Question

Explain the following observations: (a) \(\mathrm{HNO}_{3}\) is a stronger oxidizing agent than \(\mathrm{H}_{3} \mathrm{PO}_{4}\) - b) Silicon can form an ion with six fluorine atoms, \(\mathrm{SiF}_{6}{ }^{2-}\), whereas carbon is able to bond to a maximum of four, \(\mathrm{CF}_{4}\). (c) There are three compounds formed by carbon and hydrogen that contain two carbon atoms each \(\left(\mathrm{C}_{2} \mathrm{H}_{2}, \mathrm{C}_{2} \mathrm{H}_{4}\right.\), and \(\left.\mathrm{C}_{2} \mathrm{H}_{6}\right)\), whereas silicon forms only one analogous compound \(\left(\mathrm{Si}_{2} \mathrm{H}_{6}\right)\).

Short answer

The difference in observations can be explained as follows: (a) HNO3 is a stronger oxidizing agent than H3PO4 because nitrogen is more electronegative than phosphorus, enabling nitrogen to accept electrons more efficiently. (b) Silicon can form an SiF6^2- ion because it has vacant d-orbitals in its valence shell, allowing it to form 6 bonds with fluorine, unlike carbon that has no vacant d-orbitals and can form only four bonds, resulting in CF4. (c) Carbon forms three compounds with hydrogen due to its ability to form various types of hybridized bonds, leading to C2H2 (triple bond), C2H4 (double bond), and C2H6 (single bond). In contrast, silicon can only form single bonds efficiently, resulting in only one analogous compound, Si2H6, due to its larger atomic size and weaker multiple bond formation capabilities.

Step-by-step solution

Observation (a): HNO3 as a stronger oxidizing agent than H3PO4

The oxidizing strength of an agent depends on the ability of its central atom to accept electrons. Nitrogen (N) is more electronegative than phosphorus (P). The electronegativity of nitrogen in HNO3 allows it to attract electrons more efficiently than the phosphorus in H3PO4, thus making HNO3 a stronger oxidizing agent than H3PO4.

Observation (b): Silicon forming SiF6^2- while carbon forms CF4

Silicon has vacant d-orbitals in its valence shell, which allows it to form 6 bonds to fluorine (an element with high electronegativity) in the SiF6^2- ion through coordination covalent bonds. Carbon, on the other hand, has no vacant d-orbitals in its valence shell and has a maximum valency of four, which limits it to forming up to four bonds with fluorine, resulting in the CF4 molecule.

Observation (c): Three carbon-hydrogen compounds vs one silicon-hydrogen compound

Carbon can form multiple types of hybridized bonds, which are sp, sp2, and sp3 in nature. These different hybridizations result in different geometries, giving rise to a variety of carbon-hydrogen compounds: - C2H2 with a triple bond (sp hybridization) - C2H4 with a double bond (sp2 hybridization) - C2H6 with single bonds (sp3 hybridization) Silicon, however, does not efficiently form hybridized bonds like carbon does. It can only form single bonds effectively, and so there is only one analogous silicon-hydrogen compound, Si2H6 (silane). This limitation arises from silicon's larger atomic size, which weakens its ability to form multiple bonds with other atoms, unlike carbon that has a smaller atomic size, allowing it to form strong multiple bonds.

Key concepts

Electronegativity

Electronegativity is a key concept in understanding why certain compounds act as stronger oxidizing agents than others. It represents the ability of an atom to attract and hold onto electrons. The more electronegative an atom is, the more it can pull electrons towards itself in a chemical bond.
In the case of ext{HNO}_3 compared to ext{H}_3 ext{PO}_4, nitrogen is more electronegative than phosphorus. This higher electronegativity in nitrogen allows ext{HNO}_3 to more effectively attract electrons, making it a stronger oxidizing agent. This means it is better at taking electrons away from other substances, which is a hallmark of a strong oxidizing agent.

d-orbitals

d-orbitals play a significant role in the chemical bonding of elements, particularly those in the third period of the periodic table and beyond, such as silicon. These orbitals are high-energy and can hold up to 10 electrons, providing additional space for bonding.
Silicon can utilize these d-orbitals to exceed the "octet rule" and form species like ext{SiF}_{6}^{2-}. This is because it can coordinate additional electron pairs from fluorine through a process known as coordinate covalent bonding. On the other hand, carbon lacks d-orbitals in its valence shell since it is in the second period of the periodic table. This limits carbon's ability to bond only within an octet framework, such as in ext{CF}_4, where it forms only four bonds.

Hybridization

Hybridization is an important concept that explains how atoms form different types of bonds using various types and numbers of orbitals. It allows atoms to form more complex structures and is influenced by factors like atomic size and the ability to form multiple bonds.
Carbon is an excellent example of hybridization versatility. It can engage in three types of hybridized bonding states:

• sp hybridization, which allows linear structures like ext{C}_2 ext{H}_2 (acetylene) with a triple bond.
• sp2 hybridization, forming trigonal planar structures like ext{C}_2 ext{H}_4 (ethylene) with a double bond.
• sp3 hybridization, leading to tetrahedral configurations like ext{C}_2 ext{H}_6 (ethane) with single bonds.

Silicon, however, doesn't share this hybridization flexibility, primarily because its larger size leads to less effective overlap of orbitals. Silicon tends to form single bonds more readily, as seen in compounds like ext{Si}_2 ext{H}_6. This limitation is why silicon forms fewer analogs to hydrogen compounds than carbon.