which is a common trend observed across various inorganic p-block compounds.

Phosphorus pentaflluoride, antimony pentafluoride, and phosphorus trichloride can all function as Lewis acids. They possess either a high oxidation state or vacant orbitals that allow them to accept electron pairs from donors. In contrast, ammonia is a classic Lewis base because it uses its lone pair to donate to other electron-deficient species.