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In chemistry, protonation (or hydronation) is the adding of a proton (or hydron, or hydrogen cation), usually denoted by H+, to an atom, molecule, or ion, forming a conjugate acid.[1] (The complementary process, when a proton is removed from a Brønsted–Lowry acid, is deprotonation.) Some examples include

Protonation is a fundamental chemical reaction and is a step in many stoichiometric and catalytic processes. Some ions and molecules can undergo more than one protonation and are labeled polybasic, which is true of many biological macromolecules. Protonation and deprotonation (removal of a proton) occur in most acid–base reactions; they are the core of most acid–base reaction theories. A Brønsted–Lowry acid is defined as a chemical substance that protonates another substance. Upon protonating a substrate, the mass and the charge of the species each increase by one unit, making it an essential step in certain analytical procedures such as electrospray mass spectrometry. Protonating or deprotonating a molecule or ion can change many other chemical properties, not just the charge and mass, for example solubility, hydrophilicity, reduction potential or oxidation potential, and optical properties can change.

Rates

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Protonations are often rapid, partly because of the high mobility of protons in many solvents. The rate of protonation is related to the acidity of the protonating species: protonation by weak acids is slower than protonation of the same base by strong acids. The rates of protonation and deprotonation can be especially slow when protonation induces significant structural changes.[2]

Enantioselective protonations are under kinetic control, are of considerable interest in organic synthesis. They are also relevant to various biological processes.[3]

Reversibility and catalysis

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Protonation is usually reversible, and the structure and bonding of the conjugate base are normally unchanged on protonation. In some cases, however, protonation induces isomerization, for example cis-alkenes can be converted to trans-alkenes using a catalytic amount of protonating agent. Many enzymes, such as the serine hydrolases, operate by mechanisms that involve reversible protonation of substrates.[citation needed]

See also

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References

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Protonation

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Protonation is the addition of a proton (H⁺) to an atom, molecule, or ion, forming a cationic species known as the conjugate . In the Brønsted-Lowry theory of and bases, this process occurs when a base accepts a proton from an , resulting in a proton transfer that alters the charge and electronic properties of the involved species. The extent of protonation depends on the of the environment and the pKa of the conjugate acid, with the protonated form predominating in acidic conditions ( < pKa) and the deprotonated form in basic conditions ( > pKa). This equilibrium is crucial for understanding acid-base reactivity and influences properties such as , charge distribution, and molecular interactions in aqueous solutions. In , protonation often serves as an initial step in reaction mechanisms, such as the acid-catalyzed hydration of alkenes or carbonyl compounds, where it increases the electrophilicity of the substrate to facilitate nucleophilic attack. In biochemistry, protonation states regulate biological processes, including , , and pH-responsive systems, where ionizable groups on polymers or biomolecules respond to cellular gradients for targeted release. Additionally, protonation plays a role in analytical techniques like , enabling the detection of molecular masses through the formation of protonated ions.

Fundamentals

Definition and Process

Protonation is the of a proton, denoted as H⁺ (with relative isotopic mass approximately equal to 1), to an atom, , or , thereby forming a conjugate . This is fundamentally a proton transfer reaction, where the proton acceptor—often referred to as a base—binds the H⁺ ion, resulting in a positively charged product. The basic notation for protonation is \ceB+H+−>BH+\ce{B + H^+ -> BH^+}, where B represents the base. In the protonation process, the proton is transferred from a proton donor, defined as a Brønsted-Lowry acid, to the proton acceptor or base, typically occurring in a . The general reaction can be expressed as \ceB+HA<=>BH++A−\ce{B + HA <=> BH^+ + A^-}, where HA is the acid and A⁻ is its conjugate base. This transfer is most commonly observed in solution, where solvent molecules stabilize the ions involved, but protonation also occurs in the gas phase without effects, allowing for the study of intrinsic molecular interactions via techniques like . The term protonation originated within the framework of acid-base theory developed independently in 1923 by Danish chemist and English chemist , who emphasized the role of proton transfer in defining acids and bases. Their theory marked a shift from earlier definitions by focusing on the dynamic exchange of protons rather than specific chemical compositions, laying the groundwork for understanding protonation as a core mechanism in chemical reactivity.

Conjugate Acid-Base Pairs

In Brønsted-Lowry acid-base theory, a conjugate acid-base pair consists of two that differ only by a single proton (H⁺). The base (B) accepts a proton to form its conjugate acid (BH⁺), while the acid (HA) donates a proton to form its conjugate base (A⁻). This relationship arises directly from proton transfer reactions, where the conjugate acid and base are interconverted by the gain or loss of H⁺. Protonation of a base to yield its conjugate introduces a positive charge to the , assuming the base was neutral, which enhances the electrophilicity of BH⁺ relative to B and shifts its role from a proton acceptor (basic) to a proton donor (). For example, (NH₃), a neutral base, protonates to form the (NH₄⁺), which bears a +1 charge and exhibits increased electrophilicity due to the electrostatic attraction for nucleophiles. This charge alteration fundamentally changes reactivity, making the conjugate more susceptible to nucleophilic attack at sites adjacent to the protonated center. The strengths of conjugate pairs are inversely related: a stronger forms a weaker conjugate base, and a stronger base forms a weaker conjugate , as measured by their respective acid dissociation constants (pKₐ values). In protonation reactions, the equilibrium position favors the side featuring the weaker and weaker base, ensuring that proton transfer proceeds until the relative strengths balance. This governs the extent of protonation in solution. A representative protonation reaction illustrating conjugate pair formation is that of ammonia with water: NH3+H2O⇌NH4++OH−\mathrm{NH_3 + H_2O \rightleftharpoons NH_4^+ + OH^-} Here, NH₃ acts as the base and H₂O as the acid, producing the conjugate acid NH₄⁺ and conjugate base OH⁻. This equilibrium derives from the general proton transfer process, where the base's lone pair on nitrogen coordinates with the proton from water's O-H bond, cleaving it heterolytically to distribute the charge. The position of equilibrium lies far to the left (K < 1) because water is a weaker acid than NH₄⁺ (pKₐ of NH₄⁺ ≈ 9.25), reflecting the strength relation of the pairs.

Mechanisms and Sites

General Mechanism

Protonation proceeds as a bimolecular reaction in which a proton, typically from a hydronium ion (\ceH3O+\ce{H3O+}) in aqueous solution, approaches an electron-rich site on a base, such as a lone pair on a heteroatom (e.g., nitrogen or oxygen) or a π-bond in an unsaturated system. The initial step involves the diffusion-controlled encounter of the proton donor and acceptor, followed by the formation of a partial bond in the transition state, where the proton is shared between the donor and acceptor. This leads to the protonated intermediate, denoted as \ceBH+\ce{BH+}, with the concomitant release of the conjugate base (e.g., \ceH2O\ce{H2O}). In gas phase, this process often lacks a significant barrier due to the absence of solvent stabilization, allowing exothermic protonation to occur rapidly upon collision. In aqueous environments, the transition state is characterized by distinct structural motifs: the Eigen mechanism describes a fast, diffusion-limited step forming a localized hydronium-like structure (\ceH9O4+\ce{H9O4+}), where the excess proton resides on a central \ceH3O+\ce{H3O+} solvated by three water molecules in the first shell. In contrast, the Zundel mechanism involves a delocalized proton in a symmetric \ceH5O2+\ce{H5O2+} complex, shared equally between two water molecules, which serves as a key intermediate or transition state for proton transfer. These forms interconvert rapidly (on femtosecond timescales), facilitating the overall protonation pathway, with the Zundel configuration often representing the lower-energy transition state for hopping between sites. Solvation shells play a critical role in stabilizing the transition state by providing hydrogen-bonding networks that delocalize charge and lower the activation barrier, particularly in the Eigen-to-Zundel conversion. In solution, the ordered water structure around the developing \ceBH+\ce{BH+} and departing \ceH2O\ce{H2O} contrasts with gas-phase protonation, where intrinsic molecular interactions dominate without such stabilization, often resulting in higher exothermicity but less selectivity in site preference. The rate law for protonation in diffusion-limited cases follows second-order kinetics: \rate=k[\ceB][\ceH+]\rate = k [\ce{B}][\ce{H+}] where kk approaches 101010^{10} to 101110^{11} M−1^{-1} s−1^{-1}, reflecting the encounter rate governed by diffusion coefficients. An energy diagram for the process illustrates a modest activation barrier at the transition state, modulated by solvation energy, beyond which the protonated product lies at lower energy.

Preferred Protonation Sites

The preferred sites of protonation in organic molecules are governed by electronic factors, primarily the distribution of electron density, which directs the incoming proton to regions of highest availability. Lone pairs on heteroatoms like nitrogen, oxygen, and sulfur offer localized high electron density, rendering these the most common protonation sites due to favorable orbital overlap with the proton's empty orbital. π-Electron clouds in alkenes and aromatic systems can also serve as sites, though protonation there typically requires stronger acids and leads to allylic or aromatic cations. The intrinsic basicity of these sites generally decreases in the order nitrogen > oxygen > carbon, a trend driven by differences in , , and s-character of the lone-pair orbitals, with nitrogen's lower electronegativity allowing greater electron donation compared to oxygen or carbon./23%253A_Amines/23.01%253A_Relative_Basicity_of_Amines_and_Other_Compounds) Representative examples illustrate these preferences. In aliphatic amines, the lone pair is protonated to form ions (R₃NH⁺), reflecting the strong basicity of amines with pK_b values around 3–5 in . For carbonyl groups in aldehydes and ketones, protonation targets the oxygen atom, yielding oxonium ions (R₂C=OH⁺) that are stabilized by between the positive oxygen and the adjacent carbon./16%253A_Chemistry_of_Benzene_-Electrophilic_Aromatic_Substitution/16.04%253A_Alkylation_and_Acylation-_The_Friedel-Crafts_Reaction) In contrast, carbon protonation is rare under standard conditions but occurs in media, such as the conversion of to the methonium ion (CH₅⁺) in fluorosulfonic acid-antimony pentafluoride mixtures, marking a pivotal observation in chemistry. Steric factors further modulate site selectivity by influencing accessibility. Bulky groups adjacent to a potential site can create hindrance, slowing or preventing protonation there and favoring alternative, less congested positions with comparable electron density. Resonance effects in conjugated systems can enhance stability at certain sites; in protonated aromatics, delocalization of the resulting positive charge across the ring lowers the energy barrier, as seen in the resonance structures of species like the pyridinium ion. A notable case is pyridine, where the sp²-hybridized nitrogen lone pair, orthogonal to the π-system, confers higher basicity than the ring carbons, with protonation exclusively at nitrogen being over 200 kJ/mol more favorable than at carbon, as determined by quantum chemical calculations.

Thermodynamics

Equilibrium Constants and pKa

The acid dissociation constant, KaK_a, quantifies the extent to which a weak acid \ceHA\ce{HA} dissociates in aqueous solution according to the equilibrium \ceHA⇌H++A−\ce{HA ⇌ H+ + A-}, expressed as Ka=[\ceH+][\ceA−][\ceHA]K_a = \frac{[\ce{H+}][\ce{A-}]}{[\ce{HA}]}. The pKa value is defined as pKa=−log⁡10Ka\mathrm{p}K_a = -\log_{10} K_a, providing a convenient logarithmic scale where lower pKa values indicate stronger acids. For bases \ceB\ce{B}, the base dissociation constant KbK_b describes \ceB+H2O⇌BH++OH−\ce{B + H2O ⇌ BH+ + OH-} as Kb=[\ceBH+][\ceOH−][\ceB]K_b = \frac{[\ce{BH+}][\ce{OH-}]}{[\ce{B}]}, and it relates to the pKa of its conjugate acid \ceBH+\ce{BH+} via pKb=14−pKa(\ceBH+)\mathrm{p}K_b = 14 - \mathrm{p}K_a(\ce{BH+}) at 25°C in water, since Ka⋅Kb=Kw=10−14K_a \cdot K_b = K_w = 10^{-14}. In protonation equilibria, such as \ceHA+B⇌BH++A−\ce{HA + B ⇌ BH+ + A-}, the equilibrium constant K=[\ceBH+][\ceA−][\ceHA][\ceB]=Ka(\ceHA)Ka(\ceBH+)=10pKa(\ceBH+)−pKa(\ceHA)K = \frac{[\ce{BH+}][\ce{A-}]}{[\ce{HA}][\ce{B}]} = \frac{K_a(\ce{HA})}{K_a(\ce{BH+})} = 10^{\mathrm{p}K_a(\ce{BH+}) - \mathrm{p}K_a(\ce{HA})}, meaning the position favors the side with the weaker (higher pKa). Thus, a strong (low pKa for HA) fully protonates a weak base (high pKa for BH+), as the conjugate base A- is a very weak base and BH+ is a weaker than HA, driving the reaction toward products. For the direct protonation \ceB+H+⇌BH+\ce{B + H+ ⇌ BH+}, the association constant is K=[\ceBH+][\ceB][\ceH+]=1Ka(\ceBH+)=10pKa(\ceBH+)K = \frac{[\ce{BH+}]}{[\ce{B}][\ce{H+}]} = \frac{1}{K_a(\ce{BH+})} = 10^{\mathrm{p}K_a(\ce{BH+})}, highlighting that bases with high conjugate pKa values are strongly protonated at low pH. pKa values are typically measured experimentally through , where the is plotted against added base to identify the midpoint ( = pKa for the buffer region), or via spectroscopic methods such as UV-Vis or NMR, which detect shifts in absorption or chemical shifts corresponding to protonated and deprotonated forms across a range. These techniques provide accurate thermodynamic data under controlled conditions, often at 25°C in aqueous media. Representative pKa values for common species illustrate the range of acid strengths relevant to protonation:
AcidpKaConjugate Base
\ceHCl\ce{HCl}-7\ceCl−\ce{Cl-}
\ceH3O+\ce{H3O+}-1.7\ceH2O\ce{H2O}
\ceCH3COOH\ce{CH3COOH}4.76\ceCH3COO−\ce{CH3COO-}
\ceNH4+\ce{NH4+}9.25\ceNH3\ce{NH3}
\ceH2O\ce{H2O}15.7\ceOH−\ce{OH-}
These values underscore how protonation is thermodynamically favored when the proton donor much lower pKa than the conjugate of the acceptor.

Solvent Effects on Protonation

Solvents play a crucial role in modulating protonation equilibria by influencing the of charged and neutral molecules, thereby altering the effective acidity and basicity of reactants. In protic solvents such as and , the presence of hydrogen-bond donor groups enables strong solvation of protons and anions through hydrogen bonding networks, which stabilizes ionic forms of protonated and shifts the pKa values toward lower (more acidic) equilibria compared to aprotic environments. For instance, computational studies show that the proton free energy in is approximately -266 kcal/mol, closely matched by at -266 kcal/mol, enhancing the stability of hydronium-like and favoring proton transfer. In contrast, aprotic solvents like and provide weaker solvation for ions due to the absence of hydrogen-bond donors, resulting in higher pKa values for protonated and a preference for neutral, unprotonated forms. This difference arises because aprotic solvents solvate cations less effectively, with proton solvation energies around -255 kcal/mol in , leading to less favorable protonation equilibria. The dielectric constant (ε) of a solvent further dictates the extent of ion stabilization and dissociation in protonation processes. Solvents with high dielectric constants, such as water (ε ≈ 80), effectively screen electrostatic interactions between charged species, promoting the dissociation of ion pairs and lowering the effective pKa for acids by reducing the energy barrier for charge separation. This effect is particularly pronounced for reactions involving charged protonated intermediates, where high-ε media (ε > 20–30) diminish Coulombic attractions, facilitating equilibria that favor ionic products. In low-dielectric solvents like nonpolar hydrocarbons (ε ≈ 2–4), the poor screening leads to tighter ion pairing, which suppresses dissociation and raises pKa values, making protonation less favorable for species that generate free charges. Protonation constants often vary linearly with the inverse of the dielectric constant in mixed solvent systems, underscoring the dominance of electrostatic contributions to solvation energies. Specific examples illustrate these solvent dependencies in protonation site preferences and equilibria. For ketones like acetone, protonation in yields a highly acidic conjugate with a pKa of approximately -7.2, reflecting strong of the resulting but still requiring highly acidic conditions due to the energy cost of charge localization on oxygen. In contrast, superacid media such as HF-SbF₅ (a with extremely low effective pKa < -20) enable protonation of ketones to form persistent carbocations by further destabilizing neutral species and stabilizing the delocalized positive charge through weak in the low-ε, noncoordinating environment. This shift allows observation of carbon-protonated sites, which are inaccessible in aqueous media. Additionally, in low-ε s, ion pairing between protonated bases and conjugate acids can alter reactivity by shielding charges and mimicking neutral behavior, as seen in equilibrium models for ionic complexation where paired species predominate over free ions. A notable consequence of solvent basicity is the leveling effect, observed in protic solvents like water, where all acids stronger than (pKa ≈ -1.7) appear equally strong because they fully protonate the solvent, preventing differentiation of their intrinsic strengths. This phenomenon limits the resolution of protonation hierarchies in aqueous systems but is mitigated in aprotic or less basic solvents, allowing measurement of superacidic behaviors. Overall, these solvent effects highlight the interplay between solvation and electrostatics in dictating protonation outcomes, with protic, high-ε media favoring ionic stabilization and aprotic, low-ε conditions preserving neutral preferences.

Kinetics

Rate Factors

The kinetics of protonation reactions are typically governed by a second-order rate law, expressed as rate = k [B][HA], where [B] is the concentration of the base (the species being protonated), [HA] is the concentration of the acid, and k is the rate constant. This reflects the bimolecular nature of the proton transfer step, in which the base and acid collide to form the protonated product. For reactions involving strong acids in aqueous solution, the process often reaches the diffusion limit, with rate constants on the order of 10¹⁰ M⁻¹ s⁻¹, constrained by the time required for molecular encounters in water. Acid concentration directly influences the rate, increasing it linearly as per the rate law, since higher [HA] provides more frequent collisions with the base. Temperature affects protonation rates through the Arrhenius equation, where the rate constant k = A e^{-E_a / RT}, and the activation energy E_a for proton transfers is notably low, typically ranging from 0 to 5 kcal/mol in water due to the minimal barriers in favorable acid-base pairings. This low E_a arises from the near-barrierless nature of many proton transfers when the pKa difference between the acid and conjugate acid of the base is small or favorable. Acid strength plays a key role, as stronger acids (those with lower pKa values) lead to faster protonation rates primarily through higher effective [H⁺] concentrations in solution, enhancing the driving force for the reaction. However, steric hindrance around the protonation site can impede access, slowing the rate by increasing the effective activation barrier; for instance, bulky substituents near basic sites reduce collision efficiency and favor alternative pathways. Protonation rates vary significantly with the basicity of the substrate, as seen in comparisons between simple alkenes and amines: simple alkenes, with much lower basicity (pKa of conjugate acid ≈ -25), undergo protonation more slowly than amines (pKa ≈ 10–11), often by orders of magnitude under similar conditions, due to less favorable energetics in forming the carbocation intermediate. Isotope effects further illuminate the mechanism, with deuteration (using D⁺ instead of H⁺) producing a primary kinetic isotope effect where k_H / k_D ranges from 2 to 7, reflecting the involvement of proton motion in the rate-determining step and differences in zero-point energies.

Reversibility and Catalysis

Protonation reactions are typically rapid and reversible under standard conditions, with the deprotonation step facilitated by the conjugate base of the proton donor, allowing the system to equilibrate quickly. The direction and extent of this reversibility are primarily governed by the (ΔG), which correlates directly with the difference in pKa values between the acid and the conjugate acid of the base being protonated; equilibria favor the side with the weaker acid, making reactions with ΔpKa > 10 effectively irreversible in practice, while smaller differences result in partial reversibility. Irreversibility can be achieved in specialized environments, such as (e.g., HF/SbF5 mixtures with H0 < -12), where extremely low basicity prevents effective by generating highly stable conjugate bases that do not revert under typical conditions. Alternatively, trapping agents can sequester the protonated species, shifting the equilibrium irreversibly; for instance, in superacid media, protonated carbonyl compounds or heterocycles form persistent carbocations without reversal due to the absence of viable nucleophiles. In reversible cases, the of the protonated intermediate is short, often on the order of milliseconds to seconds, calculated as t1/2=ln⁡2kdeprott_{1/2} = \frac{\ln 2}{k_{\text{deprot}}}, where kdeprotk_{\text{deprot}} is determined by the reverse rate constant and the concentration of the conjugate base. Protonation plays a central role in general , where a proton donor (HA) transfers H⁺ to the substrate in the , stabilizing developing negative charge or enhancing electrophilicity without full proton transfer to form a stable intermediate. This mechanism accelerates reactions by lowering the through partial protonation at the rate-determining step. A classic example is the acid-catalyzed of esters, such as with HCl, where protonation of the carbonyl oxygen facilitates nucleophilic attack by , leading to a tetrahedral intermediate and eventual cleavage to and alcohol. Mechanisms of in such processes are classified as A-1 (unimolecular, where protonation precedes a slow dissociation of the protonated to form an acylium ) or A-2 (bimolecular, involving concerted protonation and to form a tetrahedral intermediate), with A-2 predominating for most alkyl esters under acidic conditions due to the stability of the intermediate. The catalytic rate enhancement arises from the pre-equilibrium protonation, quantified by kcatkuncat=10ΔpKa\frac{k_{\text{cat}}}{k_{\text{uncat}}} = 10^{\Delta \text{p}K_a}, where ΔpKa is the difference between the pKa of and the pKa of the conjugate of the substrate, reflecting the for protonation that boosts the concentration of the reactive protonated form.

Applications

In Organic Chemistry

In organic chemistry, protonation plays a pivotal role in electrophilic addition reactions, where it activates π-bonds in alkenes and alkynes, rendering them susceptible to nucleophilic attack. For instance, in the acid-catalyzed hydration of alkenes, the alkene undergoes protonation by a strong acid such as H₂SO₄, forming a carbocation intermediate at the more substituted carbon in accordance with Markovnikov's rule; this intermediate then reacts with water to yield an alcohol after deprotonation. Similar protonation of alkynes facilitates additions like hydrohalogenation, generating vinyl carbocations that propagate further reactivity. Protonation is also essential in the deployment of protecting groups, enabling selective manipulation of functional groups during synthesis. For alcohols, acid-catalyzed protonation converts the hydroxyl into a good , facilitating formation of ethers such as tert-butyl or methoxymethyl (MOM) ethers, which shield the alcohol from unwanted reactions while allowing orthogonal deprotection later. In the case of carbonyl compounds, protonation of the oxygen atom generates resonance-stabilized oxonium ions, promoting of alcohols to form acetals or ketals as robust protecting groups for aldehydes and ketones. Superacid media, developed through pioneering work by George A. Olah, enable the direct observation and study of protonated species that are typically fleeting under standard conditions. In these extremely acidic environments, such as (HSbF₆), alkanes can be protonated to form alkyl cations, while aromatic compounds yield arenium ions (Wheland intermediates), providing insights into mechanisms of and rearrangements. Olah's contributions, recognized with the 1994 , revolutionized the understanding of reactive intermediates in organic transformations. A quintessential example of protonation's utility is the acid-catalyzed Fischer esterification, where carboxylic acids react with alcohols to form . The mechanism commences with protonation of the carbonyl oxygen by the acid catalyst (e.g., H₂SO₄), enhancing the electrophilicity of the carbon and generating a resonance-stabilized ; this is followed by nucleophilic attack from the alcohol on the protonated carbonyl, forming a tetrahedral intermediate. Subsequent proton transfers and elimination of water from the intermediate yield the ester product, with the catalyst regenerated. This process preferentially protonates the carbonyl oxygen due to its higher basicity compared to the hydroxyl oxygen.

In Biochemistry

In biochemistry, protonation plays a critical role in enzyme catalysis, particularly through proton transfer mediated by amino acid residues in active sites. Amino acids such as histidine (His), aspartate (Asp), and glutamate (Glu) often act as proton donors or acceptors due to their ionizable side chains with pKa values near physiological pH. A prominent example is the catalytic triad in serine proteases, consisting of Asp, His, and Ser residues, where the Asp-His hydrogen bond facilitates proton shuttling to activate the Ser nucleophile for peptide bond hydrolysis. In this mechanism, the histidine residue accepts a proton from the serine hydroxyl group, polarizing it for nucleophilic attack, while the aspartate stabilizes the protonated histidine through a low-barrier hydrogen bond, enabling efficient proton transfer without significant charge separation. Another key illustration is , a zinc-dependent that catalyzes the reversible hydration of CO₂ to and protons. In this process, a Zn²⁺-bound undergoes to form a nucleophilic hydroxide ion, which attacks CO₂; subsequent protonation steps regenerate the . Residues like Thr199 and Glu106 modulate the pKa of the Zn-bound (approximately 7.0), facilitating rapid proton transfer through a network of and achieving a rate acceleration of approximately 10⁷-fold compared to the uncatalyzed reaction. Protonation states of proteins are highly sensitive to , influencing folding, stability, and activity via the Henderson-Hasselbalch , which describes the ratio of protonated to deprotonated forms of ionizable groups: pH=pKa+log⁡10([A−][HA])\text{pH} = \text{p}K_\text{a} + \log_{10} \left( \frac{[\text{A}^-]}{[\text{HA}]} \right) where HA is the protonated acid form and A⁻ is the deprotonated base. In hydrophobic pockets, pKa values can shift significantly—e.g., Asp residues buried in nonpolar environments may exhibit elevated pKa (up to 8.5) due to desolvation penalties, stabilizing protonated states and promoting conformational changes like protein unfolding or lid opening in enzymes such as nitrophorin 4. These shifts alter electrostatic interactions, affecting oligomerization and catalytic efficiency. Proton gradients also drive essential transport and signaling processes. In mitochondria, the establishes a proton motive force across the inner membrane (ΔpH ≈ 0.5–1 unit, alkaline inside), where protons flow back through to power ATP synthesis, yielding up to 2.5 ATP per NADH oxidized. This chemiosmotic coupling, proposed by Mitchell, couples proton translocation to rotational catalysis in the F₀F₁-. In neuronal signaling, protonation modulates receptors; for instance, synaptic acidification releases protons from vesicles, activating acid-sensing ion channels () like ASIC1a, which permit Na⁺ and Ca²⁺ influx to induce excitatory postsynaptic currents and regulate in the .

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