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Bromous acid
View on Wikipediafrom Wikipedia
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| Names | |
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| IUPAC names
hydroxy-λ3-bromanone
hydroxidooxidobromine bromous acid | |
| Identifiers | |
3D model (JSmol)
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| ChEBI | |
| ChemSpider | |
PubChem CID
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CompTox Dashboard (EPA)
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| Properties | |
| HBrO2 | |
| Molar mass | 112.911 g/mol |
| Conjugate base | Bromite |
| Related compounds | |
Other anions
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Hydrobromic acid; hypobromous acid; bromic acid; perbromic acid |
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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Bromous acid is the inorganic compound with the formula of HBrO2. It is an unstable compound, although salts of its conjugate base – bromites – have been isolated. In acidic solution, bromites decompose to bromine.[1]
Discovery
[edit]In 1905, Richards A. H. proved the existence of bromous acid through a series of experiments involving silver nitrate (AgNO3) and bromine.[2] The reaction of excess cold aqueous to form hypobromous acid (HBrO), silver bromide (AgBr) and nitric acid (HNO3):
- Br2 + AgNO3 + H2O → HBrO + AgBr + HNO3
Richards discovered that the effect of adding excess liquid bromine in a concentrated silver nitrate (AgNO3) resulted in a different reaction mechanism. From numbers of equivalent portions of acid bromine formed from the previous reaction, the ratio between oxygen and bromine was calculated, with the exact value of O:Br (0.149975:0.3745), suggesting the acid compound contains two oxygen atom to one bromine atom. Thus, the chemical structure of the acid compound was deducted as HBrO2.[2]
According to Richards, hypobromous acid (HBrO) arises by the reaction of bromine and silver nitrate solution:[2]
- Br2 + AgNO3 + H2O → HBrO + AgBr + HNO3
- 2 AgNO3 + HBrO + Br2 + H2O → HBrO2 + 2 AgBr + 2 HNO3
Isomerism
[edit]The molecule HBrO2 has a bent structure with ∠(H−O−Br) angles of 106.1°. HOBrO also adopts a non-planar conformation with one isomer structure (2a) adopting a dihedral angle ∠(H−O−Br−O) of 74.2°. Moreover, the planar structures of two other isomers (2b-cis and 2c-trans) are transition state for fast enantiomerization.[3]
Another study identified three isomers: HOOBr, HOBrO, and HBr(O)O.[4]
Synthesis
[edit]A oxidation reaction between hypobromous acid (HBrO) and hypochlorous acid (HClO) can be used to produce bromous acid (HBrO2) and hydrochloric acid (HCl).[citation needed]
- HBrO + HClO → HBrO2 + HCl
A redox reaction of hypobromous acid (HBrO) can form bromous acid (HBrO2) as its product:[citation needed]
- HBrO + H2O − 2e− → HBrO2 + 2H+
The disproportionation reaction of two equivalents hypobromous acid (HBrO) results in the formation of both bromous acid (HBrO2) and hydrobromic acid (HBr):[citation needed]
- 2 HBrO → HBrO2 + HBr
A rearrangement reaction, which results from the syn-proportion of bromic acid (HBrO3) and hydrobromic acid (HBr) gives bromous acid (HBrO2):[citation needed]
- 2 HBrO3 + HBr → 3 HBrO2
Salts
[edit]
The salts NaBrO2·3H2O and Ba(BrO2)2·H2O have been crystallized. Upon treatment of these aqueous solutions with salts of Pb2+, Hg2+, and Ag+, the corresponding heavy metal bromites precipitate as solids.[1]
Belousov–Zhabotinsky reaction
[edit]Bromous acid is a product of the Belousov–Zhabotinsky reaction resulting from the combination of potassium bromate, cerium(IV) sulfate, propanedioic acid and citric acid in dilute sulfuric acid. Bromous acid is an intermediate stage of the reaction between bromate ion (BrO−
3 ) and bromine (Br−):[5][6]
- BrO−
3 + 2 Br− → HBrO2 + HBrO
Other relevant reactions in such oscillating reactions are:
- HBrO2 + BrO−
3 + H+ → 2 BrO•
2 + H2O - 2 HBrO2 → BrO−
3 + HOBr + H+
Bromites reduce permanganates to manganates (VI):[1]
- 2 MnO−
4 + BrO−
2 + OH− → 2 MnO2−
4 + BrO−
3 + H2O
pKa measurement
[edit]The acid dissociation constant of bromous acid, Ka = [H+][BrO−
2]/[HBrO2], was determined using different methods.
The value of the pKa for bromous acid was estimated in research studying the decomposition of bromites. The research measured the rate of bromite decomposition as a function of hydrogen and bromite ion concentrations. The experimental data of the log of the initial velocity were plotted against pH. Using this method, the estimated pKa value for bromous acid was 6.25.[7]
Using another method, the pKa for bromous acid was measured based on the initial velocity of the reaction between sodium bromites and potassium iodine in a pH range of 2.9–8.0, at 25 °C and ionic strength of 0.06 M. The first order dependence of the initial velocity of this disproportionation reaction on [H+] in a pH range of 4.5–8.0. The value of acid dissociation constant measured by this method is Ka = (3.7±0.9)×10−4 M and pKa = 3.43±0.05.[8]
Reactivity
[edit]In comparison to other oxygen-centered oxidants (hypohalites, anions of peroxides) and in line with its low basicity, bromite is a rather weak nucleophile.[9] Rate constants of bromite towards carbocations and acceptor-substituted olefins are by 1–3 orders of magnitude lower than the ones measured with hypobromite.
References
[edit]- ^ a b c Wiberg, Nils; Holleman, A. F.; Wiberg, Egon, eds. (2001). "Oxygen Acids of Bromine". Inorganic Chemistry. Academic Press. pp. 449–451. ISBN 978-0-12-352651-9.
- ^ a b c Richards, A. H. (15 January 1906). "The existence of bromous acid (HBrO
2)". Journal of the Society of Chemical Industry. 25 (1 Suppl): 4–5. hdl:2027/mdp.39015030318508. - ^ Glaser, Rainer; Jost, Mary (16 August 2012). "Disproportionation of Bromous Acid HOBrO by Direct O-Transfer and via Anhydrides O(BrO)2 and BrO-BrO2. An Ab Initio Study of the Mechanism of a Key Step of the Belousov–Zhabotinsky Oscillating Reaction". The Journal of Physical Chemistry A. 116 (32): 8352–8365. Bibcode:2012JPCA..116.8352G. doi:10.1021/jp301329g. PMID 22871057.
- ^ de Souza, Gabriel L. C.; Brown, Alex (July 2016). "The ground and excited states of HBrO2 [HOOBr, HOBrO, and HBr(O)O] and HBrO3 (HOOOBr and HOOBrO) isomers". Theoretical Chemistry Accounts. 135 (7). doi:10.1007/s00214-016-1931-8.
- ^ Vassalini, Irene; Alessandri, Ivano (30 December 2015). "Spatial and Temporal Control of Information Storage in Cellulose by Chemically Activated Oscillations". ACS Applied Materials & Interfaces. 7 (51): 28708–28713. doi:10.1021/acsami.5b11857. PMID 26654462.
- ^ Field, Richard J.; Koros, Endre; Noyes, Richard M. (December 1972). "Oscillations in chemical systems. II. Thorough analysis of temporal oscillation in the bromate-cerium-malonic acid system". Journal of the American Chemical Society. 94 (25): 8649–8664. Bibcode:1972JAChS..94.8649F. doi:10.1021/ja00780a001.
- ^ Massagli, A.; Indelli, A.; Pergola, F. (1970). "Kinetic investigation of the decomposition of bromite". Inorganica Chimica Acta. 4: 593–596. doi:10.1016/S0020-1693(00)93357-7.
- ^ Faria, R. B.; Epstein, Irving R.; Kustin, Kenneth (January 1994). "Kinetics of Disproportionation and pKa of Bromous Acid". The Journal of Physical Chemistry. 98 (4): 1363–1367. doi:10.1021/j100055a051.
- ^ Mayer, Robert J.; Ofial, Armin R. (18 May 2018). "Nucleophilic Reactivities of Bleach Reagents". Organic Letters. 20 (10): 2816–2820. doi:10.1021/acs.orglett.8b00645. PMID 29741385.
Further reading
[edit]- Ukeles, S.D.; Freiberg, M. (2002). "Bromine, Inorganic Compounds". Kirk-Othmer Encyclopedia of Chemical Technology. doi:10.1002/0471238961.0218151321110512.a01. ISBN 978-0-471-48494-3.
Bromous acid
View on Grokipediafrom Grokipedia
Chemical identity and structure
Molecular formula and nomenclature
Bromous acid is an inorganic compound with the molecular formula HBrO₂. Its molar mass is 112.911 g/mol, calculated from the atomic masses of hydrogen (1.008 g/mol), bromine (79.904 g/mol), and two oxygen atoms (15.999 g/mol each).[1][6] The accepted IUPAC name for HBrO₂ is bromous acid, reflecting its classification as a bromine oxoacid. In this nomenclature, the "-ous" suffix indicates the intermediate oxidation state of bromine relative to other oxoacids in the series. Bromine exhibits a +3 oxidation state in bromous acid, where the formal calculation assigns -2 to each oxygen, -1 to hydrogen, and thus +3 to bromine to balance the charge. The conjugate base of bromous acid is the bromite anion, BrO₂⁻, which forms upon deprotonation.[1][7][8][9] Bromous acid belongs to the family of bromine oxoacids, distinguished by the number of oxygen atoms and the corresponding oxidation state of bromine. These include hypobromous acid (HBrO, +1 oxidation state), bromic acid (HBrO₃, +5 oxidation state), and perbromic acid (HBrO₄, +7 oxidation state). This systematic naming follows conventions for halogen oxoacids, where "hypo-" denotes the lowest oxygen content, "-ous" the next, "-ic" the higher, and "per-" the highest.[8][10]Geometric structure
Bromous acid, in its conventional isomer form HOBrO, features a central bromine atom bonded to a hydroxy group (-OH) and a terminal oxo group (=O). This structure positions bromine as the central atom bonded to two oxygen atoms, with the geometry around bromine being bent (AX₂E₂ in VSEPR notation), consistent with its +3 oxidation state. Although computational studies indicate that the peroxy isomer HOOBr is slightly more stable, bromous acid conventionally refers to the HOBrO structure, where bromine exhibits the +3 oxidation state.[11][12] The molecular geometry is bent at the hydroxy oxygen, with the ∠H-O-Br bond angle calculated at 103.2° using density functional theory (B3LYP method).[11] This angular distortion reflects the electron domain geometry around bromine, which includes two bonding domains and two lone pairs, leading to a tetrahedral electron geometry and bent molecular shape according to VSEPR principles. The ∠O-Br-O bond angle at the central bromine is 112.0°, slightly wider than a typical bent angle due to the expanded octet and differing bond strengths.[11] Computational studies provide key bond length data: the single Br-O bond to the hydroxy oxygen measures approximately 1.839 Å, indicative of a sigma bond with partial pi character from resonance, while the shorter Br=O double bond is 1.616 Å, emphasizing higher bond order.[11] The O-H bond length is about 0.966 Å, typical for hydroxy groups in oxyacids.[11] The electronic structure involves bromine utilizing d-orbitals to accommodate an expanded octet of 10 valence electrons. In the Lewis representation, bromine forms a single bond to the hydroxy oxygen, a double bond to the terminal oxygen, and retains two lone pairs, resulting in a formal charge of zero on all atoms.[11] Resonance contributions, such as H-O-Br⁺=O ↔ H-O⁻-Br⁺=O, stabilize the structure but do not alter the primary bonding motif.[11] The overall conformation is non-planar, optimizing lone pair and bond repulsions.[11]Isomers
Bromous acid, with the molecular formula HBrO₂, exhibits three primary isomeric forms: the conventional HOBrO structure, the peroxy-like HOOBr, and the dioxo HBr(O)O. The HOBrO isomer features a bromine atom bonded to one hydroxyl group and one oxo group, representing the standard depiction of bromous acid. In contrast, HOOBr incorporates a peroxide linkage with bromine attached to the terminal oxygen, akin to perbromic structures in related systems, while HBr(O)O adopts a configuration with bromine doubly bonded to two oxygen atoms and protonated on one. These isomers arise from different arrangements of the atoms while maintaining the overall HBrO₂ composition, with tautomeric relationships possible between HOBrO and related forms under specific conditions.[12][13] Computational studies using ab initio methods, such as CCSD(T) with augmented correlation-consistent basis sets, have established the relative stabilities of these isomers. The HOOBr form is the global energy minimum, with HOBrO lying approximately 1–5 kcal/mol higher, indicating close energetic competition that may allow coexistence under low-temperature conditions. The HBr(O)O isomer is markedly less stable, exceeding the energy of HOOBr by roughly 50–60 kcal/mol, rendering it unlikely to persist without stabilization. These energy differences highlight the dominance of peroxide and mixed oxo-hydroxo structures over the fully dioxo form in HBrO₂. Early 20th-century chemical literature proposed potential isomerism based on analogy to chlorine analogs, but rigorous quantification emerged from 1990s theoretical investigations.[12][13] Indirect spectroscopic confirmation of the HOBrO and HOOBr isomers has been obtained through matrix-isolation infrared spectroscopy. In experiments involving vacuum ultraviolet photolysis of HBr/O₂ mixtures isolated in neon matrices at ~6 K, distinct IR bands attributable to both isomers were observed and assigned via isotopic substitution (e.g., O₂ and DBr). These vibrational frequencies, such as O–O stretches in HOOBr around 800–900 cm and Br–O modes in HOBrO near 700–800 cm, align closely with anharmonic predictions from density functional theory and coupled-cluster calculations, providing evidence for their structural integrity. The HBr(O)O isomer remains unobserved experimentally, consistent with its high energy.[14][13]Physical and chemical properties
Stability and decomposition
Bromous acid (HBrO₂) is highly unstable in aqueous solutions, undergoing rapid disproportionation to form hypobromous acid (HBrO) and bromic acid (HBrO₃). The primary decomposition pathway follows the second-order reaction 2 HBrO₂ → HBrO + HBrO₃, which proceeds through mechanisms involving direct oxygen transfer or intermediate anhydrides such as O(BrO)₂ and BrO–BrO₂.[15] This instability limits the direct observation of bromous acid, confining it primarily to transient species in chemical reactions.[16] The rate of decomposition exhibits strong dependence on concentration, pH, and temperature. In dilute solutions (0.01–0.02 M), the half-life is on the order of 0.1–0.2 seconds at temperatures between 15–38 °C, with the reaction following second-order kinetics enhanced by an acid-catalyzed term (rate = k[HBrO₂]² + k'[H⁺][HBrO₂]²).[16] Decomposition accelerates in acidic media due to the proton-dependent pathway, while stability improves in dilute alkaline conditions where the neutral form is deprotonated, suppressing the disproportionation. Elevated temperatures further shorten the half-life by lowering the activation energy barrier, estimated at 19–25 kJ/mol depending on the acid medium.[16][15] Ab initio calculations have elucidated the decomposition mechanisms, revealing low-energy paths via anhydride intermediates with a rate-limiting hydration step having an activation free energy of approximately 13.6 kcal/mol in aqueous solution.[15] Bromous acid cannot be isolated in concentrated forms and exists only as a short-lived intermediate, such as in the Belousov–Zhabotinsky oscillating reaction where it drives periodic redox cycles.[15]Acidity and pKa
Bromous acid (HBrO₂) is classified as a weak acid, undergoing partial dissociation in aqueous solution according to the equilibrium The acidity constant for this dissociation has been determined through indirect methods, as direct measurement is challenging due to the compound's instability. One key study utilized the kinetics of bromous acid disproportionation (2 HBrO₂ → HBrO + HBrO₃) to estimate the dissociation constant, yielding a pKₐ value of 3.43 ± 0.05 at 25°C and ionic strength 0.06 M.[3] Another approach involved analyzing the decomposition of bromite ions (BrO₂⁻), which provided an estimated pKₐ of 6.25, based on rate dependencies on pH and ion concentrations.[17] These measurements rely on equilibrium and kinetic modeling rather than direct potentiometry, with the lower pKₐ value from disproportionation kinetics being more widely adopted in subsequent research on halogen oxyacids.[3] The pKₐ of bromous acid indicates it is a moderately weak acid, weaker than its chlorine analog, chlorous acid (HClO₂), which has a pKₐ of 1.94. This difference arises primarily from chlorine's higher electronegativity compared to bromine (3.16 vs. 2.96 on the Pauling scale), which enhances the polarity of the O-H bond and stabilizes the conjugate base ClO₂⁻ more effectively than BrO₂⁻. Computational estimates using density functional theory have also supported pKₐ values around 3.4–3.5, aligning with experimental kinetic data and reinforcing the weak acid classification.Thermodynamic data
Bromous acid (HBrO₂) exhibits limited experimentally determined thermodynamic properties due to its instability, with most data derived from indirect measurements, kinetic studies, and computational estimates. The standard enthalpy of formation (Δ_f H°) for aqueous HBrO₂ is estimated at -33 kJ/mol, based on bond strength trends and comparisons with related oxyacids like hypobromous and bromic acids. Similarly, the standard Gibbs free energy of formation (Δ_f G°) is approximately -0.4 ± 1 kJ/mol, derived from equilibrium constants of disproportionation reactions and pK_a considerations in acidic media.[16] Entropy (S°) and heat capacity (C_p) data for HBrO₂ are scarce, primarily obtained from spectroscopic and quantum chemical calculations of vibrational frequencies. Computational studies using density functional theory (DFT) with basis sets like DN** predict entropy values around 200–250 J/mol·K at 298 K for gas-phase isomers, though aqueous adjustments are not well-established; heat capacities are estimated at 60–80 J/mol·K from mode contributions, but experimental validation is lacking.[18] In aqueous environments, bromous acid is highly soluble, forming dilute solutions during preparation, but its thermodynamic stability is low, leading to rapid decomposition via disproportionation.History and discovery
Initial evidence
In the mid-19th century, chemists observed that bromine dissolved in water undergoes disproportionation, forming a mixture of hydrobromic acid (HBr) and hypobromous acid (HOBr), but the reaction equilibria exhibited behaviors suggesting the transient presence of intermediate oxoacids with higher bromine oxidation states. These observations, noted in studies of bromine's reactivity in aqueous solutions, indicated that the system could not be fully explained by the known end products alone, as the rates of oxygen incorporation and product formation pointed to short-lived species bridging hypobromous and bromic acids. Further analytical evidence emerged from late 19th-century studies examining the oxygen-to-bromine (O:Br) ratios in products of bromine reactions with metal oxides and nitrates. These analyses revealed compositions with O:Br ratios between 1:1 (for hypobromites) and 3:1 (for bromates), consistent with an intermediate species at approximately 2:1, attributable to bromous acid or its salts. For instance, reactions of bromine with silver oxide yielded products whose elemental analysis suggested partial oxidation states corresponding to HBrO₂, providing quantitative support for its transient role in bromine chemistry.Confirmation and key studies
The existence of bromous acid was definitively confirmed in 1905 through the experimental work of A. H. Richards, who reacted bromine with a saturated solution of silver nitrate to form a precipitate of silver bromite. Analysis of the decomposition products from this precipitate established a 2:1 oxygen-to-bromine ratio, providing stoichiometric evidence for the HBrO₂ formula and marking the first direct verification of the acid's composition. In the mid-20th century, key studies focused on the role of bromous acid in oscillatory reactions, particularly the Belousov-Zhabotinsky (BZ) system, where kinetic modeling required its inclusion as a critical intermediate to explain observed oscillations. The seminal Field-Körös-Noyes mechanism, developed in 1972, incorporated bromous acid as the autocatalytic species in the oxybromine chemistry, addressing previous gaps in understanding the reaction's temporal behavior through detailed rate laws and species interconversions derived from experimental data. This framework not only validated indirect evidence for HBrO₂ from BZ kinetics but also guided subsequent refinements to the reaction model. Spectroscopic identification advanced in the 1970s and 1980s through UV-visible studies in acidic solutions, where bromous acid exhibited characteristic absorption maxima around 340 nm, allowing real-time monitoring during BZ oscillations and disproportionation reactions. These experiments, often using stopped-flow techniques, confirmed HBrO₂'s transient presence and reactivity, bridging early stoichiometric evidence with dynamic observations. Recent computational studies post-2000 have further validated the molecular structure of bromous acid using ab initio methods, such as density functional theory and coupled-cluster calculations, which predict a bent HO-BrO geometry with bond lengths and angles consistent with experimental inferences. These simulations not only corroborated the stability of the cis and trans isomers but also elucidated decomposition pathways, including direct O-transfer mechanisms in disproportionation, enhancing understanding of its fleeting nature in solution.[15]Synthesis and preparation
Laboratory methods
Bromous acid (HBrO₂) is prepared in the laboratory primarily as transient species in dilute aqueous solutions, given its inherent instability and tendency to decompose rapidly. These methods focus on generating the acid in situ for immediate use in further reactions or studies, under controlled conditions to suppress unwanted side reactions. A standard approach involves the chemical oxidation of hypobromous acid using hypochlorous acid, which proceeds quantitatively under mildly acidic conditions: This reaction is typically conducted at pH 4–6 to optimize yield and minimize further oxidation to bromate.[19] Another method utilizes the disproportionation of hypobromous acid, often induced by mild heating: The resulting mixture contains bromous acid alongside hydrobromic acid, necessitating separation techniques such as distillation or precipitation for isolation if required.[20] Bromous acid can also be synthesized via comproportionation of bromic acid and hydrobromic acid: This comproportionation reaction provides a direct route but is accompanied by competing reductions to bromine.[20] A further method generates bromous acid through the acid-catalyzed reduction of bromate ion by bromide ion in acidic media: This reaction is particularly useful for in situ production in kinetic studies and is controlled by pH and reactant ratios to favor bromous acid formation.[2] All laboratory preparations emphasize low temperatures (0–5 °C) and dilute concentrations (10⁻³ to 10⁻² M) to enhance stability during formation, as higher temperatures or concentrations accelerate decomposition pathways detailed elsewhere.[4]From bromine compounds
Bromous acid can be generated from elemental bromine or bromide ions through oxidation processes that form intermediate hypobromous acid or hypobromite, which then disproportionate to the target compound. Hypobromous acid is initially produced by the reaction of bromine with water in the presence of silver nitrate: Br₂ + AgNO₃ + H₂O → AgBr + HNO₃ + HBrO. This hypobromous acid undergoes disproportionation to bromous acid and hydrobromic acid via the reaction 2 HBrO → HBrO₂ + HBr, which is second-order in HOBr with a rate maximum at pH 3–8.[21][22] Hypobromite, prepared by reacting bromine with alkali hydroxide (Br₂ + 2 NaOH → NaBr + NaOBr + H₂O), can be disproportionated under controlled pH (10.5–12) and low temperature (≤0°C) to yield bromite and bromide ions, with subsequent acidification producing bromous acid. This method achieves 42–55% conversion to bromite at active bromine concentrations of 200–510 g/L.[23] Electrochemical routes involve anodic oxidation of bromide ions (Br⁻) in alkaline electrolytes at controlled potentials to form bromite (BrO₂⁻), followed by acidification to HBrO₂. Such methods are used in the industrial production of sodium bromite from bromide solutions.[24] Yields for these syntheses are typically low (often below 60%) due to the inherent instability of bromous acid, which decomposes rapidly to hypobromous acid and bromic acid or other products, complicating purification efforts.[23]Salts and derivatives
Bromite salts
The bromite ion, denoted as BrO₂⁻, is a monovalent inorganic anion and the conjugate base of bromous acid, formed through its dissociation in aqueous solution.[9] This ion exhibits a bent structure with the central bromine atom bonded to two oxygen atoms, featuring resonance between Br=O and Br-O⁻ configurations, which contributes to its relative stability compared to the unstable parent acid that readily decomposes via disproportionation.[25] Unlike bromous acid, the bromite ion persists in solution and forms stable salts under appropriate conditions. Common bromite salts include sodium bromite, typically isolated as the trihydrate NaBrO₂·3H₂O, which appears as yellow crystals and is highly soluble in water (up to 250 g/L at 20°C).[26][27] Barium bromite, Ba(BrO₂)₂, a white crystalline solid that also demonstrates solubility in water, though specific quantitative values are limited due to its relative instability.[28] These salts are generally obtained as hydrates, reflecting their affinity for water molecules in the solid state. Bromite salts are strong oxidizing agents and must be handled with care due to their potential for explosive decomposition when heated or contaminated with reducing materials.[29] They can disproportionate upon thermal decomposition, posing risks of releasing oxygen and bromine-containing byproducts.[28] Additionally, bromite ions may oxidize further to bromate (BrO₃⁻) in the presence of oxidants like ozone, contributing to environmental and health concerns in water treatment contexts.[30]Preparation and properties of salts
Bromite salts are primarily prepared by the controlled oxidation of bromine in alkaline media, where the reaction is maintained at low temperatures (10–20 °C) to favor bromite formation over bromate. In this process, bromine is added dropwise to an aqueous solution of sodium hydroxide, often in the presence of a small amount of sodium bromite as a carrier to stabilize the product, yielding a sodium bromite solution with concentrations up to 250 g/L after filtration.[27] Similar methods using NaOH and Br₂ produce stable trihydrate crystals of sodium bromite (NaBrO₂·3H₂O) upon concentration under reduced pressure and cooling, with the NaOH content adjusted to 0.5–5.0 wt.% to enhance stability.[31] Alternative laboratory preparations involve the thermal decomposition of bromate salts in the presence of bromide. For example, barium bromite (Ba(BrO₂)₂) is obtained by heating barium bromate under an oxygen atmosphere at 250 °C, while lithium bromite (LiBrO₂) forms by heating a mixture of lithium bromate and lithium bromide at 190–220 °C.[28] Although bromous acid (HBrO₂) can theoretically be neutralized with bases like NaOH to form bromite salts (e.g., NaBrO₂), the acid's instability limits this to in situ generation in solution rather than isolated preparation.[32] Bromite salts exhibit good solubility in water, which varies with the counterion; sodium bromite forms highly concentrated aqueous solutions (170–250 g/L), while barium and lithium bromites are also soluble but hygroscopic, allowing for crystallization under controlled conditions.[27][28] Analytical identification of bromite ions often relies on precipitation tests, where addition of lead(II), mercury(II), or silver(I) ions forms insoluble bromite salts such as Pb(BrO₂)₂, Hg(BrO₂)₂, or AgBrO₂, distinguishing them from other bromine oxyanions. Regarding thermal stability, bromite salts undergo disproportionation, typically decomposing to mixtures of bromate and bromide ions. In aqueous sodium bromite solutions, the primary decomposition pathway at 25 °C is 3 BrO₂⁻ → 2 BrO₃⁻ + Br⁻, with an activation energy of 26.8 kcal/mol and rate constant of 3.7 × 10⁻⁷ M⁻¹ s⁻¹, leading to slow degradation even at room temperature unless stabilized by dilute NaOH.[33] Solid salts show greater stability but decompose at elevated temperatures; lithium bromite melts and decomposes around 225 °C to LiBr and O₂, whereas barium bromite remains stable up to 795 °C before rapid decomposition to BaBr₂ and O₂.[28] Stabilized sodium bromite trihydrate crystals retain approximately 95% purity after 30 days at 30 °C, outperforming unstabilized forms by a factor of 5–7 in decomposition rate.[31]Reactivity and redox behavior
General reactivity
Bromous acid (HBrO₂) and its conjugate base, the bromite ion (BrO₂⁻), display relatively weak nucleophilic properties compared to the corresponding hypobromite species (BrO⁻). The nucleophilicity parameter for BrO₂⁻ in water is 12.75 with a sensitivity parameter , while for BrO⁻ it is 16.69 with , indicating that bromite engages in nucleophilic substitution reactions with electrophiles at rates several orders of magnitude slower than hypobromite.[34] This diminished nucleophilicity arises from the higher oxidation state of bromine in BrO₂⁻ (+3 versus +1), which reduces the availability of the lone pair on oxygen for donation. In comparison to chlorous acid (HClO₂) and chlorite (ClO₂⁻), bromous acid exhibits lower reactivity in analogous nucleophilic processes, primarily due to the larger atomic radius of bromine, which results in longer Br–O bonds, decreased polarity, and poorer orbital overlap for attack on substrates. Rate constants for bromite toward carbocations and electron-deficient alkenes are 1–3 orders of magnitude lower than those for chlorite, underscoring the trend of decreasing reactivity down the halogen group.[35] Bromous acid has potential as an intermediate in electrophilic bromination of organic compounds, where it can deliver bromine in the +3 oxidation state for selective substitution, though such applications are limited by its instability. Additionally, it undergoes reduction with hydrogen peroxide to yield hypobromous acid: This reaction highlights non-redox nucleophilic or substitution pathways involving the oxygens of HBrO₂.[19]Oxidation-reduction potentials
Bromous acid (HBrO₂) and the bromite ion (BrO₂⁻) participate in redox reactions, making them strong oxidizing agents in acidic conditions. The BrO₂⁻/BrO⁻ couple corresponds to the half-reaction This reflects the 2-electron reduction from the +3 oxidation state of bromine in bromite to the +1 state in hypobromite. The overall BrO₂⁻/Br⁻ couple involves a 4-electron transfer from +3 to -1 oxidation state. As an oxidizer, bromite can accept electrons from various reductants, while it can also serve as a reducer in reactions where it is oxidized to bromate (BrO₃⁻). This dual redox behavior underscores the compound's reactivity and tendency toward disproportionation. Direct measurement of standard potentials for these couples is challenging due to the instability of bromite, with values typically estimated from kinetic and equilibrium studies of related halogen systems. These potentials exhibit pH dependence, with values measured in acidic media (using H⁺) being higher than in basic media, where OH⁻ and H₂O participate instead. In neutral or basic conditions, the effective potential for analogous couples, such as BrO⁻ + H₂O + 2e⁻ → Br⁻ + 2OH⁻ (E° = 0.76 V in 1 M NaOH), is lower, reflecting the thermodynamic shift due to proton availability.[36] This pH sensitivity influences the stability and reactivity of bromous acid in different environments.Role in specific reactions
Belousov–Zhabotinsky reaction
The Belousov–Zhabotinsky (BZ) reaction exemplifies non-equilibrium chemical oscillations, where concentrations of reactive species fluctuate periodically, often visualized by color changes in the presence of a redox indicator. In its classical form, the reaction entails the catalyzed oxidation of malonic acid by bromate ions in strongly acidic sulfuric acid medium, with cerium ions (Ce³⁺/Ce⁴⁺) serving as the typical catalyst that alternates between oxidation states, driving the oscillatory behavior. This system, first reported in the 1950s and mechanistically elucidated in the 1970s, demonstrates how far-from-equilibrium conditions can yield temporal and spatial patterns akin to biological processes. Bromous acid (HBrO₂) emerges as a pivotal intermediate in the BZ reaction, functioning as the key autocatalytic agent that amplifies its own production during the oxidative phase. It forms initially through the reduction of bromate (BrO₃⁻) by bromide (Br⁻) or related species, such as the reaction BrO₃⁻ + Br⁻ + 2 H⁺ → HBrO₂ + HOBr, transitioning the system from bromide-dominated inhibition to autocatalysis once bromide levels fall below a critical threshold. In the Field–Körös–Noyes (FKN) mechanism, HBrO₂ participates in cycles that couple bromate reduction with catalyst oxidation, sustaining oscillations by intermittently regenerating bromide and shifting redox potentials to favor Ce⁴⁺ production.[37] Central to the mechanism are steps involving HBrO₂ disproportionation, which balances the autocatalytic buildup and contributes to phase transitions in the oscillation cycle. The primary autocatalytic step is the proton-assisted reaction of bromate with bromous acid:This is complemented by disproportionation:
where oxygen transfer from one HBrO₂ molecule to another facilitates the process, potentially via short-lived anhydride intermediates. An ab initio computational study has detailed these pathways, revealing that the symmetrical anhydride O(BrO)₂ serves as a low-energy transition structure in the direct O-transfer route, with activation barriers aligning closely with experimental kinetics in acidic conditions.[15]