Question

Explain the following observations: (a) The highest fluoride compound formed by nitrogen is \(\mathrm{NF}_{3}\), 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

In short, the observations can be explained as follows: (a) Nitrogen forms NF3 and phosphorus forms PF5 due to the difference in their electron configurations and the availability of d-orbitals for phosphorus. (b) CO exists but SiO doesn't under ordinary conditions because of the electronegativity differences and bond strengths between the elements. SiO2 is preferentially formed instead. (c) AsH3 is a stronger reducing agent than NH3 due to the weaker As-H bond compared to the N-H bond, making it easier for AsH3 to donate its electron pair.

Step-by-step solution

Observation (a): Nitrogen forms NF3 while Phosphorus forms PF5

The highest fluoride compounds formed by nitrogen and phosphorus are NF3 and PF5, respectively. This difference can be explained by considering the electron configurations of these elements and their willingness to form chemical bonds. Nitrogen has the electron configuration of 1s² 2s² 2p³, with three unpaired electrons in the 2p orbitals. When nitrogen forms bonds with other atoms, it can only form three covalent bonds, as it has three unpaired electrons available for sharing. Phosphorus, in contrast, has the electron configuration of 1s² 2s² 2p⁶ 3s² 3p³, with three unpaired electrons in the 3p orbitals. However, unlike nitrogen, phosphorus is in the third period of the periodic table, so it has access to the 3d orbitals as well. When forming bonds, phosphorus can promote one of its electrons from the 3s orbital to one of the empty 3d orbitals, thus enabling it to form five covalent bonds. As a result, phosphorus can form PF5, while nitrogen can only form NF3.

Observation (b): Existence of CO but not SiO under ordinary conditions

Carbon and silicon are both group 14 elements, meaning they have four electrons in their outermost energy level. However, the bonding in CO and SiO is different due to the electronegativity differences between the elements and the strength of the bonds they form. Carbon and oxygen are both highly electronegative elements, with electronegativity values of around 2.55 and 3.44, respectively. When they combine to form CO, each bond is highly polar, which creates a very stable compound. Moreover, the shorter length of C-O bond (due to smaller atomic radii of C and O) makes the bond stronger. On the other hand, silicon and oxygen have lower electronegativity values, around 1.90 and 3.44, respectively. Despite having four valence electrons similar to carbon, silicon's larger atomic radius (due to being in the third period) results in weaker, less polar Si-O bonds compared to C-O bonds. Consequently, SiO does not readily form under ordinary conditions. Moreover, silicon prefers to form SiO2, which is a stable compound with a tetrahedral structure, consisting of one Si atom and two O atoms. The O atoms in SiO2 are bound in double bonds to the Si atom. Therefore, SiO2 is formed preferentially as opposed to SiO.

Observation (c): AsH3 is a stronger reducing agent than NH3

Arsine (AsH3) and ammonia (NH3) are both group 15 hydrides. However, their reducing abilities depend on the difference in electronegativity and bond strength between the central atom and hydrogen. The electronegativity value of nitrogen (3.04) is higher than that of arsenic (2.18). As a result, the N-H bond in NH3 is more polar and stronger as compared to the As-H bond in AsH3. In a reducing reaction, the central atom must donate its electron pair to reduce another species. A weaker central atom-hydrogen bond would make it easier to donate this electron pair. As the As-H bond is weaker than the N-H bond, AsH3 can more easily donate its electron pair, making it a stronger reducing agent than NH3. In conclusion, the differences in the observed behaviors of these compounds can be explained by factors such as electron configurations, electronegativity, bond strengths, and atomic radii of the involved elements.

Key concepts

Understanding Electron Configurations

The heart of chemical behavior lies in an element's electron configuration, which is the distribution of electrons across different orbital levels. When we mention electron configurations, we're referring to the way electrons are arranged around an atom's nucleus, which directly impacts the element's chemical properties and the types of bonds it can form.

For example, nitrogen's tendency to form \textbf{NF}\(_3\), rather than a compound like \textbf{NF}\(_5\), stems from its electron configuration, which supports the formation of three covalent bonds corresponding to its three unpaired electrons. This limitation is a direct consequence of the lack of available d orbitals in nitrogen's valence shell.

In contrast, phosphorus, being in the third period, can extend its valence shell to include the 3d orbitals. This ability to access an additional orbital allows phosphorus to accommodate more electrons and form five covalent bonds, leading to the formation of \textbf{PF}\(_5\).

Electron Configurations and Periodic Trends

It's essential to recognize that electron configurations also underpin periodic trends such as atomic size, ionization energy, and electronegativity. All these properties influence how an element will interact with others. For instance, the larger size of silicon compared to carbon results in less effective overlapping with oxygen's orbitals, ultimately explaining why SiO is not stable under ordinary conditions whereas CO is a well-known stable molecule.

Using bullet points, here are key takeaways about electron configurations:

• Electron configurations determine the possible number of covalent bonds an element can make.
• Elements in higher periods have access to more orbitals, like d orbitals, which can expand their covalent bonding capabilities.
• Periodic trends such as atomic size and electronegativity can be predicted based on an element's electron configuration.

Covalent Bonds: When Atoms Share Electrons

Covalent bonding is a type of chemical bond formed when two atoms share pairs of electrons. The number of covalent bonds an atom can form is influenced by its electron configuration, particularly by the number of unpaired electrons available for bonding.

For elements like nitrogen with a limited number of unpaired electrons, only a certain number of covalent bonds can be formed, which explains why nitrogen can't form \textbf{NF}\(_5\). Phosphorus, however, can form \textbf{PF}\(_5\) due to its ability to utilize d orbitals and hence accommodate more shared electron pairs.

Polar and Nonpolar Covalent Bonds

Not all covalent bonds are created equal. The difference in electronegativity between the bonded atoms can lead to either nonpolar covalent bonds, where electrons are shared equally, or polar covalent bonds, which have an unequal sharing of electrons. In the case of CO, the strong polarization contributes to the stability of the molecule.

Key aspects of covalent bonding include:

• The number of covalent bonds is related to the number of an atom's unpaired electrons.
• Covalent bonds can be nonpolar or polar, affecting molecule properties.
• The stability and existence of molecules are influenced by how atoms share their electrons.

Reducing Agents: Donating Electrons in Chemical Reactions

Reducing agents play a crucial role in chemical reactions, particularly in redox (reduction-oxidation) processes where electron transfer occurs. A reducing agent, by definition, loses electrons and, in the process, reduces another substance.

The strength of a reducing agent is influenced by its bond polarity and bond strength. For instance, arsine (\textbf{AsH}\(_3\)) is a stronger reducing agent than ammonia (\textbf{NH}\(_3\)) because the As-H bond is weaker and less polar than the N-H bond, making it easier for arsine to donate electrons.

Factors Affecting Reducing Strength

Several factors determine the reducing strength of a compound, including:

• Electronegativity: A lower electronegativity means that the atom holds onto its bonding electrons less tightly, which facilitates their transfer.
• Bond Strength: Weaker bonds are more easily broken, allowing for electron donation.
• Atomic Size: Larger atoms have less control over their valence electrons, potentially making them better reducing agents.

These factors help us understand why arsine is a more potent reducer than ammonia and emphasize the interconnectedness of electron configurations, periodic trends, and chemical behavior.