Question

Explain the following observations: (a) for a given oxidation state, the acid strength of the oxyacid in aqueous solution decreases in the order chlorine \(>\) bromine \(>\) iodine. (b) Hydrofluoric acid cannot be stored in glass bottles. (c) Hi cannot be prepared by treating NaI with sulfuric acid. (d) The interhalogen \(\mathrm{ICl}_{3}\) is known, but \(\mathrm{BrCl}_{3}\) is not.

Short answer

The acid strength of oxyacids decreases from chlorine to iodine due to decreasing electronegativity, which affects O-H bond polarity and proton-donating ability. Hydrofluoric acid can't be stored in glass bottles because it reacts with silicon dioxide, corroding the glass. HI cannot be prepared by treating NaI with sulfuric acid, as the iodide ions get oxidized to elemental iodine instead. ICl3 exists but BrCl3 does not due to differences in electronegativity and bonding capacity between iodine and bromine, making iodine more likely to form hypervalent bonds.

Step-by-step solution

Observation (a): Acid strength of oxyacids in aqueous solution

The acid strength of oxyacids (a compound containing hydrogen, oxygen, and another element, such as HClOx) in aqueous solution decreases in the order chlorine > bromine > iodine. This is because the electronegativity of halogens decreases down the group (from chlorine to iodine). Higher electronegative elements attract oxygen's electrons more strongly, leading to a stronger O-H bond polarity. As a result, in oxyacids with higher electronegative central atoms, the O-H bond is more easily broken, and the acid can donate protons (H+) more easily, making it a stronger acid.

Observation (b): Hydrofluoric acid cannot be stored in glass bottles

Hydrofluoric acid (HF) cannot be stored in glass bottles because it reacts with the main component of glass, which is silicon dioxide (SiO2). The reaction between HF and SiO2 forms hexafluorosilicic acid (H2SiF6) and water: \[ SiO_2 + 6HF \to H_2SiF_6 + 2H_2O \] This reaction causes the glass to corrode and can lead to leakage or breakage of the glass container.

Observation (c): HI cannot be prepared by treating NaI with sulfuric acid

Hydrogen iodide (HI) cannot be prepared by treating sodium iodide (NaI) with sulfuric acid (H2SO4) because it would lead to the oxidation of iodide ions (I-) to elemental iodine (I2) by the highly oxidizing sulfuric acid: \[ 2I^{-} + H_2SO_4 \to I_2 + SO_2 + 2H_2O \] This side reaction prevents the formation of significant amounts of HI.

Observation (d): ICl3 is known, but BrCl3 is not

The interhalogen ICl3 exists, but BrCl3 does not. This is due to the difference in electronegativity and bonding capacity between iodine and bromine. Iodine is less electronegative and can form bonds with five or seven electron pairs (known as hypervalent bonds); hence, it can form compounds like ICl3. On the other hand, bromine is more electronegative and prefers to form only three electron-pair bonds. Since bromine has a lower capacity to form hypervalent bonds, BrCl3 does not exist.

Key concepts

Oxyacids

Oxyacids are acids that contain oxygen, hydrogen, and another element, usually a non-metal. The strength of an oxyacid in aqueous solution is influenced by the electronegativity of the central non-metal atom. As electronegativity increases, the ability to attract electrons from the O-H bond also increases. This makes it easier for the acid to donate a proton, thus increasing its acid strength.

In the case of halogen oxyacids, such as those with chlorine, bromine, and iodine, the acid strength varies due to differences in electronegativity:

• Chlorine has the highest electronegativity among these halogens, resulting in the strongest oxyacids.
• Bromine follows, with moderate electronegativity, resulting in weaker oxyacids compared to chlorine.
• Iodine, being the least electronegative, forms the weakest oxyacids.

This trend supports the observation that oxyacid strength decreases in the order: chlorine > bromine > iodine.

Interhalogen Compounds

Interhalogen compounds are molecules consisting of two or more different halogen atoms. These compounds can exhibit fascinating behaviors due to their variations in size and electronegativity.

For example,

• The compound ICl₃ is stable and well-known because iodine can accommodate more electron pairs in its outer shell, allowing the formation of hypervalent bonds.
• This is due to iodine's larger atomic size and ability to form five or seven electron-pair bonds.
• However, BrCl₃ does not exist because bromine prefers forming bonds with three electron pairs due to its higher electronegativity, preventing the stability required for such a compound.

The absence of BrCl₃ highlights the importance of considering both size and electronegativity when predicting the stability of interhalogen compounds.

Hydrogen Halides

Hydrogen halides are diatomic molecules consisting of hydrogen and a halogen. They are known for their acidic properties and vary greatly in strength and reactivity.

One common reaction involves sodium iodide (NaI) and sulfuric acid (H₂SO₄):

• When NaI is treated with sulfuric acid, the intended product, hydrogen iodide (HI), does not form in significant quantities.
• This is because sulfuric acid is a strong oxidizing agent and converts iodide ions (I^-) into elemental iodine (I₂), preventing HI production.
• This reaction creates a side product that competes with the desired reaction, showcasing the necessity of understanding reaction conditions in hydrogen halide chemistry.

Such interactions redefine how we approach creating these compounds to avoid unwanted side reactions.

Glass Corrosion by Acids

Glass corrosion occurs when certain acids react with the components of glass. Silicon dioxide (SiO₂), a primary constituent of glass, is particularly vulnerable to hydrofluoric acid (HF).

The reaction between HF and SiO₂ is

• SiO₂ reacts with HF, forming hexafluorosilicic acid (H₂SiF₆) and water.
• This leads to the degradation of the glass, making HF storage in glass containers unsafe as it results in corrosion and potential leakages.
• This reaction showcases the unique ability of HF to etch glass, a property not observed with other hydrohalic acids.

Understanding this reaction warns against using glass for HF storage, emphasizing the need for alternative materials.