Question

The Group 5 \(\mathrm{A}\) elements can form molecules or ions that involve three, five, or six covalent bonds; \(\mathrm{NH}_{3}, \mathrm{AsCl}_{5},\) and \(\mathrm{PF}_{6}-\) are examples. Draw the Lewis structure for each of these substances, and predict the molecular structure and hybridization for each. Why doesn't \(\mathrm{NF}_{5}\) or \(\mathrm{NCl}_{6}-\) form?

Short answer

The Lewis structures for NH₃, AsCl₅, and PF₆⁻ are as follows: NH₃: N: •• | H:-(N)-H: | H: AsCl₅: Cl: | Cl:-(As)-Cl: | Cl: PF₆⁻: F: | F:-(P)-F: | | F+-+-F: The molecular structures and hybridizations are NH₃ (trigonal pyramidal, sp³), AsCl₅ (trigonal bipyramidal, sp³d), and PF₆⁻ (octahedral, sp³d²). NF₅ and NCl₆⁻ don't form because nitrogen's small size and electronegativity restrict it to only accommodate four electron groups, due to the absence of available d-orbitals in its second energy level.

Step-by-step solution

Drawing the Lewis structures

First, we will draw the Lewis structures for NH₃, AsCl₅, and PF₆⁻. For NH₃: N has 5 valence electrons, and H has 1 valence electron. Since there are three H atoms, we have a total of 5 + (3×1) = 8 valence electrons. In NH₃, N forms three single bonds with three H atoms. N: •• | H:-(N)-H: | H: For AsCl₅: As has 5 valence electrons, and Cl has 7 valence electrons. Since there are five Cl atoms, we have a total of 5 + (5×7) = 40 valence electrons. In AsCl₅, As forms five single bonds with five Cl atoms. Cl: | Cl:-(As)-Cl: | Cl: For PF₆⁻: P has 5 valence electrons, F has 7 valence electrons, and since it has a negative charge, we add 1 more electron. Since there are six F atoms, we have a total of 5 + 1 + (6×7) = 48 valence electrons. In PF₆⁻, P forms six single bonds with six F atoms. F: | F:-(P)-F: | | F+-+-F:

Predicting the molecular structures

Using VSEPR theory, we predict the molecular structures for each: NH₃: N is surrounded by three bonding pairs (with H) and one lone pair. So, its molecular structure is trigonal pyramidal. AsCl₅: As is surrounded by five bonding pairs (with Cl). So, its molecular structure is trigonal bipyramidal. PF₆⁻: P is surrounded by six bonding pairs (with F). So, its molecular structure is octahedral.

Determining hybridization

Now, we'll determine the hybridization for each central atom: NH₃: N has three sigma bonds and one lone pair, which gives it an electron group count of 4. Thus, N undergoes sp³ hybridization. AsCl₅: As has five sigma bonds, which gives it an electron group count of 5. Thus, As undergoes sp³d hybridization. PF₆⁻: P has six sigma bonds, which gives it an electron group count of 6. Thus, P undergoes sp³d² hybridization.

Why NF₅ and NCl₆⁻ don't form

Nitrogen, which is in Group 5A, has only five valence electrons. Given its relatively small size and electronegativity, it can only expand its valence shell to accommodate four electron groups by forming three single bonds and having one lone pair (as seen in NH₃). Therefore, it cannot form NF₅ (which would require five single bonds) or NCl₆⁻ (which would require six single bonds). Nitrogen's inability to expand its valence shell beyond four electron groups is mainly due to the absence of available d-orbitals in its second energy level.

Key concepts

VSEPR Theory

VSEPR theory, which stands for Valence Shell Electron Pair Repulsion theory, is a model used to predict the geometry of molecules. The theory posits that electron pairs around a central atom will arrange themselves as far apart as possible to minimize repulsion. This arrangement results in distinct molecular shapes.
For example, in the molecule \( ext{NH}_3\), the nitrogen atom is surrounded by three hydrogen atoms and one lone pair of electrons. Due to these electron pair repulsions, \( ext{NH}_3\) adopts a trigonal pyramidal shape. Similarly, \( ext{AsCl}_5\) is encircled by five chlorine atoms, leading to a trigonal bipyramidal structure.
Finally, \( ext{PF}_6^-\) is accompanied by six fluorine atoms, shaping into an octahedral form. Each molecular shape is the direct result of minimizing electron pair repulsions through the application of VSEPR theory.

Molecular Structure

Understanding molecular structure is crucial as it dictates the physical and chemical properties of a molecule. Molecular structure describes the arrangement of atoms within a molecule. This arrangement influences how a molecule interacts with other molecules.
In \( ext{NH}_3\), the molecular structure is trigonal pyramidal because of the lone pair on nitrogen that creates an uneven distribution of electron density. This affects its polarity and reaction behavior.
For \( ext{AsCl}_5\), the atoms are arranged in a trigonal bipyramidal structure which is crucial for understanding its symmetry and reaction types.
The octahedral structure of \( ext{PF}_6^-\) means that it’s highly symmetrical, typically making it more stable in reactions. By identifying molecular structure, we predict properties like angle, polarity, and reactivity.

Hybridization

Hybridization is a concept in chemistry where atomic orbitals mix to form new hybrid orbitals. These orbitals influence molecule shape and bonding properties. For \( ext{NH}_3\), the hybridization is sp³ because nitrogen forms three sigma bonds and contains one lone pair, comprising four electron groups.
In \( ext{AsCl}_5\), the central arsenic atom has sp³d hybridization, arising from five sigma bonds shared with chlorine atoms. This hybridization allows for five electron groups.
Meanwhile, \( ext{PF}_6^-\) involves sp³d² hybridization. The phosphorus atom forms six sigma bonds, leading to an electron group count of six. Hybridization offers insight into the geometry and dimensional shape of molecules, essential for predicting behavior and interactions at the atomic level.

Covalent Bonds

Covalent bonds occur when atoms share electrons. This sharing results in molecules like \( ext{NH}_3\), \( ext{AsCl}_5\), and \( ext{PF}_6^-\).
The key characteristic of covalent bonds is that nonmetals like nitrogen, arsenic, phosphorus, hydrogen, chlorine, and fluorine prefer to achieve a stable electron configuration by sharing electrons.
In the case of \( ext{NH}_3\), nitrogen shares its electrons with hydrogen atoms forming strong covalent bonds.
Similarly, in the molecules \( ext{AsCl}_5\) and \( ext{PF}_6^-\), the arsenic and phosphorus atoms share electrons respectively with chlorine and fluorine atoms to form stable compounds. Understanding covalent bonds is vital in predicting molecular stability and possible interactions with other chemicals.

Valence Electrons

Valence electrons are the outermost electrons of an atom and are crucial in determining how atoms bond in molecules. These electrons participate in forming bonds to create stable electron configurations.
In \( ext{NH}_3\), nitrogen has five valence electrons, which it shares with three hydrogen atoms. This results in a complete valence shell, following the octet rule.
For \( ext{AsCl}_5\), arsenic's five valence electrons are shared with five chlorine atoms, highlighting its ability to expand its valence shell due to available d-orbitals.
Phosphorus in \( ext{PF}_6^-\) holds five valence electrons, sharing them with six fluorine atoms inclusive of an additional electron due to the negative charge. Understanding valence electrons helps in drawing Lewis structures and predicting how molecules will form covalent bonds.