Question

Which of the following statements are true? (a) Both nitrogen and phosphorus can form a pentafluoride compound. (b) Although CO is a well-known compound, SiO does not exist under ordinary conditions. (c) \(\mathrm{Cl}_{2}\) is easier to oxidize than \(\mathrm{I}_{2}\). (d) At room temperature, the stable form of oxygen is \(\mathrm{O}_{2}\), whereas that of sulfur is \(S_{8}\)

Short answer

Statements (a), (b), and (d) are true; statement (c) is false. Nitrogen and phosphorus can form pentafluoride compounds (NF5 and PF5), and while CO exists, SiO does not under ordinary conditions. O₂ is the stable form of oxygen at room temperature, while S8 is the stable form of sulfur. Cl₂ is not easier to oxidize than I₂.

Step-by-step solution

(a) Nitrogen and phosphorus form pentafluoride compounds

The elements nitrogen (N) and phosphorus (P), belonging to group 15 of the periodic table, form pentafluoride compounds when they bond with five fluorine (F) atoms. Nitrogen forms NF5 and phosphorus forms PF5. Therefore, this statement is true.

(b) CO is a well-known compound, but SiO does not exist under ordinary conditions

Carbon monoxide (CO) is a well-known compound that is made up of carbon (C) and oxygen (O). Carbon and silicon (Si) are in the same group 14 in the periodic table, but they behave differently due to different electronegativity values. CO forms easily because C forms double bond with O due to its smaller size and higher electronegativity; however, silicon monoxide (SiO) does not typically exist under ordinary conditions since silicon, being a larger atom with a lower electronegativity than carbon, doesn't form a stable compound with oxygen at ordinary conditions. Therefore, this statement is also true.

(c) Cl2 is easier to oxidize than I2

The oxidizing ability of an element is determined by the ease with which it gains electrons. Halogens like chlorine (Cl) and iodine (I) have a higher electronegativity due to the fact that they are just one electron short of achieving a completed shell. They usually have a strong oxidizing ability. However, as we move down the group, the electronegativity value decreases and the atom size increases, causing a decrease in oxidizing ability. Therefore, iodine (I2) is a weaker oxidizing agent compared to chlorine (Cl2). This statement is false; it is actually I2 that is easier to oxidize than Cl2.

(d) O2 is the stable form of oxygen at room temperature, and S8 is the stable form of sulfur

At room temperature, oxygen gas exists in the diatomic molecular form (O2) because O atoms form a double bond with each other, making them stable. In contrast, sulfur exists in its elemental form as an octatomic molecule (S8) due to the fact that sulfur atoms form a ring structure in which each S atom is connected to two other sulfur atoms by single bonds, resulting in a more stable configuration. Therefore, this statement is true. In summary, statements (a), (b), and (d) are true, while statement (c) is false.

Key concepts

Nitrogen Pentafluoride

Nitrogen pentafluoride, or NF extsubscript{5}, is a chemical compound formed when one nitrogen atom bonds with five fluorine atoms. Both nitrogen and phosphorus belong to group 15 of the periodic table, also known as the nitrogen group.
This group is characterized by its ability to form bonds with multiple fluorine atoms, resulting in pentafluoride compounds like NF extsubscript{5} for nitrogen and PF extsubscript{5} for phosphorus. The bonding in NF extsubscript{5} occurs because nitrogen has five valence electrons, allowing it to share these electrons with five fluorine atoms, each of which needs one more electron to achieve a full outer shell. The compound results from covalent bond formation, which is quite typical for molecules containing these nonmetals. If you find yourself wondering why nitrogen can form this specific pentafluoride compound, remember these key points:

• Group 15 elements, like nitrogen and phosphorus, commonly bond with highly electronegative elements like fluorine.
• The small size of nitrogen allows it to efficiently form strong bonds with multiple fluorine atoms.

However, NF extsubscript{5} is quite unstable, unlike phosphorus pentafluoride (PF extsubscript{5}), which is used in various chemical reactions due to its stability and reactivity.

Carbon and Silicon Compounds

Carbon and silicon are both in group 14 of the periodic table, which consists of elements known for their four valence electrons. They are similar in some ways but differ significantly in their ability to form compounds.
Carbon monoxide (CO) is a widely known compound, commonly formed by carbon readily bonding with oxygen. This is largely because carbon is smaller and more electronegative than silicon, allowing it to form a stable double bond with oxygen. In contrast, silicon monoxide (SiO) is rare and unstable under normal conditions. Silicon is larger, has a lower electronegativity, and does not form stable structures similar to CO. Instead, silicon tends to form structures like silicon dioxide (SiO extsubscript{2}), which is very stable and commonly found as quartz. Here are some things to remember about carbon and silicon compounds:

• Carbon forms strong covalent bonds due to its smaller size and ability to make double bonds.
• Silicon, being larger, forms compound structures typically with more atoms in comparison to carbon's smaller compounds.

This difference underpins why carbon and silicon do not always create similar compounds despite their proximity in the periodic table.

Oxidizing Ability

Oxidizing ability refers to a substance's tendency to gain electrons in a chemical reaction, often described as the ability to make other substances lose electrons. This concept is crucial in understanding the reactivity of elements, especially in nonmetals like the halogens.
Chlorine (Cl extsubscript{2}) and iodine (I extsubscript{2}) are both halogens, yet their oxidizing abilities differ significantly. Chlorine is a stronger oxidizing agent compared to iodine because it is higher in the group of halogens and is more electronegative. Higher electronegativity means chlorine is more likely to gain an electron to form a stable ion. It is also smaller in size compared to iodine, which allows greater attraction of electrons towards its nucleus. Consider these points when thinking about oxidizing abilities:

• Elements higher in a group tend to be stronger oxidizing agents due to higher electronegativity.
• Electronegativity decreases as you move down the group because atomic size increases, which weakens the pull on additional electrons.

Hence, this is why I extsubscript{2} oxidizes less readily than Cl extsubscript{2}, contrary to what one might initially believe. This nuance is especially important in redox reactions where selecting the right oxidizing agent can significantly impact the reaction outcome.