Halides and oxygen-containing halogen compounds

Preparation and properties of metal halides

Halides are synthesized by two main routes: direct combination of simple substances (halogen and metal) or via double-displacement reactions in solutions.

The general rule is that metal chlorides, bromides, and iodides are highly soluble in water. However, there is a group of exceptions: salts of silver, monovalent copper, mono- and divalent mercury, as well as lead(II).

Interestingly, fluorine exhibits reverse properties compared to its group neighbors. Where chlorides and bromides precipitate (e.g., with a silver ion), silver fluoride turns out to be readily soluble. Conversely, soluble calcium chlorides and bromides are contrasted by practically insoluble calcium fluoride.

Due to lone electron pairs, halogen ions frequently act as ligands in complex compounds. Classic examples include the formation of potassium hexaferrate(III) or tetrachloroauric acid upon the interaction of simple salts with an excess of halide.

Oxygen-containing acids of chlorine

The properties of oxygen-containing chlorine compounds strictly follow patterns related to the valence of the central atom.

In the series from hypochlorous to perchloric acid, two opposing trends are observed:

1. Acid strength increases. $\text{HClO}$ is the weakest in the series, while $\text{HClO}_4$ is one of the strongest inorganic acids.
2. Oxidizing capacity and reaction rates decrease. Hypochlorous acid is a powerful and fast oxidizing agent (instantly displacing iodine from potassium iodide), whereas dilute perchloric acid practically does not react with the same reagent.

Hypochlorous acid and disproportionation reactions

When dissolved in water, chlorine forms hypochlorous acid ($\text{HClO}$). This compound is extremely unstable and its chemical transformations depend on conditions:

• In the light, the molecule decomposes, releasing active atomic oxygen. This very process accounts for the bactericidal action of chlorine.
• Upon water loss (if dehydrating agents such as $\text{CaCl}_2$ are used), the acid turns into its anhydride — chlorine(I) oxide.
• Upon heating, disproportionation occurs: hydrochloric and chloric acids are formed.

Such self-oxidation-self-reduction is also typical for the interaction of chlorine with alkalis. In a cold solution, chlorides and hypochlorites are formed (yielding Eau de Javelle), while upon heating, hypochlorites convert into chlorates.

Redox features

In solutions, halogens and their compounds behave predictably if their reducing capacity is evaluated. For instance, in a mixture of bromides and chlorides, a strong oxidizing agent (potassium permanganate) will react primarily with the bromide ion, as it is a stronger reducing agent than chloride.

It is impossible to chemically oxidize a fluoride ion in an aqueous solution. Fluorine itself is the strongest oxidizing agent, with its standard electrode potential being $+2.87$ V, which significantly exceeds the potential of permanganate.

The interaction of fluorine with oxygen gives rise to a unique compound — oxygen difluoride ($\text{OF}_2$). Since fluorine is more electronegative, it takes the electrons, acquiring an oxidation state of -1, while oxygen is forced to exhibit a positive oxidation state of +2, which is atypical for it. The molecule has polar covalent bonds and an angular geometry similar to the structure of water.

Class and exam questions

What else to study on this topic

• Oxidation of sulfites by halogens in an aqueous medium
• Mechanism of hydrogen ion binding during chlorine disproportionation in the presence of carbonates
• Sulfate method for producing hydrogen halides from solid salts
• Laboratory methods for synthesizing chlorine gas using manganese dioxide
• Comparison of the molecular geometry of oxygen difluoride and water

Other topics in the section "Chapter 9. Group VIIA Elements — Halogens"

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