Question

All the group 15 elements form trifluorides, but nitrogen is the only group 15 element that does not form a pentafluoride. (a) Suggest a reason why nitrogen does not form a pentafluoride. (b) The observed bond angle in \(\mathrm{NF}_{3}\) is approximately \(102.5^{\circ} \mathrm{C} .\) Use VSEPR theory to rationalize the structure of the \(\mathrm{NF}_{3}\) molecule.

Short answer

Nitrogen does not form a pentafluoride because its outer shell can only accommodate eight electrons. The \(\mathrm{NF}_{3}\) molecule forms a trigonal pyramidal shape due to three bonded electron pairs and one lone pair of electrons around the nitrogen atom, resulting in an observed bond angle of approximately \(102.5^{\circ} \mathrm{C}\).

Step-by-step solution

Analyze the Electron Configuration of Nitrogen

Nitrogen belongs to group 15 of the periodic table and has 5 electrons in its outermost shell. According to the octet rule, atoms tend to combine in such a way that each atom has eight electrons in its valence shell, giving it the same electron configuration as a noble gas.

Reason for Nitrogen Not Forming Pentafluoride

To form a pentafluoride, nitrogen would need to share five of its electrons with five fluorine atoms, which would result in a total of ten electrons around the nitrogen atom. However, nitrogen's outer shell can only accommodate eight electrons, so it does not form a pentafluoride.

Analyze the VSEPR Theory

The Valence Shell Electron Pair Repulsion (VSEPR) theory states that the electron pairs around a central atom arrange themselves as far apart as possible to minimize electron-electron repulsion.

Apply the VSEPR Theory to NF3

In the \(\mathrm{NF}_{3}\) molecule, nitrogen is the central atom. It is surrounded by three bonded electron pairs (due to the three fluorine atoms) and one lone pair of electrons. According to the VSEPR theory, this forms a trigonal pyramidal shape, resulting in a bond angle of approximately \(102.5^{\circ} \mathrm{C}\), which is observed in \(\mathrm{NF}_{3}\).

Key concepts

Group 15 Elements

Elements in Group 15 of the periodic table are also known as pnictogens. These elements include nitrogen (N), phosphorus (P), arsenic (As), antimony (Sb), and bismuth (Bi). Each of these elements has five electrons in their outermost shell. This configuration makes them quite interesting because they often seek three more electrons to achieve a stable octet configuration, similar to noble gases.

• Nitrogen is distinct among the group due to its ability to form multiple bonds, such as in N₂ gas, which is a triple bond.
• Phosphorus and the heavier elements, however, tend to form single bonds more readily.

The chemistry of Group 15 elements is heavily influenced by the number of oxidation states they can achieve, with nitrogen showing a range from -3 to +5. While the heavier elements also display this potential, their chemistry leans more toward forming compounds with oxidation states of +3 and +5. Nitrogen, by contrast, rarely forms compounds at the +5 oxidation state.

VSEPR Theory

VSEPR stands for Valence Shell Electron Pair Repulsion theory, which is essential for predicting the shapes of molecules. According to VSEPR theory, the shape of a molecule is largely determined by repulsions between electron pairs located in the valence shell of the central atom. These electron pairs can be bonded pairs, which are shared between atoms, or lone pairs, which are unshared and belong fully to the central atom.

• Lone pairs take up more space than bonded pairs because they are not shared between atoms and therefore exert greater electron-electron repulsion.
• This repulsion influences the geometry of molecules significantly.

A good example is ammonia (NH₃), where the nitrogen atom has three bonded pairs and one lone pair. According to VSEPR, this results in a tetrahedral electron pair shape, but the molecule itself is trigonal pyramidal due to the lone pair exerting extra repulsion, compressing the bond angles from the expected 109.5° to about 107°.

Trifluorides

Trifluorides are compounds where a central element, in this case from Group 15, is bonded to three fluorine atoms. An example is nitrogen trifluoride ( F₃), a stable compound due to the strong F bond with a bond angle of approximately 102.5°. This specific angle arises from the combined effects of bonded pairs and a lone pair on nitrogen.

• The molecular shape is determined by both the bonded pairs and the lone pair of electrons around nitrogen.
• Trigonal pyramidal geometry is observed due to the lone pair of electrons pushing the bonded pairs closer together.

The bond strength in these trifluorides varies across different elements in Group 15, influencing their reactivity and properties. Nitrogen trifluoride, for example, is a gas at room temperature and is used in various industrial applications, such as plasma etching in the semiconductor industry.

Pentafluorides

Unlike trifluorides, pentafluorides involve five fluorine atoms bonded to a central Group 15 element. These are typically formed by heavier elements like phosphorus ( P₅). To accommodate the five fluorine atoms, the central atom must be capable of expanding its valence shell beyond the octet configuration, using d-orbitals.

• Phosphorus pentafluoride ( P₅) has a trigonal bipyramidal shape due to the involvement of five bonded pairs.
• This geometry results from the electron pair arrangement minimizing repulsion while allowing the extra electrons to reside in d-orbitals.

Nitrogen does not form a stable pentafluoride because it lacks d-orbitals in its valence shell. Its inability to expand the valence shell and accommodate more than eight electrons makes pentafluoride formation impossible for nitrogen.