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Deprotonation of acetic acid by a hydroxide ion

Deprotonation (or dehydronation) is the removal (transfer) of a proton (or hydron, or hydrogen cation), (H+) from a Brønsted–Lowry acid in an acid–base reaction.[1][2] The species formed is the conjugate base of that acid. The complementary process, when a proton is added (transferred) to a Brønsted–Lowry base, is protonation (or hydronation). The species formed is the conjugate acid of that base.

A species that can either accept or donate a proton is referred to as amphiprotic. An example is the H2O (water) molecule, which can gain a proton to form the hydronium ion, H3O+, or lose a proton, leaving the hydroxide ion, OH−.

The relative ability of a molecule to give up a proton is measured by its pKa value. A low pKa value indicates that the compound is acidic and will easily give up its proton to a base. The pKa of a compound is determined by many aspects, but the most significant is the stability of the conjugate base. This is primarily determined by the ability (or inability) of the conjugated base to stabilize negative charge. One of the most important ways of assessing a conjugate base's ability to distribute negative charge is using resonance. Electron withdrawing groups (which can stabilize the molecule by increasing charge distribution) or electron donating groups (which destabilize by decreasing charge distribution) present on a molecule also determine its pKa. The solvent used can also assist in the stabilization of the negative charge on a conjugated base.

Bases used to deprotonate depend on the pKa of the compound. When the compound is not particularly acidic, and, as such, the molecule does not give up its proton easily, a base stronger than the commonly known hydroxides is required. Hydrides are one of the many types of powerful deprotonating agents. Common hydrides used are sodium hydride and potassium hydride. The hydride forms hydrogen gas with the liberated proton from the other molecule. The hydrogen is dangerous and could ignite with the oxygen in the air, so the chemical procedure should be done in an inert atmosphere (e.g., nitrogen).

Deprotonation can be an important step in a chemical reaction. Acid–base reactions typically occur faster than any other step which may determine the product of a reaction. The conjugate base is more electron-rich than the molecule which can alter the reactivity of the molecule. For example, deprotonation of an alcohol forms the negatively charged alkoxide, which is a much stronger nucleophile.

To determine whether or not a given base will be sufficient to deprotonate a specific acid, compare the conjugate base with the original base. A conjugate base is formed when the acid is deprotonated by the base. In the image above, hydroxide acts as a base to deprotonate the carboxylic acid. The conjugate base is the carboxylate salt. In this case, hydroxide is a strong enough base to deprotonate the carboxylic acid because the conjugate base is more stable than the base because the negative charge is delocalized over two electronegative atoms compared to one. Using pKa values, the carboxylic acid is approximately 4 and the conjugate acid, water, is 15.7. Because acids with higher pKa values are less likely to donate their protons, the equilibrium will favor their formation. Therefore, the side of the equation with water will be formed preferentially. If, for example, water, instead of hydroxide, was used to deprotonate the carboxylic acid, the equilibrium would not favor the formation of the carboxylate salt. This is because the conjugate acid, hydronium, has a pKa of -1.74, which is lower than the carboxylic acid. In this case, equilibrium would favor the carboxylic acid.

References

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Deprotonation

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Deprotonation is the chemical process by which a proton (H⁺) is removed from a , typically from an by a base, yielding the conjugate base of the and the conjugate of the base. This proton transfer is the core of Brønsted-Lowry acid-base theory, which defines as proton donors and bases as proton acceptors, extending the concept beyond aqueous solutions to any or even gas phase. The reaction is reversible, with the equilibrium favoring the side of the weaker and weaker base. The extent of deprotonation depends on the acid's strength, measured by its pKa value—the lower the pKa, the more readily the proton is donated. For instance, strong acids like (HCl, pKa ≈ -7) fully deprotonate in water to form chloride ions (Cl⁻) and (H₃O⁺), while weaker acids like acetic acid (CH₃COOH, pKa ≈ 4.76) partially deprotonate, establishing an equilibrium with (CH₃COO⁻). In , deprotonation generates reactive intermediates such as enolates from carbonyl compounds. Deprotonation is equally vital in biochemistry, where it facilitates enzymatic mechanisms and proton transport. residues like (pKa ≈ 6.0) undergo rapid deprotonation/ to shuttle protons in , as seen in serine proteases. In cellular environments, pH-dependent deprotonation influences protein stability and folding, as well as proton gradients in ATP synthases that drive energy production.

Fundamentals

Definition

Deprotonation is the chemical process involving the removal of a proton (H⁺) from a Brønsted-Lowry acid, thereby forming its conjugate base in an acid-base reaction. This process is fundamental to acid dissociation, represented by the general equilibrium equation: HA⇌H++A−\text{HA} \rightleftharpoons \text{H}^+ + \text{A}^- where HA denotes the acid and A⁻ its conjugate base. The reverse process, known as , occurs when a Brønsted-Lowry base accepts a proton to form its conjugate . In this framework, and bases are defined by their roles in proton transfer: a Brønsted-Lowry is a proton donor, while a Brønsted-Lowry base is a proton acceptor. Certain chemical species, termed amphiprotic, possess the ability to act as both acids and bases, undergoing either deprotonation or depending on the reaction conditions. For instance, (H₂O) exemplifies an amphiprotic species, as it can donate a proton to form the hydroxide (OH⁻) or accept one to form the (H₃O⁺). The tendency for deprotonation in such systems is often assessed using the pKₐ value, which indicates the acidity strength of the species. The pK_a is defined as pK_a = −log_{10} K_a, where K_a is the measuring the equilibrium extent of deprotonation.

Relation to Acid-Base Theories

Deprotonation is fundamentally described within the Brønsted-Lowry theory of acids and bases, which defines acids as proton donors and bases as proton acceptors, thereby framing deprotonation as the essential proton transfer process in acid-base reactions. This theory, developed independently in 1923 by Danish chemist and English chemist , expanded the understanding of acid-base behavior beyond aqueous environments by emphasizing proton exchange rather than ion production. In this framework, deprotonation represents the core mechanism by which an acid loses a proton to form its conjugate base, establishing a reversible equilibrium that underpins the theory's predictive power for reaction tendencies. In contrast to the Brønsted-Lowry emphasis on proton transfer, the Lewis of acids and bases, proposed by in , views acid-base interactions more broadly as the donation and acceptance of s, where deprotonation involves the Lewis base donating an electron pair to the proton (a Lewis acid) detached from the Brønsted acid, forming a new bond while the conjugate base is released, though the proton transfer aspect remains secondary to the electron-sharing perspective. This Lewis approach encompasses Brønsted-Lowry reactions as a but extends to non-protonic processes, highlighting that while deprotonation inherently involves proton movement, its theoretical interpretation in Lewis terms prioritizes the electronic reorganization over the proton itself. The distinction underscores how Brønsted-Lowry provides a more specific lens for deprotonation events, particularly in contexts where proton mobility is key. The Arrhenius theory, introduced by in 1884, offers a more limited view of deprotonation, restricting it to aqueous solutions where acids dissociate to produce hydrogen ions (H⁺) and bases yield hydroxide ions (OH⁻), thus deprotonation equates to the generation of these ions in water. This model, while foundational, confines deprotonation explanations to hydration-dependent ion formation and fails to account for reactions in non-aqueous media or those without explicit H⁺/OH⁻ production, rendering it less versatile than subsequent theories for broader theoretical context.

Mechanism and Kinetics

General Mechanism

Deprotonation occurs through a proton transfer process in which a base (B:) abstracts a proton (H⁺) from an acid (HA), yielding the conjugate acid (BH⁺) and conjugate base (A⁻), as defined within the Brønsted-Lowry framework of acid-base reactions. This stepwise mechanism typically begins with the formation of a hydrogen bond between the base's lone pair and the acidic proton, facilitating the subsequent cleavage of the E-H bond in HA (where E is the atom bearing the proton) and the formation of a new bond to B, resulting in rapid proton relocation. The process is inherently equilibrium-driven, where the forward deprotonation competes with the reverse , governed by the K=[A−][BH+][HA][B]K = \frac{[A^-][BH^+]}{[HA][B]}. This constant reflects the relative strengths of HA and BH⁺ as acids, with larger K favoring deprotonation and formation of dissociated species. plays a pivotal role in modulating this equilibrium; protic solvents, such as or alcohols, stabilize the charged products BH⁺ and A⁻ through extensive hydrogen bonding and shells, promoting . In contrast, aprotic solvents like provide weaker stabilization of the charged products, shifting the equilibrium toward the undissociated form. The energy profile of deprotonation features a characteristically low for the proton transfer step, frequently on the order of a few kcal/mol, owing to the stabilizing influence of bonding that pre-aligns the reactants and reduces the barrier to proton abstraction.

Factors Influencing Deprotonation

The ease of deprotonation is fundamentally governed by the intrinsic acidity of the proton donor, which depends on the strength of the E-H bond (where E is an electronegative atom, such as oxygen, , or carbon in certain contexts) and the of the atom bearing the proton. Weaker E-H bonds facilitate proton removal, as less energy is required to break them, thereby increasing the acidity and promoting deprotonation. Higher of E enhances acidity by better stabilizing the resulting negative charge on the conjugate base through inductive withdrawal of . For instance, the O-H bond in alcohols is more acidic than the C-H bond in alkanes due to oxygen's greater . Stability of the conjugate base plays a crucial role in determining the feasibility of deprotonation, as more stable anions lower the energy barrier for proton loss. delocalization of the negative charge across multiple atoms significantly stabilizes the conjugate base, making deprotonation more favorable; for example, in ions, the charge is spread over two oxygen atoms. Inductive effects from nearby electron-withdrawing groups further enhance this stability by polarizing the E-H bond and dispersing the charge. in protic solvents, such as , also influences conjugate base stability, with smaller, more concentrated anions experiencing stronger hydrogen bonding interactions that stabilize them relative to larger ones. The strength of the base used in the deprotonation reaction directly impacts its effectiveness, particularly for weaker acids that require a sufficiently strong proton acceptor to drive the process forward. A base must be stronger than the conjugate base of the to achieve complete deprotonation, ensuring the equilibrium favors the products. For example, (NaH), a very strong, , is commonly employed to deprotonate weakly acidic compounds like alcohols or terminal alkynes by generating the anion (H⁻), which readily abstracts the proton. Environmental conditions further modulate the rate and extent of deprotonation. Elevated temperatures accelerate the by increasing molecular , allowing more collisions to surpass the barrier for proton transfer. In aqueous media, higher promotes deprotonation by elevating the concentration of available bases like ions, shifting the process toward the deprotonated form. For reactions involving highly reactive bases such as NaH, an inert atmosphere (e.g., or ) is essential to prevent side reactions, including hydrogen gas evolution upon exposure to or oxygen, which could lead to ignition or incomplete deprotonation.

Examples

In Organic Chemistry

In , deprotonation plays a pivotal role in synthetic transformations by generating nucleophilic species from carbon-based acids, enabling carbon-carbon bond formation and manipulation. Terminal alkynes, with their acidic C-H bonds due to sp-hybridization, undergo deprotonation using strong bases such as (NaNH₂) in liquid ammonia to form acetylide ions (RC≡C⁻). These acetylide anions serve as potent nucleophiles for addition to carbonyl compounds like aldehydes and ketones, yielding propargyl alcohols that are valuable intermediates in the synthesis of complex molecules. The reaction exemplifies the utility of deprotonation in extending carbon chains, as illustrated by the equation: RC≡CH+NaNH2→RC≡CNa+NH3\text{RC}\equiv\text{CH} + \text{NaNH}_2 \rightarrow \text{RC}\equiv\text{CNa} + \text{NH}_3 Carbonyl compounds, particularly ketones and aldehydes, are deprotonated at the α-position to form s, which are key to reactions like the . (LDA), a bulky , is commonly employed for this purpose in (THF) at low temperatures, ensuring selective kinetic enolate formation without self-condensation. This deprotonation facilitates nucleophilic attack on another carbonyl, building β-hydroxy carbonyl products essential for synthesis and pharmaceutical intermediates. Deprotonation of carboxylic acids with aqueous (NaOH) generates water-soluble salts (RCOO⁻ Na⁺), which enhance the purification of organic mixtures through acid-base extractions. This process converts the neutral acid into an ionic , improving in aqueous phases and allowing separation from non-acidic impurities, followed by acidification to regenerate the acid. The equation for this transformation is: RCOOH+NaOH→RCOO−Na++H2O\text{RCOOH} + \text{NaOH} \rightarrow \text{RCOO}^- \text{Na}^+ + \text{H}_2\text{O} Such salts are routinely used in salt formation to stabilize and handle reactive carboxylic derivatives in synthesis. The enhanced nucleophilicity of these deprotonated underpins their broader reactivity in subsequent transformations.

In Inorganic Chemistry

In inorganic chemistry, deprotonation often occurs in ionic systems involving metal hydrides, which serve as strong bases due to the high basicity of the hydride ion (H⁻). For instance, (NaH) reacts with protic compounds such as or alcohols by abstracting a proton, yielding gas (H₂) and the corresponding conjugate base. The reaction is typically represented as NaH + RH → NaR + H₂, where RH is the proton donor; this process is exothermic and drives the formation of H₂, making metal hydrides valuable for generating conditions in synthetic applications. Polyprotic acids, such as (H₂SO₄), undergo stepwise deprotonation in aqueous media, reflecting the decreasing acidity of successive protons. The first deprotonation is complete, producing the bisulfate ion (HSO₄⁻) with a pKₐ₁ ≈ -3, while the second step to (SO₄²⁻) is partial with pKₐ₂ = 1.99, establishing an equilibrium. A representative neutralization reaction for the initial step is H₂SO₄ + OH⁻ → HSO₄⁻ + H₂O, which occurs quantitatively in basic conditions and is fundamental to acid-base titrations and industrial processes like fertilizer production. In coordination chemistry, deprotonation transforms aqua ligands in metal complexes to hydroxo ligands, particularly in aqueous solutions where the metal ion's influences ligand acidity. For hexaaqua complexes [M(H₂O)₆]^(m+), addition of initiates sequential deprotonation: [M(H₂O)₆]^(m+) + OH⁻ → [M(H₂O)₅(OH)]^((m-1)+) + H₂O, often leading to precipitation of metal hydroxides or formation of polynuclear species for divalent metals like Co²⁺ or Zn²⁺. This process is pH-dependent and irreversible under typical conditions, playing a key role in the of metal ions in natural waters and synthetic design.

Applications and Measurement

Practical Applications

In the industrial production of soap, deprotonation plays a central role through the process, where fats or oils—primarily triglycerides—are hydrolyzed using (NaOH). This base-catalyzed reaction cleaves the ester bonds in the triglycerides, yielding and fatty acid carboxylates, which are the deprotonated forms of s that confer soap's properties. The deprotonation step ensures the formation of water-soluble sodium salts rather than insoluble carboxylic acids, enabling the emulsification of oils in for cleaning applications. This process has been utilized for centuries and remains a cornerstone of the soap industry, producing approximately 20 million tons annually as of 2024 for household and industrial use. In pharmaceutical formulations, deprotonation enhances drug solubility and bioavailability, particularly for weakly acidic compounds like aspirin (acetylsalicylic acid). Aspirin itself has limited water solubility due to its neutral carboxylic acid form, but conversion to its sodium salt via deprotonation with a base such as sodium bicarbonate or NaOH creates a charged carboxylate ion that is highly soluble in aqueous environments. This approach is employed in effervescent tablets, where the deprotonated salt dissolves rapidly, facilitating faster absorption and onset of analgesic effects. Such salt formation strategies are widely adopted to improve the delivery of ionizable drugs, addressing solubility challenges in oral and injectable formulations. Biologically, deprotonation is essential in , as exemplified by serine proteases such as , where the facilitates the deprotonation of the serine hydroxyl group. The residue acts as a base, abstracting a proton from serine-195, enhancing its nucleophilicity for attacking bonds in proteins during . This proton transfer mechanism accelerates the reaction by orders of magnitude, enabling efficient and cellular processes. Additionally, in blood pH regulation, the bicarbonate/CO₂ buffer system relies on the deprotonation of (H₂CO₃) to form ion (HCO₃⁻), which absorbs excess protons to maintain physiological between 7.35 and 7.45. This equilibrium, governed by the enzyme , counters or from metabolic or respiratory activities, ensuring . In applications, deprotonation supports proton conduction in (PEM) fuel cells, where polymer like sulfonated variants facilitate the transport of protons from to . Acid-base interactions within the membrane create protonation-deprotonation loops, particularly with moderately basic groups (e.g., benzotriazolium, pKa ≈ 1.6), which lower barriers for proton hopping and enhance conductivity under low-humidity conditions. This mechanism is critical for efficient oxidation and oxygen reduction, enabling PEM fuel cells to achieve power densities up to several hundred mW/cm² while operating at temperatures around 80°C. Advances in such deprotonation-enabled conduction have improved performance for automotive and portable power systems.

Quantification Methods

The extent of deprotonation in acid-base equilibria is quantified primarily through the KaK_a, defined as the for the dissociation of an acid HA into its conjugate base A⁻ and a proton H⁺: Ka=[A−][H+][HA]K_a = \frac{[A^-][H^+]}{[HA]}. The pKa value, expressed as pKa=−log⁡10Ka\mathrm{p}K_a = -\log_{10} K_a, provides a convenient to assess acidity, where a lower pKa indicates a stronger acid and thus easier deprotonation under given conditions. For instance, acetic acid (CH₃COOH) has a pKa of 4.76 in at 25°C, meaning it partially deprotonates, while itself has a pKa of 15.7 for the reaction H₂O ⇌ H⁺ + OH⁻, indicating much weaker acidity.
AcidpKa ValueSolvent/Temperature
(HCl)-6.3, 25°C
Acetic acid (CH₃COOH)4.76, 25°C
(H₂O)15.7, 25°C
(CH₃CH₂OH)15.9, 25°C
This table illustrates relative acid strengths, with values sourced from compilations of experimental data; note that pKa can vary with solvent and temperature. Experimentally, pKa values are determined via , where the is monitored as a base is added to the acid solution, yielding a sigmoid curve from which the at half-equivalence corresponds to the pKa. For kinetic aspects of deprotonation, such as proton exchange rates, NMR tracks changes in proton or carbon signals as a function of , allowing derivation of pKa through Henderson-Hasselbalch analysis of the profile. Computational prediction of pKa has advanced with (DFT) methods, which calculate free energy differences between protonated and deprotonated species in models, achieving accuracies within 1-2 units for organic acids when calibrated against experimental data. The thermodynamic basis for deprotonation extent links KaK_a to the standard change via the equation ΔG∘=−RTln⁡Ka\Delta G^\circ = -RT \ln K_a where RR is the (8.314 J/mol·K) and TT is temperature in ; a more negative ΔG∘\Delta G^\circ (higher KaK_a) favors deprotonation at equilibrium. To achieve complete deprotonation in practice, a base is selected such that the pKa of its conjugate exceeds that of the target by at least 2-3 units, ensuring the equilibrium shifts toward the deprotonated form; for example, using a base like ethoxide (conjugate ethanol, pKa 15.9) for alcohols (pKa ~15-18) promotes efficient deprotonation.

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