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Potassium sulfite
Potassium sulfite
Potassium sulfite
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Potassium sulfite
Names
IUPAC name
Potassium sulfite
Other names
  • E225
Identifiers
3D model (JSmol)
ChemSpider
ECHA InfoCard 100.030.279 Edit this at Wikidata
UNII
  • InChI=1S/2K.H2O3S/c;;1-4(2)3/h;;(H2,1,2,3)/q2*+1;/p-2 ☒N
    Key: BHZRJJOHZFYXTO-UHFFFAOYSA-L ☒N
  • InChI=1/2K.H2O3S/c;;1-4(2)3/h;;(H2,1,2,3)/q2*+1;/p-2
    Key: BHZRJJOHZFYXTO-NUQVWONBAU
  • [O-]S(=O)[O-].[K+].[K+]
Properties
K2SO3
Molar mass 158.26 g/mol
Appearance White solid
Density 2.49 g/cm3[1]
Soluble
Acidity (pKa) 8
−64.0·10−6 cm3/mol
Hazards
Flash point Non-flammable
Related compounds
Other anions
 Potassium sulfate
Potassium selenite
Other cations
 Sodium sulfite
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
☒N verify (what is checkY☒N ?)

Potassium sulfite is the inorganic compound with the formula K2SO3. It is the salt of potassium cation and sulfite anion. It is a white solid that is highly soluble in water. Potassium sulfite is used for preserving food and beverages.[2]

History

[edit]

Potassium sulfite was first obtained by Georg Ernst Stahl in the early 18th century,[3] and was therefore known afterwards as Stahl's sulphureous salt. It became the first discovered sulfite and was first properly studied along with other sulfites by French chemists in the 1790s, and it was called sulphite of potash in the early 19th century.[4] Gilles-François Boulduc also discovered the salt in water of Passy in the 1720s.[5]

Production and reactions

[edit]
Main article: Sulfite § Reactions
See also: Wellman–Lord process

Potassium sulfite is produced by the thermal decomposition of potassium metabisulfite at 190 °C:[6]

K2S2O5 → K2SO3 + SO2

Structure

[edit]

The structure of solid K2SO3, as assessed by X-ray crystallography. The S-O distances are 1.515 Å, and the O-S-O angles are 105.2°[1]

References

[edit]
Revisions and contributorsEdit on WikipediaRead on Wikipedia

Potassium sulfite

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from Grokipedia
Potassium sulfite is an inorganic salt with the K₂SO₃, consisting of two cations and one anion, and a molecular weight of 158.26 g/mol. It appears as a white, odorless crystalline powder that is highly soluble in water (approximately 1060 g/L at 25°C) but only slightly soluble in , and it decomposes at temperatures above 590°C without a distinct . As a versatile and , potassium sulfite is widely employed in the as the additive E225 to inhibit , prevent oxidation, and maintain color and freshness in products such as wines, dried fruits, and processed vegetables. Its preservative properties stem from the release of in acidic conditions, which acts against microbes and enzymatic browning, though its use is regulated due to potential sensitivity in asthmatics. Beyond food applications, it serves as a developing agent in , a bleaching aid in textiles, and a component in to remove excess . Potassium sulfite can be prepared by reacting with or through the thermal decomposition of , yielding the compound alongside . Chemically, it exhibits strong reducing characteristics, oxidizing to in the presence of oxygen, and it reacts with acids to produce and the corresponding salt. While approved for use as a by regulatory bodies such as the FDA and EFSA, it may cause irritation to the eyes, skin, and upon direct exposure, and its solutions should be handled with care to avoid of fumes.

Properties

Physical properties

Potassium sulfite is a white crystalline powder or solid. Its is 158.26 g/mol. The of potassium sulfite is 2.35 g/cm³ at 20 °C. Potassium sulfite exhibits high solubility in , approximately 106 g per 100 mL at 25 °C, while it is slightly soluble in . It decomposes at temperatures above 590 °C without . As a hygroscopic substance, potassium sulfite readily absorbs from the air, potentially forming hydrates. Potassium sulfite is non-flammable and has no .

Chemical properties

Potassium sulfite is an ionic compound composed of cations (K⁺) and anions (SO₃²⁻). The anion acts as the conjugate base of hydrogen (HSO₃⁻), the second dissociation step of (H₂SO₃), which has a pKₐ of approximately 7.2; this renders aqueous solutions of potassium sulfite weakly basic. In , the compound undergoes partial , where SO₃²⁻ reacts with H₂O to form HSO₃⁻ and OH⁻ ions, typically resulting in a of around 8–9 for dilute solutions. Potassium sulfite exhibits moderate stability under ambient conditions but slowly oxidizes in the presence of air to (K₂SO₄), a process driven by the reducing nature of the ion. As a mild , the ion readily donates electrons in processes, contributing to its utility in preventing oxidation in various applications.

Structure

Molecular structure

Potassium sulfite, with the ionic formula , consists of cations and the anion (SO₃²⁻). The anion exhibits trigonal pyramidal , where a central atom is covalently bonded to three oxygen atoms, accompanied by a on the that causes deviation from ideal tetrahedral . In this , the S–O is approximately 1.515 Å, and the O–S–O bond angles are around 105.2°, reflecting the influence of the lone pair repelling the bonding pairs. The of the sulfite anion depicts the atom with three S–O bonds and one , often represented with to show equivalent bond lengths: one form includes a to one oxygen (formal charges: 0, double-bonded oxygen 0, each single-bonded oxygen –1), averaging to –1 per oxygen across hybrids. The ions (K⁺) do not form covalent bonds but coordinate electrostatically with the negatively charged oxygen atoms of the anions within the ionic lattice.

Crystal structure

Potassium sulfite (K₂SO₃) crystallizes in the trigonal crystal system with space group P̅3m1 (No. 164). The unit cell is hexagonal with lattice parameters a = 5.915 Å and c = 6.968 Å, corresponding to two formula units per unit cell (Z = 2). These parameters were determined through single-crystal X-ray diffraction studies, confirming the three-dimensional arrangement of the ions. In the crystal lattice, each potassium ion occupies sites with varying coordination geometries, typically surrounded by six to nine oxygen atoms from multiple sulfite (SO₃²⁻) anions. The sulfite anions adopt a pyramidal geometry, with the sulfur atom bonded to three oxygen atoms, and these units bridge the potassium cations to form a cohesive network. This ionic packing results in a density of approximately 2.49 g/cm³. The overall structure exhibits a layered character, akin to related sulfites, where alternating planes of cations and anions contribute to the stability of the solid phase. remains the primary method for elucidating these details, with computational optimizations providing refined bond lengths, such as K–O distances ranging from 2.65 to 3.24 and S–O distance of 1.54 .

Synthesis

Industrial production

Potassium sulfite is primarily produced industrially by absorbing gas into aqueous solutions of or , followed by and to obtain the solid product. The reaction with proceeds as 2KOH + SO₂ → K₂SO₃ + H₂O, an exothermic process requiring cooling to maintain control, typically conducted using high-purity KOH feedstock with low iron content (less than 0.3 ppm) to ensure clarity and quality. This method allows for the production of clear, high-purity solutions suitable for applications like photographic chemicals. With , the reaction is K₂CO₃ + SO₂ → K₂SO₃ + , where SO₂ is bubbled through the solution until evolution ceases, after which the mixture is concentrated. An alternative industrial route involves the of (K₂S₂O₅) at approximately 190°C, yielding potassium sulfite and gas via the reaction K₂S₂O₅ → K₂SO₃ + SO₂. Industrial-grade potassium sulfite typically achieves 95-98% purity, with efforts to minimize impurities such as sulfates through careful feedstock selection and process controls. It is manufactured on a commercial scale primarily for the and chemical industries.

Laboratory preparation

Potassium sulfite is commonly prepared in the laboratory by neutralizing with , where is formed from the dissolution of gas in . The balanced equation for the reaction is: 2\ceKOH+\ceH2SO3→\ceK2SO3+2\ceH2O2 \ce{KOH} + \ce{H2SO3} \rightarrow \ce{K2SO3} + 2 \ce{H2O} A standard procedure involves dissolving 100 g of in 200 mL of freshly boiled , and passing through it with a stream until the solution is just acidic to indicator. An equivalent amount of additional solution is then added, and the mixture is evaporated to boiling under the stream, followed by and of the crystals. This approach yields approximately 200 g of potassium sulfite. To minimize oxidation to during synthesis, the reaction is often conducted under an inert atmosphere, such as or gas. All preparations must be performed in a well-ventilated due to the potential evolution of toxic gas during gas introduction or if the solution becomes acidic.

Chemical reactions

Reactions with acids

Potassium sulfite undergoes acid-base reactions with strong acids, leading to the of the and subsequent decomposition to gas. For example, the reaction with proceeds as follows: K2SO3+2HCl→2KCl+H2O+SO2\mathrm{K_2SO_3 + 2HCl \rightarrow 2KCl + H_2O + SO_2} This evolution of SO₂ occurs readily upon addition of dilute strong acids, as the anion (SO₃²⁻) is protonated to form (H₂SO₃), which is unstable and decomposes. In reactions with strong acids like , full decomposition to SO₂ occurs: K2SO3+H2SO4→K2SO4+SO2+H2O\mathrm{K_2SO_3 + H_2SO_4 \rightarrow K_2SO_4 + SO_2 + H_2O} The ion exhibits stepwise in general, where the first yields the intermediate (HSO₃⁻) with weaker acids or under limited proton availability, before further acidification drives SO₂ release. can be formed by reacting with or absorbing SO₂ gas in a solution under controlled conditions. The release of SO₂ from potassium sulfite is highly -dependent, with rapid gas evolution occurring below 5 due to the shift in equilibrium toward undissociated . These reactions are kinetically fast at , typically completing within seconds to minutes depending on acid concentration and mixing.

reactions

Potassium sulfite acts as a in reactions, primarily through the oxidation of the sulfite ion (SO₃²⁻) to (SO₄²⁻). Under aerobic conditions, it undergoes slow oxidation by molecular oxygen, following the balanced : 2K2SO3+O2→2K2SO42\mathrm{K_2SO_3} + \mathrm{O_2} \rightarrow 2\mathrm{K_2SO_4} This reaction is kinetically controlled and occurs in aqueous solutions or air-exposed environments, converting the compound to potassium sulfate. The standard electrode potential for the SO₃²⁻/SO₄²⁻ couple in basic media, SO₄²⁻ + H₂O + 2e⁻ ⇌ SO₃²⁻ + 2OH⁻, is -0.93 V, indicating the strong reducing nature of sulfite relative to many common oxidants. Potassium sulfite also reacts with halogens such as chlorine in aqueous solution, where the sulfite ion is oxidized to sulfate while reducing Cl₂ to chloride. The reaction proceeds as: K2SO3+Cl2+H2O→K2SO4+2HCl\mathrm{K_2SO_3} + \mathrm{Cl_2} + \mathrm{H_2O} \rightarrow \mathrm{K_2SO_4} + 2\mathrm{HCl} This process is utilized in water treatment to neutralize excess chlorine, highlighting sulfite's role in halogen dechlorination. In photographic applications, potassium sulfite serves as a reducing agent in developing solutions, where it helps reduce silver halides (AgX) to metallic silver, contributing to image formation while preventing oxidation of the primary developing agents. The stability of potassium sulfite solutions against oxidation is influenced by environmental factors; the reaction rate accelerates with heat, light exposure, or trace transition metals (e.g., Cu²⁺, Fe²⁺), which catalyze the process via radical intermediates. Conversely, certain antioxidants, such as phenols, inhibit oxidation by scavenging reactive oxygen species or interrupting the radical chain mechanism.

Applications

Food and beverage preservation

Potassium sulfite, designated as E225, serves as a key in the and beverage industry, particularly in regions where it remains authorised, such as and , with maximum permitted levels varying by category, such as up to 3000 mg/kg expressed as SO₂ in dried fruits and 250 mg/L in low-sugar wines. While authorized in and , its direct use as a is prohibited in the and , though permitted in certain processed ingredients like production. The compound functions primarily as an through the release of (SO₂) in aqueous solutions, which effectively inhibits microbial growth by disrupting bacterial and fungal enzymes while also preventing enzymatic reactions that degrade . In , potassium sulfite is applied to prevent oxidation and maintain clarity and flavor stability during and storage. For dried fruits such as apricots and raisins, it controls mold proliferation and preserves color and texture by limiting oxidative damage. It is also incorporated into certain beverages, like fruit juices and , to extend and inhibit spoilage organisms. Compared to , potassium sulfite offers advantages including greater water solubility—approximately 106 g/100 mL at 25°C versus 28 g/100 mL for the sodium analog at 20°C—facilitating easier dissolution in formulations, and it provides a sodium-free alternative beneficial for low-sodium dietary needs. Detection of potassium sulfite in preserved products relies on analytical methods that quantify free or total SO₂ content, typically via followed by or chromatographic techniques, ensuring compliance with regulatory limits.

Industrial and other uses

Potassium sulfite serves as a versatile and in various industrial applications, leveraging its ability to react with oxidizing substances like and . In , it functions as an in systems, where it rapidly reacts with dissolved to form , thereby minimizing of metal components in high-pressure environments. This application is particularly valuable in processes, ensuring equipment longevity and operational efficiency. In the , potassium sulfite acts as both a bleaching agent and an antichlor during fabric processing. As a bleaching agent, it helps remove impurities and colorants from fibers like , enhancing whiteness without damaging the material. More critically, it neutralizes residual after bleaching, preventing degradation of cellulosic fibers and maintaining fabric strength. Its reducing properties make it preferable in eco-conscious and operations, where it stabilizes dyes and reduces environmental loads. The compound also plays a key role in paper manufacturing through the sulfite pulping process, where potassium sulfite is one of the salts employed to cook wood chips under acidic conditions, selectively dissolving while preserving fibers. This yields high-quality pulp suitable for fine papers, tissues, and writing materials, with the process allowing recovery of cooking liquors for chemical reuse. In , potassium sulfite is incorporated into developer formulations as a to inhibit aerial oxidation of developing agents and to acidify fixing baths, ensuring consistent image quality. In , potassium sulfite solutions are utilized as reagents for the qualitative detection and purification of aldehydes and ketones, forming reversible adducts that precipitate or allow separation from mixtures, due to equilibrium with ions in aqueous media. Beyond these, potassium sulfite finds use as a and in , where it prevents oxidation in formulations like hair straighteners and lotions. Its application in is limited due to potential sensitivity reactions.

History

Discovery

Potassium sulfite emerged from early 18th-century investigations into compounds and mineral waters, amid the development of pneumatic chemistry—the study of gases and their reactions. The term "Stahl's sulphureous salt" was used in 18th-century to refer to impure potassium sulfite. These findings contributed to broader pneumatic explorations of "fixed air" () and volatile gases, influencing subsequent work on acid salts and atmospheric reactions in the .

Early studies and nomenclature

In the 1790s, French chemists including and Louis-Bernard Guyton de Morveau conducted systematic analyses of potassium sulfite, classifying it as a neutral salt derived from the combination of sulfureous acid—a product of 's partial oxygenation—with , thereby establishing it as a sulfur-oxygen compound within emerging chemical frameworks. These studies built on experimental observations of 's combustion in oxygen, which produced acidic gases that reacted with bases like to form stable salts, advancing the understanding of oxygenation degrees in sulfur compounds. The for the compound evolved significantly during this period, transitioning from the pre-Lavoisierian term "Stahl's sulphureous salt"—honoring early 18th-century chemist —to "sulphite of " under the standardized system proposed by de Morveau, Lavoisier, Claude-Louis Berthollet, and Antoine François de Fourcroy in 1787. This emphasized the acid-base reaction, with "sulphite" denoting the anion from sulfureous acid and "potash" referring to the base derived from vegetable ashes. Following John Dalton's atomic theory in 1808, which introduced quantitative elemental proportions, the name was refined to "potassium sulfite" in early 19th-century literature, aligning with binary compound notations and facilitating precise chemical descriptions. Key publications in chemical treatises from 1810 to 1820 further elaborated on its properties. By the mid-19th century, investigations into sulfur compounds like sulfites transitioned toward industrial relevance, as their antimicrobial qualities—rooted in sulfur dioxide release—drew attention for wine and food preservation amid Europe's expanding beverage trade.

Safety and environmental considerations

Health hazards

Potassium sulfite can cause irritation to the skin, eyes, and respiratory tract upon contact or inhalation, with symptoms including redness, itching, and inflammation. The release of sulfur dioxide gas, particularly when the compound reacts with moisture or acids, exacerbates respiratory risks by further irritating the mucous membranes and potentially leading to coughing or shortness of breath. In cases of eye exposure, immediate flushing is necessary to prevent conjunctivitis or more severe damage. Sulfite sensitivity, a known reaction, is particularly prevalent among individuals with , where exposure to potassium sulfite can trigger , urticaria, or even in severe cases. This sensitivity affects approximately 5-13% of asthmatic patients, though it occurs in less than 2% of the general population. Symptoms may include swelling of the tongue, difficulty swallowing, , and rapid heartbeat, necessitating prompt medical intervention. Ingestion of potassium sulfite leads to gastrointestinal upset even at low doses, causing , , , and due to the formation of in the . The oral LD50 in rats is approximately 1,420 mg/kg, indicating moderate . Larger doses may result in effects, such as initial stimulation followed by depression or seizures. No specific OSHA PEL is established for potassium sulfite dust. Exposures should be controlled in accordance with general standards for particulates not otherwise regulated (PNOR), which set limits of 15 mg/m³ (total dust) and 5 mg/m³ (respirable fraction) as 8-hour time-weighted averages. for exposure involves rinsing affected skin or eyes with plenty of water for at least 15 minutes and removing contaminated clothing. For , move the individual to fresh air and provide oxygen if breathing is difficult; seek immediate attention in all cases of or persistent symptoms.

Environmental impact

Potassium sulfite poses a moderate risk to aquatic ecosystems primarily through its potential toxicity when released in elevated concentrations. Acute toxicity studies indicate an LC50 of 220–460 mg/L for fish (Leuciscus idus) after 96 hours of exposure. For aquatic invertebrates, the 48-hour EC50 for immobilization of Daphnia magna is 245.5 mg K2SO3/L, derived from sulfite anion data. Chronic effects include growth inhibition in algae, with a 72-hour EC10 of 55.3 mg K2SO3/L for Desmodesmus subspicatus. These values suggest it is acutely harmful to aquatic organisms but does not warrant classification as an environmental hazard under Regulation (EC) No 1272/2008, as thresholds exceed 1 mg/L. In environmental settings, potassium sulfite dissociates into potassium and ions, with the latter being unstable under aerobic conditions. anions rapidly oxidize to , integrating into the natural mediated by microorganisms that regulate states. This process limits long-term persistence, and is not applicable as an . The substance shows no bioaccumulation potential due to its ionic nature and high in water. High mobility in and increases the risk of dispersion, but proper wastewater management mitigates impacts from industrial or uses. Localized releases should be avoided to prevent oxygen depletion or shifts that could indirectly affect aquatic life.

References

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