Chlorite (ClO2–)
Table of Contents
The chlorite ion (ClO2–) consists of one chlorine atom bonded to two oxygen atoms and carries a single negative charge. In this ion, chlorine has an oxidation state of +3, placing chlorite between less oxidized and more highly oxidized chlorine–oxygen species such as:
- Hypochlorite: ClO– (oxidation state = +1)
- Chlorate: ClO3– (+5)
- Perchlorate: ClO4– (+7)
Across this series, the oxidation state of chlorine increases from +1 to +7 as the number of bonded oxygen atoms increases, which strongly influences their chemical behavior.
Chlorite is essential because it has moderate oxidizing and antimicrobial properties. Its most common salt, sodium chlorite (NaClO2), is widely used in all major cities to generate chlorine dioxide for drinking-water treatment. Chlorine dioxide removes harmful microorganisms and unpleasant odors without producing large amounts of chlorinated by-products.
Structure [3]
Molecular Geometry
The chlorite ion adopts a bent molecular geometry due to the presence of two lone pairs on the central chlorine atom, as shown in its Lewis structure. This results in a bond angle of approximately 111° between the two Cl–O bonds, consistent with valence shell electron pair repulsion (VSEPR) theory.
Resonance and Bonding
The negative charge in ClO2– is delocalized over the two oxygen atoms via resonance structures. As a result:
- Both Cl–O bonds have partial double-bond character in the resonance hybrid structure.
- The bond lengths are equivalent.
- Increases the stability relative to a structure with a localized charge.
This delocalization also contributes to the chlorite ion’s reactivity as an oxidant, because the π-bonding permits easier transfer of electron density during redox reactions.
Physical Properties of Chlorite Salts [4]
| Property | Description |
|---|---|
| Mass of the ClO2– ion | 67.45 g/mol |
| Physical State and Appearance | The chlorite ion exists only in solution and cannot be isolated. Its salts, such as sodium chlorite (NaClO2) and calcium chlorite (Ca(ClO2)2), are white to pale yellow crystalline solids. |
| Solubility in Water | — Sodium chlorite is highly soluble in water. — Calcium chlorite has limited stability and solubility. |
| Stability | — Sodium chlorite is stable under dry, neutral-to-alkaline conditions but decomposes in acids or on heating. — Calcium chlorite is much less stable and decomposes readily in water. |
| Moisture absorption | — Sodium chlorite is moderately moisture-absorbing. — Calcium chlorite is highly moisture-sensitive. |
Preparation of Chlorite Salts [5]
1. Industrial Method
Commercial sodium chlorite (NaClO2) is produced from sodium chlorate (NaClO3) through chlorine dioxide (ClO2) as a key intermediate. In an acidic solution, sodium chlorate is reduced by sulfur dioxide:
2 ClO3– + SO2 + 2 H+→ 2 ClO2 + SO42– + H2O
The chlorine dioxide is then absorbed into an alkaline solution, where it undergoes controlled disproportionation:
2 ClO2 + 2 OH– → ClO2– + ClO3– + H2O
This reaction forms chlorite ions (ClO2–) and chlorate ions (ClO3–). In the presence of sodium ions, these become sodium chlorite (NaClO2) and sodium chlorate (NaClO3). The chlorate produced is recycled back into the first step, making the process efficient and minimizing waste.
2. Laboratory Method
In laboratory chemistry, chlorite can be prepared by reducing sodium chlorate with hydrogen peroxide in an alkaline solution:
ClO3− + H2O2 → ClO2− + O2 + H2O
Although written as a simple equation, this reaction proceeds through chlorine dioxide intermediates. The peroxide acts as a selective reducing agent that stops the reduction at the chlorite level rather than proceeding all the way to chloride.
Chemical Behavior [4,6]
1. Oxidizing Ability
Chlorite is a moderate oxidizing agent. It is stronger than hypochlorite but weaker than chlorate or perchlorate, a trend that mirrors the increasing oxidation states of chlorine. One illustrative redox reaction is the oxidation of iodide:
ClO2– + 4 I– + 4 H+ → Cl– + 2 I2 + 2 H2O
This reaction is used in chemistry to study how fast the reaction occurs and how strongly chlorite acts as an oxidizing agent, showing how it transfers oxygen to other substances.
2. Disproportionation
Because chlorite contains chlorine in an intermediate oxidation state, it can participate in redox reactions, including disproportionation:
3 ClO2– → 2 ClO3– + Cl–
Here, some chlorine atoms are oxidized from +3 to +5 (chlorate), while others are reduced from +3 to −1 (chloride). This tendency to undergo self-oxidation and self-reduction makes chlorite chemically unstable under certain conditions, which is why chlorite-containing industrial formulations must be carefully stabilized.
Health, Safety, and Handling
Chlorite compounds are strong oxidants and can react violently with acids, organic materials, or reducing agents. Contact with acids releases chlorine dioxide, a toxic and explosive gas. For this reason, sodium chlorite must be stored in cool, dry, alkaline conditions and handled using appropriate protective equipment. [7,8]
