Question

Explain each of the following observations: (a) At room temperature I \(_{2}\) is a solid, Br \(_{2}\) is a liquid, and \(C l_{2}\) and \(F_{2}\) are both gases. (b) \(F_{2}\) cannot be prepared by electrolytic oxidation of aqueous \(\mathrm{F}^{-}\) solutions. (c) The boiling point of HF is much higher than those of the other hydrogen halides. (d) The halogens decrease in oxidizing power in the order \(\mathrm{F}_{2}>\mathrm{Cl}_{2}>\mathrm{Br}_{2}>\mathrm{I}_{2} .\)

Short answer

The different states of halogens at room temperature can be explained by the increase in intermolecular London dispersion forces as we go down the halogen group due to an increase in atomic size and number of electrons. F₂ cannot be prepared by electrolytic oxidation of aqueous F⁻ solutions because of its strong oxidizing nature and tendency to oxidize water. The boiling point of HF is higher than other hydrogen halides due to the presence of hydrogen bonding between HF molecules. The oxidizing power of halogens decreases in the order F₂ > Cl₂ > Br₂ > I₂, due to a decrease in electron affinity as atomic size increases down the group.

Step-by-step solution

Explanation for observation (a)

In order to explain why I₂ is a solid, Br₂ is a liquid, and Cl₂ and F₂ are both gases at room temperature, we need to consider the intermolecular forces acting in each of these substances. All the halogens are nonpolar molecules, so the primary intermolecular force between them is the London dispersion force (also called van der Waals force). As we go down the halogen group (from F₂ to I₂), the size and number of electrons in the halogen atoms increase, which in turn increases the London dispersion forces between the molecules. This results in a higher boiling and melting point for the heavier halogens. Therefore, at room temperature, I₂ (the heaviest halogen) is a solid, Br₂ is liquid, and Cl₂ and F₂ (lighter halogens) are gases.

Explanation for observation (b)

The second observation states that F₂ cannot be prepared by electrolytic oxidation of aqueous fluoride ion solutions. This phenomenon is due to the strong oxidizing nature of F₂ and its highly negative electron affinity. In an aqueous solution, water molecules can potentially act as an oxidizing agent. Since F₂ is a stronger oxidizing agent compared to water, it tends to oxidize water rather than allowing the oxidation of F⁻ to F₂. Consequently, it is challenging to prepare F₂ by electrolytic oxidation of aqueous fluoride ion solutions.

Explanation for observation (c)

The boiling point of HF is much higher than those of other hydrogen halides because of the presence of hydrogen bonding in HF. The F atom in HF is highly electronegative, creating a polar bond with hydrogen. This difference in electronegativity leads to a partial positive charge on the H atom and a partial negative charge on the F atom. Due to the resulting partial charges, the positively charged H atom in one HF molecule is attracted to the negatively charged F atom in another HF molecule, creating a strong intermolecular hydrogen bond. This type of bond is stronger than the dipole-dipole interactions in other hydrogen halides, which results in a higher boiling point for HF.

Explanation for observation (d)

The halogens decrease in oxidizing power in the order: F₂ > Cl₂ > Br₂ > I₂. As mentioned earlier, the oxidizing ability of a halogen depends on their electron affinity – the ability of an atom to attract an electron. The smaller the halogen atom, the stronger its attraction towards electrons due to the proximity of the nucleus to the valence electron shell. Therefore, a smaller halogen atom will have a higher electron affinity, making it a better oxidizing agent. Since fluorine (F) is the smallest halogen, it has the highest electron affinity and the strongest oxidizing power. As we move down the halogen group, the atomic size increases, and the electron affinity decreases, resulting in a decrease in their oxidizing power. Hence, the order of oxidizing power for the halogens is F₂ > Cl₂ > Br₂ > I₂.

Key concepts

Intermolecular Forces

When exploring the different states (solid, liquid, gas) of halogens at room temperature, it's crucial to understand how intermolecular forces function. Halogens exhibit London dispersion forces, a type of van der Waals force, which are weak interactions arising due to the temporary polarization of electron clouds around atoms or molecules. These forces are more pronounced in larger molecules with more electrons because there is a greater chance of instantaneous dipoles occurring.

As a result, heavier halogen molecules, like iodine (I₂), exhibit stronger dispersion forces leading to a solid state at room temperature. Conversely, lighter halogens like fluorine (F₂) and chlorine (Cl₂) are not able to maintain as much intermolecular cohesion and are found as gases under the same conditions. Bromine (Br₂) stands in between, with dispersion forces strong enough to exist as a liquid.

Electron Affinity

Electron affinity is a measure of the energy change when an electron is added to a neutral atom to form a negatively charged ion. Among the halogens, fluorine has the highest electron affinity due to its small size and high electronegativity, which allows it to exert a stronger pull on added electrons. This high electron affinity contributes to fluorine's strong oxidizing properties, as it readily accepts electrons from other species.

However, as we descend the halogen group, atomic size increases, and the added electrons are farther from the nucleus, experiencing less electrostatic pull from the positively charged center. Consequently, electron affinities decrease down the group, with iodine having the least tendency to gain an extra electron compared to other halogens.

Hydrogen Bonding

Significance of Hydrogen Bonds in HF

Hydrogen fluoride (HF) showcases the remarkable impact of hydrogen bonding on the physical properties of a compound. Hydrogen bonds form when hydrogen is covalently bonded to a highly electronegative atom such as fluorine, leading to a significant dipole. These bonds are a specific, strong type of dipole-dipole interaction.

Essentially, each HF molecule harbors a strong attraction between the slightly positive hydrogen atom and the slightly negative fluorine atom of another molecule. This intermolecular attraction leads to an unusually high boiling point for HF, far exceeding those of other hydrogen halides, which do not partake in hydrogen bonding to such an extent due to the lower electronegativity of their respective halogens.

Oxidizing Agents

Halogens are well-known oxidizing agents, with fluorine (F₂) being the most potent among them. An oxidizing agent is a substance that can accept electrons from another species during a chemical reaction, thereby oxidizing that species. Fluorine's high electron affinity and small atomic radius make it extremely effective at pulling electrons away from other substances.

Down the group, chlorine (Cl₂), bromine (Br₂), and iodine (I₂) exhibit progressively weaker oxidizing powers due to their larger atomic sizes and lower electron affinities. This trend explains the sequence of oxidizing strength of halogens being F₂ > Cl₂ > Br₂ > I₂. Such knowledge is invaluable in predicting and understanding the outcomes of redox reactions involving halogens.