Question

Explain the following observations: (a) The highest fluoride compound formed by nitrogen is \(\mathrm{NF}_{y}\) whereas phosphorus readily forms \(\mathrm{PF}_{5}\). (b) Although \(\mathrm{CO}\) is a well-known compound, SiO does not exist under ordinary conditions. (c) \(\mathrm{AsH}_{3}\) is a stronger reducing agent than \(\mathrm{NH}_{3}\).

Short answer

Nitrogen forms NF_3 instead of NF_5 because it has only three 2p orbitals available for bonding, unlike phosphorus, which also has access to 3d orbitals. SiO does not exist under ordinary conditions because of the weaker covalent bond and poorer orbital overlap with oxygen compared to the strong and stable bonding in CO. AsH_3 is a stronger reducing agent than NH_3 due to its weaker As-H bond, which allows it to release hydrogen more easily.

Step-by-step solution

(a) Nitrogen vs Phosphorus Fluorides

Nitrogen and phosphorus are both elements in Group 15 of the periodic table, so they have similar electronic configurations. The highest fluoride compound formed by nitrogen is NF_3, while phosphorus forms PF_5. The reason for this difference is related to the number of unpaired electrons and available orbitals for bonding. Nitrogen has one 2s and three 2p orbitals, providing a total of four orbitals in its valence shell. In NF_3, nitrogen forms three single bonds with three fluorine atoms, occupying its three 2p orbitals while leaving the 2s orbital filled. Phosphorus has one 3s, three 3p, and five 3d orbitals, providing a total of nine orbitals in its valence shell. In PF_5, phosphorus forms five single bonds with five fluorine atoms, utilizing its three 3p and two 3d orbitals while leaving the 3s orbital filled. Since nitrogen does not have d orbitals in its valence shell like phosphorus, it cannot form more than three bonds, which is why the highest fluoride it forms is NF_3.

(b) CO vs SiO Stability

Carbon and silicon are both elements in Group 14 (IVA) of the periodic table, so they have similar electronic configurations. While CO is a well-known compound, SiO does not exist under ordinary conditions. This observed difference can be attributed to the difference in electronegativity and bonding between the two elements. Carbon and oxygen have similar electronegativities (2.5 and 3.5, respectively), which leads to the formation of strong, stable covalent bonds in CO. Additionally, CO has a double bond that involves 2s and 2p orbitals, which provide good overlap and strong bonding between the atoms. Silicon, on the other hand, has an electronegativity of 1.9, which creates a greater electronegativity difference with oxygen (3.5). This difference results in a weaker covalent bond between silicon and oxygen. Furthermore, silicon's 3s and 3p orbitals are significantly larger than carbon's, resulting in poorer overlap and weaker bonding. Therefore, SiO does not exist under ordinary conditions.

(c) Reducing Power of AsH_3 vs NH_3

The reducing power of a hydride refers to its ability to donate electrons and reduce other compounds. The question is asking us to compare the reducing power of AsH_3 (arsine) and NH_3 (ammonia). The reducing power of a hydride is directly related to the bond strength between the central atom and hydrogen. The weaker the bond, the more easily hydrogen can be released as a reducing agent. Arsenic and nitrogen are both in Group 15 of the periodic table, but arsenic is in the fourth period while nitrogen is in the second period. As a result, the As-H bond in AS_h3 is weaker than the N-H bond in NH_3. This is primarily due to the larger size of arsenic compared to nitrogen, which results in weaker bonding between arsenic and hydrogen because of poorer overlap. Thus, AsH_3 is a stronger reducing agent than NH_3 because it releases hydrogen more easily due to the weaker As-H bond.

Key concepts

Electronegativity

Electronegativity describes how strongly an atom attracts electrons in a chemical bond. It is crucial for understanding why certain compounds form or don't form. Generally, electronegativity increases across a period and decreases down a group in the periodic table.

• Carbon, with an electronegativity of about 2.5, forms strong bonds with oxygen (electronegativity of 3.5) due to their similar values.
• Silicon has a lower electronegativity, around 1.9, making its bonding with oxygen weaker.

This difference is why compounds like CO (carbon monoxide) are stable, while SiO is not readily formed. The stronger the electronegativity match, the more robust the bonds formed.

Covalent Bonds

Covalent bonds occur when atoms share electrons to achieve a full valence shell, creating a stable molecule. In the context of nitrogen ( NF_3) and phosphorus ( PF_5) fluorides, phosphorus can form more covalent bonds because of available d orbitals which nitrogen lacks.

• Nitrogen forms NF_3 by sharing electrons using its 2p orbitals with just three *fluorine atoms.*
• Phosphorus, having access to 3d orbitals, can share electrons with five *fluorine atoms*, forming PF_5 and achieving a higher coordination number.

The ability to form more covalent bonds leads to more complex molecule formations.

Valence Shell Orbitals

Valence shell orbitals are the outermost orbitals that participate in bonding. The electronic configuration of an element dictates how these orbitals are utilized.
Nitrogen, in the second period, uses 2s and 2p orbitals. It lacks 3d orbitals, limiting it to form NF_3:

• 2s and three 2p orbitals combine to make three bonds.

Phosphorus, from the third period, utilizes 3s, 3p, and 3d orbitals, allowing for molecule PF_5:

• 3p and two of its five 3d orbitals form up to five bonds.

The availability of more orbitals means more complex interaction possibilities.

Reducing Agent

A reducing agent is a substance that donates electrons to another compound, and is oxidized in the process. Its strength primarily depends on how easily it can release electrons. AsH_3 (arsine) versus NH_3 (ammonia) provides a great example:

• Arsine, due to weaker As-H bonds, releases hydrogen more readily, thus behaving as a stronger reducing agent.
• This weaker bonding comes from arsenic's larger atomic size and poorer orbital overlap compared to nitrogen.

In chemical reactions, the more easily a molecule can donate electrons (or hydrogen), the stronger the reducing agent it is.

Periodic Table Trends

Periodic table trends help explain and predict element behavior in bonding, reactivity, and other properties. Understanding these trends allows us to anticipate why certain elements form specific compounds or not.

• **Electronegativity** increases across a period and decreases down a group, affecting bond strength and stability.
• **Atomic size** increases down a group, impacting bond overlaps and reducing power.
• **Orbital availability** determines the kinds and number of bonds an element can form, such as in the difference between nitrogen and phosphorus.

By leveraging these trends, chemists can rationalize the formation and stability of many compounds.