Question

The triiodide ion \(\left(\mathrm{I}_{3}^{-}\right)\) in which the \(\mathrm{I}\) atoms are arranged as III is stable, but the corresponding \(\mathrm{F}_{3}^{-}\) ion does not exist. Explain.

Short answer

The non-existence of F3- and the stability of I3- can be explained based on atomic size, electronegativity, and lone pair repulsion. Iodine atoms are larger with lesser electronegativity and lesser lone pair repulsion than Fluorine, leading to the formation of a stable I3- ion. Conversely, due to its significantly smaller atomic size, higher electronegativity, and greater lone pair repulsion, a corresponding F3- ion does not exist.

Step-by-step solution

Consider the Atomic Size

Examine the trends in atomic size in the periodic table. One can observe that iodine (I) is significantly larger than fluorine (F). The larger size of the Iodine atom allows the formation of more than one bond, while the smaller Fluorine atom does not permit the existence of multiatomic anions.

Consider the Electronegativity

Fluorine is the most electronegative element in the periodic table, meaning it holds on to its electrons very tightly. As a result, it is reluctant to share electrons with another fluorine atom for them to form F3- ion. Contrarily, Iodine is less electronegative, and thus forms the I3- anion easily.

Understand the Role of Lone Pair Repulsion

Due to the smaller size of Fluorine, there is a high electron density, causing large repulsions between the fluorine lone pairs. This makes the structure very unstable compared to I3-, where the repulsions between lone pairs are smaller due to the larger size of the Iodine atom. This makes I3- more stable compared to F3- being non-existent.

Key concepts

Atomic Size

When we talk about atomic size, we're discussing the relative space an atom occupies. In the periodic table, as we move down a group, the size of the atoms generally increases. Iodine (I) sits below fluorine (F) in the same group, making iodine atoms significantly larger. This difference in size is crucial when it comes to forming bonds in structures like the triiodide ion.
Iodine's larger atomic size allows it to comfortably accommodate and bind with more atoms, forming the \( \mathrm{I}_{3}^{-} \) ion. In contrast, fluorine's atomic size is quite small. This restricts its ability to form similar multiatomic species like the non-existent \( \mathrm{F}_{3}^{-} \). The steric hindrance caused by small atomic sizes in tightly-packed environments makes such formations unfavorable for fluorine.
The atomic size directly influences how atoms can fit together in a molecule. Larger atoms like iodine can accommodate more atoms around them without experiencing excessive repulsion. This is not the case for the smaller fluorine atoms, which limits their capability to form such structures.

Electronegativity

Electronegativity refers to an atom's ability to attract and hold onto electrons. In the periodic table, fluorine is the most electronegative element, meaning it strongly attracts electrons towards itself. This intense pull means that fluorine atoms prefer to hold onto their own electrons rather than sharing them, making the formation of a \( \mathrm{F}_{3}^{-} \) ion highly unlikely.
Unlike fluorine, iodine has a much lower electronegativity. This difference means iodine atoms are more willing to share their electrons, facilitating the bonding process necessary to form the triiodide ion \( \mathrm{I}_{3}^{-} \). The lesser electronegativity of iodine makes it a better candidate for creating anions with more than one atom.
In simpler terms, the higher the electronegativity, the less likely an atom is to participate in complex anion formations, as it prefers to hold onto its electrons. Whereas, atoms with lower electronegativity are more open to sharing or distributing their electron density among other atoms.

Lone Pair Repulsion

Lone pairs are pairs of valence electrons not involved in bonding. The concept of lone pair repulsion is vital in the stability of molecular structures. Lone pair repulsions are caused by the concentrated electron density of these pairs, leading to significant repulsive forces between them.
In fluorine atoms, the small size leads to a high electron density as many electron pairs are squeezed into a small space. This results in strong repulsive forces between the lone pairs, destabilizing any potential formation of a \( \mathrm{F}_{3}^{-} \) ion. The crowded environment simply makes it impossible for such a structure to exist.
On the flip side, iodine's larger atomic size allows lone pairs to be spaced out more effectively. This reduced repulsion between iodine's lone pairs contributes to the stability of the \( \mathrm{I}_{3}^{-} \) ion. Thus, fewer repulsions and better accommodation of electron pairs in larger atoms make such formations possible.