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Soaps are weak bases formed by the reaction of fatty acids with sodium hydroxide or potassium hydroxide.
Acids and bases
Diagrammatic representation of the dissociation of acetic acid in aqueous solution to acetate and hydronium ions.
Acid types
Base types

In chemistry, there are three definitions in common use of the word "base": Arrhenius bases, Brønsted bases, and Lewis bases. All definitions agree that bases are substances that react with acids, as originally proposed by G.-F. Rouelle in the mid-18th century.

In 1884, Svante Arrhenius proposed that a base is a substance which dissociates in aqueous solution to form hydroxide ions OH−. These ions can react with hydrogen ions (H+ according to Arrhenius) from the dissociation of acids to form water in an acid–base reaction. A base was therefore a metal hydroxide such as NaOH or Ca(OH)2. Such aqueous hydroxide solutions were also described by certain characteristic properties. They are slippery to the touch, can taste bitter[1] and change the color of pH indicators (e.g., turn red litmus paper blue).

In water, by altering the autoionization equilibrium, bases yield solutions in which the hydrogen ion activity is lower than it is in pure water, i.e., the water has a pH higher than 7.0 at standard conditions. A soluble base is called an alkali if it contains and releases OH− ions quantitatively. Metal oxides, hydroxides, and especially alkoxides are basic, and conjugate bases of weak acids are weak bases.

Bases and acids are seen as chemical opposites because the effect of an acid is to increase the hydronium (H3O+) concentration in water, whereas bases reduce this concentration. A reaction between aqueous solutions of an acid and a base is called neutralization, producing a solution of water and a salt in which the salt separates into its component ions. If the aqueous solution is saturated with a given salt solute, any additional such salt precipitates out of the solution.

In the more general Brønsted–Lowry acid–base theory (1923), a base is a substance that can accept hydrogen cations (H+)—otherwise known as protons. This does include aqueous hydroxides since OH− does react with H+ to form water, so that Arrhenius bases are a subset of Brønsted bases. However, there are also other Brønsted bases which accept protons, such as aqueous solutions of ammonia (NH3) or its organic derivatives (amines).[2] These bases do not contain a hydroxide ion but nevertheless react with water, resulting in an increase in the concentration of hydroxide ion.[3] Also, some non-aqueous solvents contain Brønsted bases which react with solvated protons. For example, in liquid ammonia, NH2− is the basic ion species which accepts protons from NH4+, the acidic species in this solvent.

G. N. Lewis realized that water, ammonia, and other bases can form a bond with a proton due to the unshared pair of electrons that the bases possess.[3] In the Lewis theory, a base is an electron pair donor which can share a pair of electrons with an electron acceptor which is described as a Lewis acid.[4] The Lewis theory is more general than the Brønsted model because the Lewis acid is not necessarily a proton, but can be another molecule (or ion) with a vacant low-lying orbital which can accept a pair of electrons. One notable example is boron trifluoride (BF3).

Some other definitions of both bases and acids have been proposed in the past, but are not commonly used today.

Properties

[edit]

General properties of bases include:

  • Concentrated or strong bases are caustic on organic matter and react violently with acidic substances.
  • Aqueous solutions or molten bases dissociate in ions and conduct electricity.
  • Reactions with indicators: bases turn red litmus paper blue, phenolphthalein pink, keep bromothymol blue in its natural colour of blue, and turn methyl orange-yellow.
  • The pH of a basic solution at standard conditions is greater than seven.
  • Bases are bitter.[5]

Reactions between bases and water

[edit]

The following reaction represents the general reaction between a base (B) and water to produce a conjugate acid (BH+) and a conjugate base (OH−):[3]The equilibrium constant, Kb, for this reaction can be found using the following general equation:[3]

In this equation, the base (B) and the extremely strong base (the conjugate base OH−) compete for the proton.[6] As a result, bases that react with water have relatively small equilibrium constant values.[6] The base is weaker when it has a lower equilibrium constant value.[3]

Neutralization of acids

[edit]
Ammonia fumes from aqueous ammonium hydroxide (in test tube) reacting with hydrochloric acid (in beaker) to produce ammonium chloride (white smoke).

Bases react with acids to neutralize each other at a fast rate both in water and in alcohol.[7] When dissolved in water, the strong base sodium hydroxide ionizes into hydroxide and sodium ions:

and similarly, in water the acid hydrogen chloride forms hydronium and chloride ions:

When the two solutions are mixed, the H
3O+
 and OH−
 ions combine to form water molecules:

If equal quantities of NaOH and HCl are dissolved, the base and the acid neutralize exactly, leaving only NaCl, effectively table salt, in solution.

Weak bases, such as baking soda or egg white, should be used to neutralize any acid spills. Neutralizing acid spills with strong bases, such as sodium hydroxide or potassium hydroxide, can cause a violent exothermic reaction, and the base itself can cause just as much damage as the original acid spill.

Alkalinity of non-hydroxides

[edit]

Bases are generally compounds that can neutralize an amount of acid. Both sodium carbonate and ammonia are bases, although neither of these substances contains OH−
 groups. Both compounds accept H+ when dissolved in protic solvents such as water:

From this, a pH, or acidity, can be calculated for aqueous solutions of bases.

A base is also defined as a molecule that has the ability to accept an electron pair bond by entering another atom's valence shell through its possession of one electron pair.[7] There are a limited number of elements that have atoms with the ability to provide a molecule with basic properties.[7] Carbon can act as a base as well as nitrogen and oxygen. Fluorine and sometimes rare gases possess this ability as well.[7] This occurs typically in compounds such as butyl lithium, alkoxides, and metal amides such as sodium amide. Bases of carbon, nitrogen and oxygen without resonance stabilization are usually very strong, or superbases, which cannot exist in a water solution due to the acidity of water. Resonance stabilization, however, enables weaker bases such as carboxylates; for example, sodium acetate is a weak base.

Strong bases

[edit]

A strong base is a basic chemical compound that can remove a proton (H+) from (or deprotonate) a molecule of even a very weak acid (such as water) in an acid–base reaction. Common examples of strong bases include hydroxides of alkali metals and alkaline earth metals, like sodium hydroxide and calcium hydroxide, respectively. Due to their low solubility, some bases, such as alkaline earth hydroxides, can be used when the solubility factor is not taken into account.[8]

One advantage of this low solubility is that "many antacids were suspensions of metal hydroxides such as aluminium hydroxide and magnesium hydroxide";[9] compounds with low solubility and the ability to stop an increase in the concentration of the hydroxide ion, preventing the harm of the tissues in the mouth, oesophagus, and stomach.[9] As the reaction continues and the salts dissolve, the stomach acid reacts with the hydroxide produced by the suspensions.[9]

Strong bases hydrolyze in water almost completely, resulting in the leveling effect.[7] In this process, the water molecule combines with a strong base, due to the water's amphoteric ability; and, a hydroxide ion is released.[7] Very strong bases can even deprotonate very weakly acidic C–H groups in the absence of water. Here is a list of several strong bases:

Lithium hydroxide LiOH
Sodium hydroxide NaOH
Potassium hydroxide KOH
Rubidium hydroxide RbOH
Cesium hydroxide CsOH
Magnesium hydroxide Mg(OH)
2
Calcium hydroxide Ca(OH)
2
Strontium hydroxide Sr(OH)
2
Barium hydroxide Ba(OH)
2
Tetramethylammonium hydroxide N(CH
3)
4OH
Guanidine HNC(NH
2)
2

The cations of these strong bases appear in the first and second groups of the periodic table (alkali and earth alkali metals). Tetraalkylated ammonium hydroxides are also strong bases since they dissociate completely in water. Guanidine is a special case of a species that is exceptionally stable when protonated, analogously to the reason that makes perchloric acid and sulfuric acid very strong acids.

Acids with a pKa of more than about 13 are considered very weak, and their conjugate bases are strong bases.

Superbases

[edit]
Main article: Superbase

Group 1 salts of carbanions, amide ions, and hydrides tend to be even stronger bases due to the extreme weakness of their conjugate acids, which are stable hydrocarbons, amines, and dihydrogen. Usually, these bases are created by adding pure alkali metals such as sodium into the conjugate acid. They are called superbases, and it is impossible to keep them in aqueous solutions because they are stronger bases than the hydroxide ion (see the leveling effect). For example, the ethoxide ion (conjugate base of ethanol) undergoes this reaction quantitatively in presence of water.[10]

Examples of common superbases are:

Strongest superbases are synthesised in only gas phase:

Weak bases

[edit]
Main article: weak base

A weak base is one which does not fully ionize in an aqueous solution, or in which protonation is incomplete. For example, ammonia transfers a proton to water according to the equation[11]

NH3(aq) + H2O(l) → NH+
4(aq) + OH−(aq)

The equilibrium constant for this reaction at 25 °C is 1.8 x 10−5,[12] such that the extent of reaction or degree of ionization is quite small.

Lewis bases

[edit]

A Lewis base or electron-pair donor is a molecule with one or more high-energy lone pairs of electrons which can be shared with a low-energy vacant orbital in an acceptor molecule to form an adduct. In addition to H+, possible electron-pair acceptors (Lewis acids) include neutral molecules such as BF3 and high oxidation state metal ions such as Ag2+, Fe3+ and Mn7+. Adducts involving metal ions are usually described as coordination complexes.[13]

According to the original formulation of Lewis, when a neutral base forms a bond with a neutral acid, a condition of electric stress occurs.[7] The acid and the base share the electron pair that formerly belonged to the base.[7] As a result, a high dipole moment is created, which can only be decreased to zero by rearranging the molecules.[7]

Solid bases

[edit]

Examples of solid bases include:

  • Oxide mixtures: SiO2, Al2O3; MgO, SiO2; CaO, SiO2[14]
  • Mounted bases: LiCO3 on silica; NR3, NH3, KNH2 on alumina; NaOH, KOH mounted on silica on alumina[14]
  • Inorganic chemicals: BaO, KNaCO3, BeO, MgO, CaO, KCN[14]
  • Anion exchange resins[14]
  • Charcoal that has been treated at 900 degrees Celsius or activates with N2O, NH3, ZnCl2-NH4Cl-CO2[14]

Depending on a solid surface's ability to successfully form a conjugate base by absorbing an electrically neutral acid, basic strength of the surface is determined.[15] The "number of basic sites per unit surface area of the solid" is used to express how much basic strength is found on a solid base catalyst.[15] Scientists have developed two methods to measure the amount of basic sites: one, titration with benzoic acid using indicators and gaseous acid adsorption.[15] A solid with enough basic strength will absorb an electrically neutral acidic indicator and cause the acidic indicator's color to change to the color of its conjugate base.[15] When performing the gaseous acid adsorption method, nitric oxide is used.[15] The basic sites are then determined by calculating the amount of carbon dioxide that is absorbed.[15]

Bases as catalysts

[edit]

Basic substances can be used as insoluble heterogeneous catalysts for chemical reactions. Some examples are metal oxides such as magnesium oxide, calcium oxide, and barium oxide as well as potassium fluoride on alumina and some zeolites. Many transition metals make good catalysts, many of which form basic substances. Basic catalysts are used for hydrogenation, the migration of double bonds, in the Meerwein-Ponndorf-Verley reduction, the Michael reaction, and many others. Both CaO and BaO can be highly active catalysts if they are heated to high temperatures.[15]

Uses of bases

[edit]
  • Sodium hydroxide is used in the manufacture of soap, paper, and the synthetic fiber rayon.
  • Calcium hydroxide (slaked lime) is used in the manufacture of bleaching powder.
  • Calcium hydroxide is also used to clean the sulfur dioxide, which is caused by the exhaust, that is found in power plants and factories.[9]
  • Magnesium hydroxide is used as an 'antacid' to neutralize excess acid in the stomach and cure indigestion.
  • Sodium carbonate is used as washing soda and for softening hard water.
  • Sodium bicarbonate (or sodium hydrogen carbonate) is used as baking soda in cooking food, for making baking powders, as an antacid to cure indigestion and in soda acid fire extinguisher.
  • Ammonium hydroxide is used to remove grease stains from clothes

Monoprotic and polyprotic bases

[edit]

Bases with only one ionizable hydroxide (OH−) ion per formula unit are called monoprotic since they can accept one proton (H+). Bases with more than one OH- per formula unit are polyprotic.[16]

The number of ionizable hydroxide (OH−) ions present in one formula unit of a base is also called the acidity of the base.[17][18] On the basis of acidity bases can be classified into three types: monoacidic, diacidic and triacidic.

Monoacidic bases

[edit]
Sodium hydroxide

When one molecule of a base via complete ionization produces one hydroxide ion, the base is said to be a monoacidic or monoprotic base. Examples of monoacidic bases are:

Sodium hydroxide, potassium hydroxide, silver hydroxide, ammonium hydroxide, etc.

Diacidic bases

[edit]

When one molecule of base via complete ionization produces two hydroxide ions, the base is said to be diacidic or diprotic. Examples of diacidic bases are:

Barium hydroxide

Barium hydroxide, magnesium hydroxide, calcium hydroxide, zinc hydroxide, iron(II) hydroxide, tin(II) hydroxide, lead(II) hydroxide, copper(II) hydroxide, etc.

Triacidic bases

[edit]

When one molecule of base via complete ionization produces three hydroxide ions, the base is said to be triacidic or triprotic. Examples of triacidic bases are:

Aluminium hydroxide, ferrous hydroxide, Gold Trihydroxide,[18]

Etymology of the term

[edit]

The concept of base stems from an older alchemical notion of "the matrix":

The term "base" appears to have been first used in 1717 by the French chemist, Louis Lémery, as a synonym for the older Paracelsian term "matrix." In keeping with 16th-century animism, Paracelsus had postulated that naturally occurring salts grew within the earth as a result of a universal acid or seminal principle having impregnated an earthy matrix or womb. ... Its modern meaning and general introduction into the chemical vocabulary, however, is usually attributed to the French chemist, Guillaume-François Rouelle. ... In 1754 Rouelle explicitly defined a neutral salt as the product formed by the union of an acid with any substance, be it a water-soluble alkali, a volatile alkali, an absorbent earth, a metal, or an oil, capable of serving as "a base" for the salt "by giving it a concrete or solid form." Most acids known in the 18th century were volatile liquids or "spirits" capable of distillation, whereas salts, by their very nature, were crystalline solids. Hence it was the substance that neutralized the acid which supposedly destroyed the volatility or spirit of the acid and which imparted the property of solidity (i.e., gave a concrete base) to the resulting salt.

— William B. Jensen, The origin of the term "base"[19]

See also

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References

[edit]
[edit]
Revisions and contributorsEdit on WikipediaRead on Wikipedia

Base (chemistry)

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from Grokipedia
In chemistry, a base is defined as a substance capable of accepting protons (H⁺ ions) according to the Brønsted–Lowry or donating an electron pair according to the Lewis , serving as the counterpart to acids in fundamental chemical reactions. Under the earlier Arrhenius model, a base is any compound that increases the hydroxide ion (OH⁻) concentration in by dissociating to release OH⁻ ions. These definitions encompass a wide range of compounds, from strong bases like (NaOH) that fully ionize in to weak bases like (NH₃) that only partially ionize, and they enable bases to neutralize acids by forming and salts. Bases exhibit distinct physical and chemical properties that distinguish them from other substances. Physically, they often have a bitter , a slippery or soapy texture due to their ability to dissolve fats and greases, and they can irritate the skin or eyes upon contact. Chemically, bases turn red paper blue, react with acids in neutralization reactions to produce and ionic salts, and in aqueous solutions, they raise the above 7 by increasing OH⁻ concentration relative to H⁺ ions. Strong bases, such as (KOH) and (Ca(OH)₂), ionize completely and are highly reactive, while weak bases like (Mg(OH)₂) ionize only partially and are less corrosive. The concept of bases is essential across chemistry, influencing everything from everyday applications like soap production and antacids to advanced fields such as biochemistry and , where they facilitate regulation, , and synthesis processes.

Definitions and Properties

Arrhenius and Brønsted-Lowry Definitions

The Arrhenius definition of a base, introduced by Swedish Svante in his 1884 dissertation on electrolytic dissociation, identifies a base as a substance that increases the concentration of hydroxide ions (OH⁻) when dissolved in . This theory built on observations of behavior and explained the common properties of bases, such as their bitter taste and slippery feel, through the production of OH⁻ ions, which can neutralize H⁺ from acids to form . A classic example is (NaOH), which fully dissociates in as NaOH → Na⁺ + OH⁻, directly elevating the OH⁻ concentration. In contrast, the Brønsted-Lowry theory, proposed independently by Danish chemist and British chemist in 1923, defines a base as any species capable of accepting a proton (H⁺) from an acid./Acids_and_Bases/Acid/Bronsted_Concept_of_Acids_and_Bases) This proton-transfer model introduces the concept of conjugate acid-base pairs, where a base accepts a proton to become its conjugate acid, and an acid donates a proton to become its conjugate base. For instance, (NH₃) functions as a Brønsted-Lowry base in without initially containing OH⁻: \ceNH3+H2O⇌NH4++OH−\ce{NH3 + H2O ⇌ NH4+ + OH-} Here, NH₃ accepts H⁺ from H₂O (acting as the acid), forming the conjugate acid NH₄⁺ and conjugate base OH⁻. The Brønsted-Lowry definition expands upon the Arrhenius model by applying to any proton-transfer reaction, not limited to aqueous environments or hydroxide production, thus accommodating bases in non-aqueous solvents or gas-phase reactions. While all Arrhenius bases qualify as Brønsted-Lowry bases (since OH⁻ accepts protons), the reverse is not true, as demonstrated by NH₃, which generates OH⁻ indirectly rather than directly. This broader scope marked a significant advancement in understanding acid-base chemistry beyond water-based systems.

Lewis Definition

In 1916, Gilbert N. Lewis proposed the concept of the chemical bond as a shared pair of electrons between atoms, laying the groundwork for his later acid-base theory. In 1923, Lewis formalized the electron-pair theory of acids and bases in his monograph Valence and the Structure of Atoms and Molecules, defining a base as any species capable of donating an electron pair to form a coordinate covalent bond with an acid, which acts as the electron-pair acceptor. This framework broadens the understanding of acid-base interactions beyond proton transfer, emphasizing electron donation in diverse chemical contexts./15%3A_Equilibria_of_Other_Reaction_Classes/15.03%3A_Lewis_Acids_and_Bases) Under the Lewis definition, a base is characterized by the presence of a lone pair of electrons or a pi bond that can be shared to form a dative bond, resulting in a Lewis acid-base adduct. A classic example is ammonia (\ceNH3\ce{NH3}), which serves as a Lewis base by donating its nitrogen lone pair to boron trifluoride (\ceBF3\ce{BF3}), an electron-deficient Lewis acid, to yield the stable adduct \ceNH3⋅BF3\ce{NH3 \cdot BF3}. Similarly, the hydroxide ion (\ceOH−\ce{OH^-}) functions as a Lewis base when it donates an electron pair to a proton (\ceH+\ce{H^+}), forming a water molecule./15%3A_Equilibria_of_Other_Reaction_Classes/15.03%3A_Lewis_Acids_and_Bases) In coordination chemistry, amines such as ethylamine act as Lewis bases by providing electron pairs to metal cations, forming complexes like those in transition metal ammine compounds. The Lewis definition distinguishes itself from the Brønsted-Lowry theory—where bases accept protons and thus involve a specific subset of electron-pair donations—by encompassing acid-base reactions that do not require protons, such as the formation of adducts in the gas phase or non-aqueous organic solvents./15%3A_Equilibria_of_Other_Reaction_Classes/15.03%3A_Lewis_Acids_and_Bases) This generality makes it particularly advantageous for analyzing reactions in aprotic environments and coordination compounds, where traditional proton-based definitions fall short.

Physical and Chemical Properties

Bases in chemistry exhibit distinct physical properties that distinguish them from other substances. Aqueous solutions of bases typically have a bitter taste, though tasting chemicals in a laboratory setting is strongly discouraged due to safety risks. They also produce a slippery or soapy feel when touched, resulting from the saponification reaction between the base and the fatty oils on the skin, which forms soap-like compounds. Additionally, bases are corrosive to biological tissues, capable of causing severe chemical burns upon contact, as seen with common strong bases like lye (sodium hydroxide, NaOH). Chemically, bases in aqueous solutions have a greater than 7, indicating a higher concentration of ions relative to ions. They react exothermically with acids in neutralization reactions, releasing heat and forming salts and . The intensity of these properties, such as corrosiveness and pH elevation, is more pronounced in strong bases, which fully dissociate in , compared to weak bases that partially ionize. Common tests for identifying bases rely on acid-base indicators. Bases turn red paper , a classic qualitative test for basicity. , another widely used indicator, remains colorless in acidic or neutral solutions but turns pink or magenta in basic conditions ( > 8.2). solutions display a range of colors for basic values, shifting from green (neutral, 7) to and violet ( 8–14) in increasingly basic environments.

Classification of Bases

Strong Bases

Strong bases are chemical compounds that completely dissociate in water, producing a high concentration of ions (OH⁻) and thereby exhibiting strong basic properties. This full distinguishes them from weak bases, which only partially dissociate. For instance, dissociates according to the equation \ceNaOH−>Na++OH−\ce{NaOH -> Na+ + OH-}. The base dissociation constant (KbK_b) for strong bases is extremely large, often approaching infinity, indicating negligible reverse reaction in aqueous solutions. Common examples of strong bases include the hydroxides of alkali metals, such as (LiOH), (NaOH), (KOH), (RbOH), and cesium hydroxide (CsOH), all of which are highly soluble in . Hydroxides of heavier alkaline earth metals, like (Ca(OH)₂), (Sr(OH)₂), and (Ba(OH)₂), also qualify as strong bases due to complete of the dissolved portion, though their solubility is limited—Ca(OH)₂, for example, has a solubility of approximately 0.173 g/100 mL at 20°C. These compounds are widely used in industrial applications, such as production and . Solutions of strong bases exhibit high electrical conductivity owing to the abundance of free from complete dissociation. Dissolution of these bases in is typically highly exothermic, generating significant that can cause burns or splattering if not managed properly—for NaOH, the process releases about 44.5 kJ/mol. In concentrated solutions, such as 1 M NaOH, the pH approaches 14, reflecting a concentration of 1 M. Strong bases pose significant hazards due to their extreme corrosiveness, capable of causing severe chemical burns to , eyes, and mucous membranes upon contact, as well as corroding metals and organic materials. Solutions with ≥ 12.5 are particularly dangerous, necessitating immediate rinsing with for at least 15 minutes in case of exposure. In industrial settings, handling requires (PPE) including gloves, goggles, and protective clothing, along with storage in dedicated corrosive-resistant cabinets and use of local exhaust ventilation to mitigate risks.

Weak Bases

Weak bases are compounds that do not completely dissociate in , instead establishing an equilibrium in which only a small fraction to produce ions or accept protons. This partial ionization results in a base dissociation constant KbK_b less than 1, distinguishing them from strong bases that fully ionize. The equilibrium for a generic weak base \ceB\ce{B} is given by \ceB+H2O⇌BH++OH−\ce{B + H2O ⇌ BH+ + OH-} with the expression Kb=[\ceBH+][\ceOH−][\ceB]K_b = \frac{[\ce{BH+}][\ce{OH-}]}{[\ce{B}]} at 25°C, where concentrations are in moles per liter. The strength of the base is conveniently quantified by pKb=−log⁡Kb\mathrm{p}K_b = -\log K_b, where lower pKb\mathrm{p}K_b values indicate stronger bases among the weak category. Representative examples include , \ceNH3\ce{NH3}, which reacts as \ceNH3+H2O⇌NH4++OH−\ce{NH3 + H2O ⇌ NH4+ + OH-} with Kb=1.8×10−5K_b = 1.8 \times 10^{-5} (pKb=4.74\mathrm{p}K_b = 4.74), and various amines such as , \ce(CH3)2NH\ce{(CH3)2NH}, with Kb=5.4×10−4K_b = 5.4 \times 10^{-4} (pKb=3.27\mathrm{p}K_b = 3.27), and , \ceC6H5NH2\ce{C6H5NH2}, a much weaker base with Kb=4.3×10−10K_b = 4.3 \times 10^{-10} (pKb=9.37\mathrm{p}K_b = 9.37). Inorganic examples encompass the carbonate ion, \ceCO32−\ce{CO3^2-}, which acts as \ceCO32−+H2O⇌HCO3−+OH−\ce{CO3^2- + H2O ⇌ HCO3- + OH-} with Kb≈2.1×10−4K_b \approx 2.1 \times 10^{-4} (pKb≈3.68\mathrm{p}K_b \approx 3.68), calculated from Kw/KaK_w / K_a where KaK_a for \ceHCO3−\ce{HCO3-} is 4.7×10−114.7 \times 10^{-11}. These values illustrate how KbK_b quantifies the position of equilibrium, with smaller constants reflecting weaker basicity. The basicity of weak bases, particularly amines, is influenced by inductive effects, where electron-donating alkyl groups increase on the , enhancing proton acceptance compared to electron-withdrawing groups like those in arylamines. in aqueous media further modulates strength by stabilizing the charged conjugate acid through hydrogen bonding, though this effect diminishes for larger or less polarizable bases.

Superbases

Superbases represent a class of compounds with extraordinarily high basicity, typically defined as bases whose conjugate acids possess pKa values greater than 25 in non-aqueous solvents, with some exceeding 40, surpassing the strength of typical strong bases such as hydroxides. These agents are particularly valued in non-aqueous environments where standard bases fail to deprotonate highly stable carbon-hydrogen bonds or other weak acids. The term "superbase" emerged in the context of to describe reagents capable of generating carbanions from substrates with pKa values around 30–45, enabling reactions inaccessible to milder bases. Organic superbases, such as phosphazenes and guanidines, are prominent due to their tunable structures and thermal stability, often designed to avoid nucleophilicity while maximizing proton abstraction. Phosphazenes, exemplified by Schwesinger's P4 base (1-tert-butyl-4,4,4-tris(dimethylamino)-2,2,6,6-tetramethyl-2,6-diphosphacyclohexane-1,3,5-triene, or t-Bu-P4), exhibit a conjugate acid pKa of approximately 42 in , making it one of the strongest non-ionic organic bases available. Guanidines, like 1,5,7-triazabicyclo[4.4.0]dec-5-ene (TBD), achieve high basicity through delocalization of the positive charge in their protonated forms, with pKa values for conjugate acids around 26 in . These organic variants are favored for their solubility in organic solvents and reduced tendency to form salts that complicate workups. Inorganic superbases include species like the amide (NH₂⁻), generated from (NaNH₂), where the conjugate acid (NH₃) has a pKa of about 38 in , classifying it as a for deprotonating terminal alkynes and other C-H acids with pKa >30. Another notable example is potassium tert-butoxide (t-BuOK) in (DMSO), which functions as an effective system; the low solvating ability of DMSO enhances the basicity of the beyond its aqueous pKa of 19 for t-BuOH, allowing of compounds with pKa up to 35. These inorganic systems are often employed in aprotic solvents to minimize proton donation from the medium. In , superbases facilitate the of exceptionally weak acids, such as hydrocarbons or activated methylene compounds, to form reactive carbanions for subsequent bond-forming reactions like or , without the side reactions common to metal-based bases. This selective proton abstraction is crucial for constructing complex carbon frameworks in synthesis.

Reactions Involving Bases

Dissociation in Water

In aqueous solutions, Brønsted-Lowry bases undergo dissociation by accepting a proton from molecules, forming their conjugate and ions according to the general equilibrium reaction: B+H2O⇌BH++OH−\text{B} + \text{H}_2\text{O} \rightleftharpoons \text{BH}^+ + \text{OH}^- This process increases the concentration of ions (OH⁻) in the solution, which is characteristic of basic conditions. The extent of this reaction varies; strong bases dissociate nearly completely, while weak bases reach only partial equilibrium. This base dissociation interacts with water's autoionization, where two water molecules react to produce hydronium (H₃O⁺) and hydroxide ions: 2H₂O ⇌ H₃O⁺ + OH⁻. The equilibrium constant for this autoionization, known as the ion product of water (K_w), is defined as K_w = [H⁺][OH⁻] = 1.0 × 10^{-14} at 25°C. In basic solutions, the OH⁻ generated from base dissociation suppresses the autoionization to maintain this constant, resulting in lower [H⁺] than in pure water. The acidity or basicity of these solutions is quantified using , defined as pH = -log[H⁺]. For basic solutions at 25°C, where [OH⁻] > 10^{-7} M, the relationship pH + pOH = 14 holds, with pOH = -log[OH⁻]; thus, pH = 14 - pOH provides a direct way to calculate pH from measured or known concentrations. The dissociation behavior of bases in water is temperature-dependent because K_w increases with rising —for instance, K_w ≈ 2.5 × 10^{-14} at 37°C—leading to greater autoionization and a shift in the base-water equilibrium toward more OH⁻ production for weak bases. This endothermic nature of autoionization means higher temperatures reduce the neutrality pH from 7 at 25°C to approximately 6.8 at 37°C.

Neutralization Reactions

Neutralization reactions occur when acids and bases react to form a salt and water, effectively reducing the concentrations of hydrogen ions (H⁺) and hydroxide ions (OH⁻) in solution. According to the Arrhenius definition, this process involves the combination of H⁺ from the acid and OH⁻ from the base to produce water, leaving behind the salt derived from the cations and anions of the reactants. The general equation for such reactions is acid + base → salt + H₂O. A classic example is the reaction between hydrochloric acid and sodium hydroxide: \ceHCl(aq)+NaOH(aq)−>NaCl(aq)+H2O(l)\ce{HCl(aq) + NaOH(aq) -> NaCl(aq) + H2O(l)} This reaction exemplifies a complete neutralization where stoichiometric amounts of acid and base yield neutral products. The of neutralization reactions depends on the number of ionizable protons in the and hydroxyl groups in the base. For monoprotic s and monobasic bases, such as HCl and NaOH, the reaction follows a 1:1 molar , meaning one mole of neutralizes one mole of base. In cases involving polyprotic s or bases, multiple equivalence points may appear due to stepwise neutralization of each proton or group. For instance, (H₂SO₄), a diprotic , requires two moles of NaOH for complete neutralization. These stoichiometric relationships are crucial for quantitative analysis in chemical reactions. Neutralization reactions are exothermic, releasing heat as the ionic bonds in form. For acids and bases in dilute aqueous solutions, the heat of neutralization is approximately 57 kJ/mol of produced, reflecting the consistent change for the formation of H₂O from H⁺ and OH⁻ regardless of the specific acid or base involved. This value arises because electrolytes are fully dissociated, leading to a uniform net ionic : H⁺(aq) + OH⁻(aq) → H₂O(l). Deviations occur with weak acids or bases due to partial dissociation and additional enthalpic contributions from . In , neutralization reactions form the basis of , where a base of known concentration is added to an solution (or vice versa) to determine the unknown concentration. curves plot against the volume of titrant added, revealing characteristic shapes based on and base strengths. For a strong titrated with a strong base, the occurs at 7, where equal moles of and base have reacted, resulting in a neutral solution. In contrast, titrating a weak with a strong base yields an above 7, as the conjugate base of the weak hydrolyzes to produce excess OH⁻. These curves enable precise endpoint detection using indicators or meters, with the steep rise near equivalence highlighting the reaction's completion.

Reactions with Non-Aqueous Solvents

In non-aqueous solvents, the that limits the observable strength of strong bases in is absent, enabling the differentiation and utilization of bases with intrinsic basicities exceeding that of the hydroxide ion. This absence arises because many non-aqueous solvents lack the amphoteric nature of , which autoionizes to set an upper limit on basicity; instead, these solvents allow bases to exhibit their full proton-accepting capacity without rapid proton transfer to the solvent. For instance, liquid ammonia serves as a protophilic solvent where the conjugate base, amide ion (NH₂⁻), defines the leveling threshold, permitting the study of superbases like alkali metal amides that would be indistinguishable in aqueous media. Specific examples illustrate varying base behaviors across non-aqueous media. In , a weaker as a proton donor than due to its lower constant and reduced hydrogen-bonding capacity, bases such as alkoxides display enhanced apparent basicity relative to their aqueous counterparts, allowing finer control in reactions like deprotonations. Conversely, in (DMSO), an aprotic polar solvent with high donor number, systems like (KOH) or potassium tert-butoxide generate superbasic environments by minimizing proton , facilitating reactions inaccessible in protic solvents. These differences highlight how solvent polarity and proticity dictate base reactivity, with aprotic solvents like DMSO promoting higher basic strengths for applications in . Solvent effects on basicity are profoundly influenced by donor-acceptor properties, which modulate the pKa of a base's conjugate through differential stabilization of charged . Protic solvents with strong hydrogen-bond donor abilities, such as alcohols, solvate anions less effectively than , often increasing measured pKa values and thus apparent basicity, while aprotic solvents with electron-donor capabilities like DMSO enhance anion stability via coordination, lowering pKa and amplifying basic strength. These interactions arise from the solvent's ability to act as a Lewis base or , altering the free energy of proton transfer without the autoionization constraints of . In some contexts, bases function as Lewis bases by coordinating directly to solvent molecules, further tuning reactivity. Spectroscopic techniques provide essential tools for detecting and quantifying basicity in non-aqueous systems, where traditional measurements fail due to varying autoionization constants. UV-visible using acidochromic indicator probes monitors color shifts from equilibria, enabling pKa determination by observing absorbance changes as base concentration varies. () complements this by tracking perturbations in proton or heteronuclear signals during acid-base interactions, offering insights into and proton transfer dynamics in solvents like DMSO or . These methods ensure precise characterization without relying on aqueous standards.

Specialized Types of Bases

Non-Hydroxide Bases

Non-hydroxide bases are chemical compounds that demonstrate basic properties by accepting protons or generating ions through reactions such as , without inherently containing the (OH⁻) group in their molecular structure. These bases expand the concept of beyond traditional metal hydroxides, encompassing a variety of including metal oxides, carbonates, and amides that interact with or acids to produce basic solutions. A prominent example is metal oxides, such as (CaO), commonly known as quicklime, which reacts exothermically with water to form :
\ceCaO+H2O−>Ca(OH)2\ce{CaO + H2O -> Ca(OH)2}
This reaction underscores the basic nature of metal oxides, as the resulting hydroxide ions increase the of the solution. Similarly, (NaNH₂) serves as a strong non-hydroxide base, acting as a proton acceptor in non-aqueous environments due to the high basicity of the ion (NH₂⁻). Carbonates, like (Na₂CO₃), exhibit basicity through of the ion in :
\ceCO32−+H2O⇌HCO3−+OH−\ce{CO3^2- + H2O ⇌ HCO3^- + OH^-}
This equilibrium shifts to produce ions, making solutions alkaline, with the extent of depending on the . (NH₃) and amines function as weak bases by accepting a proton to form ions:
\ceNH3+H2O⇌NH4++OH−\ce{NH3 + H2O ⇌ NH4^+ + OH^-}
These examples illustrate how non- bases rely on conjugate base strength or reactive anions to manifest . In environmental contexts, non-hydroxide bases like carbonates contribute significantly to , particularly in arid regions where accumulation leads to high levels that affect nutrient availability and microbial activity. This natural occurrence of carbonate-driven influences remediation strategies and agricultural practices.

Solid Bases

Solid bases refer to insoluble materials that manifest basic properties primarily through surface sites, distinguishing them from soluble bases by their role in heterogeneous systems. These materials, often metal oxides, possess basic sites such as oxygen anions (O²⁻), groups (OH⁻), or phenoxide species on their surfaces, which interact with acidic probes or reactants. Representative examples include oxides like (MgO) and (CaO), as well as oxides such as zinc oxide (ZnO), where the basicity arises from low-coordinated surface oxygen atoms that can donate pairs. The characterization of solid basicity focuses on assessing both the strength and density of these surface sites. Hammett indicators, a colorimetric method adapted for solids, measure basic strength by titrating the material with indicators of known pK_B values (e.g., for H_ = 9.3 or for weaker sites), where color changes indicate the point at which the indicator's conjugate acid is deprotonated by the solid's basic sites. This technique classifies basic strength on the Hammett H_ scale, with superbasic sites exceeding H_ > 18. Complementarily, CO₂ adsorption techniques, including temperature-programmed desorption (TPD), quantify basic site density and strength; CO₂, acting as a Lewis acid, adsorbs selectively on strong basic sites to form surface carbonates or bicarbonates, with desorption temperatures correlating to site strength (e.g., >400°C for strong sites on MgO). Solid bases are categorized into several types based on composition and structure. Single metal oxides, such as MgO and ZnO, offer uniform basic sites but limited tunability, with MgO exhibiting strong basicity due to its ionic lattice. Mixed metal oxides, including those derived from hydrotalcites—calcined like Mg-Al oxides—provide enhanced basicity and site diversity through their defective structures and adjustable metal ratios, enabling moderate to strong basic sites. Zeolites, typically acidic, can be engineered as basic solids via exchange or desilication, creating isolated O²⁻ sites within their microporous frameworks for shape-selective basic . In industrial applications, solid bases demonstrate notable thermal and , often withstanding temperatures up to 800°C without significant loss of activity. Regeneration is achieved through or oxidative treatments to remove adsorbed species like CO₂ or , restoring surface basic sites; for instance, hydrotalcite-derived oxides can be reused multiple cycles after such processes with minimal leaching. These materials' surface basic sites facilitate in reactions requiring base activation, such as aldol condensations.

Polyprotic Bases

Polyprotic bases are that can accept more than one proton (H⁺) due to the presence of multiple basic sites, such as lone pairs on or negatively charged oxygen atoms. These bases undergo stepwise in aqueous solutions, with each step characterized by a distinct base dissociation constant (K_b). Unlike monoprotic bases, polyprotic bases exhibit more complex acid-base behavior, influencing their reactivity and utility in various chemical systems. The protonation of a polyprotic base proceeds sequentially, with the first step typically being much stronger than subsequent ones because the addition of the initial proton reduces the electron density available for further protonation, increasing electrostatic repulsion. For a diprotic base denoted as B²⁻, the equilibria are: B2−+H2O⇌HB−+OH−Kb1=[HB−][OH−][B2−]\text{B}^{2-} + \text{H}_2\text{O} \rightleftharpoons \text{HB}^- + \text{OH}^- \quad K_{b1} = \frac{[\text{HB}^-][\text{OH}^-]}{[\text{B}^{2-}]} HB−+H2O⇌H2B+OH−Kb2=[H2B][OH−][HB−]\text{HB}^- + \text{H}_2\text{O} \rightleftharpoons \text{H}_2\text{B} + \text{OH}^- \quad K_{b2} = \frac{[\text{H}_2\text{B}][\text{OH}^-]}{[\text{HB}^-]} Here, K_{b1} \gg K_{b2}, often by several orders of magnitude, meaning the first protonation dominates in moderately basic solutions. These constants are related to the acid dissociation constants (K_a) of the conjugate acids via K_b = K_w / K_a, where K_w is the product of . Representative examples include the carbonate ion (CO₃²⁻), a diprotic base that protonates first to (HCO₃⁻) and then to (H₂CO₃), with K_{b1} \approx 2.1 \times 10^{-4} and K_{b2} \approx 2.3 \times 10^{-8}. The (PO₄³⁻) serves as a triprotic base, sequentially forming HPO₄²⁻, H₂PO₄⁻, and H₃PO₄, exhibiting weak basicity overall but with stepwise K_b values decreasing from K_{b1} \approx 2.1 \times 10^{-2} to K_{b3} \approx 1.3 \times 10^{-12}. An organic example is (H₂NCH₂CH₂NH₂), a diprotic base with two groups, where the pK_b values are approximately 4.07 and 7.15, corresponding to K_{b1} \approx 8.5 \times 10^{-5} and K_{b2} \approx 7.1 \times 10^{-8}. Polyprotic bases are particularly valuable in buffer systems, where conjugate acid-base pairs from different steps provide buffering capacity across multiple ranges. For instance, the / pair (HCO₃⁻/CO₃²⁻) effectively buffers near 10, while systems offer buffers around 7 (H₂PO₄⁻/HPO₄²⁻) and 12 (HPO₄²⁻/PO₄³⁻), making them essential in biological and analytical applications. In neutralization reactions, these bases require multiple equivalents of acid for complete , complicating curves with multiple equivalence points.

Applications of Bases

Catalytic Roles

Bases function as catalysts in organic reactions primarily by deprotonating substrates to generate nucleophilic intermediates such as carbanions or , which then react with electrophiles to form new bonds. This mechanism lowers the by stabilizing the through charge delocalization in the enolate, where the negative charge is shared between the alpha-carbon and the carbonyl oxygen. The process is reversible, with the base regenerating after proton transfer back to the product, ensuring catalytic turnover. Base catalysis occurs in two main forms: homogeneous and heterogeneous. In homogeneous catalysis, the base is dissolved in the reaction medium, such as aqueous NaOH, which provides high mobility and uniform reactivity in solution-phase organic transformations. Heterogeneous catalysis, conversely, utilizes insoluble solid bases like MgO, where surface basic sites facilitate ; these systems enable easy catalyst recovery and are preferred for large-scale processes due to reduced and . Prominent examples illustrate these catalytic roles. In the , OH⁻ deprotonates the alpha-carbon of an or to form an , which nucleophilically adds to a second carbonyl, yielding a β-hydroxy carbonyl compound; subsequent often affords α,β-unsaturated products under basic conditions. The employs bases to generate ester , promoting intermolecular attack on another carbonyl to construct β-keto esters, a key step in synthesis. Similarly, the uses a strong base to abstract a β-proton from a quaternary salt, inducing E2 elimination to produce alkenes with favoring the less substituted (Hofmann) product due to the bulky . Basic media often impart and significant rate enhancements to these reactions. geometry, influenced by the base and conditions, directs diastereoselective additions, as seen in aldol reactions where Z-enolates yield products via Zimmerman-Traxler transition states. Rate accelerations arise from base-assisted stabilization of charged intermediates.

Industrial and Uses

Bases are integral to numerous , enabling the production of everyday materials and the treatment of resources on a large scale. In soap manufacturing, (NaOH) serves as the key reagent in , where it reacts with animal fats or vegetable oils to break bonds, yielding (sodium salts of fatty acids) and as a ; this process accounts for a significant portion of global NaOH consumption, with modern continuous kettle or split methods enhancing efficiency. In the paper industry, alkali pulping—often using NaOH in the —dissolves from wood chips, liberating fibers for ; approximately 20% of NaOH production is directed toward this sector, facilitating the separation of fibrous material while minimizing fiber degradation. Water treatment relies on lime, or (Ca(OH)₂), for adjustment and softening, where it precipitates calcium and magnesium as carbonates, raising to around 10 to enhance and removal in municipal supplies. The economic scale of base production underscores their industrial importance, with global NaOH output reaching about 83 million tons annually as of 2024, driven primarily by the chlor-alkali process and supporting sectors like chemicals, pulp, and textiles. In laboratory settings, bases are routinely employed to prepare buffers, such as those combining NaOH with weak s like acetic acid, which stabilize reaction conditions in biochemical assays and maintain optimal environments for activity. -base extractions utilize bases like NaOH to deprotonate acidic organic compounds, converting them to water-soluble salts for separation from non-polar solvents, a technique central to purifying natural products and pharmaceuticals. For , bases drive ester hydrolysis, as in the conversion of to and under NaOH , providing a mild route to alcohols and acids without harsh acidic conditions. Neutralization reactions with bases also form the basis of analyses for quantifying acids in samples. Environmental considerations in base utilization focus on managing alkaline wastes, which arise from processes like pulping and chlor-alkali production; these high-pH effluents can raise receiving water , potentially harming aquatic life, so treatment often involves neutralization with acids or dilution to comply with discharge standards. Effective strategies, including recovery of byproducts like from , reduce ecological footprints and support circular economies in base-intensive industries.

Historical Context

Etymology

The term "base" in chemistry originates from the Latin basis, meaning "foundation" or "pedestal," which entered English via base around the 14th century to denote a supporting structure. In chemical contexts, it was first applied by French chemist Louis Lémery in 1717, who used it as a for the Paracelsian alchemical of a "matrix"—a substance capable of combining with acids to form compounds, analogous to a foundational element in reactions. This early usage emphasized the role of such substances in stabilizing or "supporting" acidic components, much like a base supports a structure. The modern chemical meaning solidified through the work of Guillaume-François Rouelle, who in 1744 described "bases" as diverse substances—such as , lime, or metallic oxides—that react with acids to produce neutral salts, distinguishing them from fixed alkalis. In 1754, Rouelle proposed a of salts into neutral, acidic, and basic types, further developing the concept of bases as substances that react with acids. A related historical term, "alkali," frequently overlapped with "base" in early usage and derives from the Arabic al-qaliy (الْقِلْي), meaning "the calcined ashes," referring to alkaline extracts from burned plants or wood ashes used in soap-making and dyeing since ancient times. Over time, the terminology evolved with advances in acid-base theory; by the early , following Søren Sørensen's introduction of the scale in 1909, solutions with pH greater than 7 became designated as "basic," denoting their opposition to acidic conditions (pH less than 7) and alignment with hydroxide-producing properties. This shift integrated the foundational etymological sense of "base" into quantitative measures of , standardizing its use in modern chemistry.

Evolution of the Concept

The concept of bases in chemistry originated from empirical observations in ancient civilizations, where substances like lime (calcium oxide) and wood ashes (containing ) were recognized for their ability to neutralize sour tastes and react with fats to produce soaps. Lime was used as a and mortar as early as 7000 BCE in the , with evidence from sites indicating its slaking to form for plasters and coatings. Wood ashes, leached to produce (), were employed in soap-making and cleaning by cultures such as the ancient Babylonians and , highlighting early practical understanding of basic properties without theoretical framework. In the late , proposed a linking acidity to the presence of oxygen, suggesting that acids formed by combining oxygen with non-metallic elements, while bases were oxides of metals that could neutralize them. This oxygen , articulated around 1777, represented an early systematic attempt to classify chemical behaviors but was later disproven by the discovery of as an oxygen-free acid in 1810. The focus remained on observable reactions rather than underlying mechanisms until the . Svante Arrhenius advanced the understanding in 1884 with his ionic theory, defining bases as substances that dissociate in aqueous solution to produce hydroxide ions (OH⁻), thereby increasing the solution's basicity. This water-centric model explained neutralization as the combination of H⁺ from acids and OH⁻ from bases to form water, earning Arrhenius the 1903 Nobel Prize in Chemistry for his contributions to electrolyte dissociation. The early 20th century saw broader definitions emerge independently in 1923: and described bases as proton (H⁺) acceptors in any proton-transfer reaction, extending beyond aqueous contexts to include gas-phase and non-aqueous systems. Concurrently, proposed that bases are electron-pair donors, forming coordinate bonds with electron-pair acceptors (acids), which unified acid-base interactions with valence theory and applied to coordination chemistry. Post-1950s developments addressed solvent limitations of earlier theories by establishing solvent-independent basicity scales, such as pKa measurements in (DMSO), pioneered by Frederick G. Bordwell in the 1960s–1980s to quantify acid strengths of weak bases across organic solvents. These scales enabled comparisons of intrinsic basicity without water's leveling effect. Since the 1960s, quantum chemical models have further extended basicity concepts, notably through Ralph G. Pearson's Hard-Soft Acid-Base (, which classifies bases by their and to predict reactivity preferences with acids, influencing modern computational predictions of basicity.

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