Base (chemistry)

In chemistry, bases are substances that, in aqueous solution, release hydroxide (OH) ions, are slippery to the touch, can taste bitter if an alkali,[1] change the color of indicators (e.g., turn red litmus paper blue), react with acids to form salts, promote certain chemical reactions (base catalysis), accept protons from any proton donor or contain completely or partially displaceable OH ions. Examples of bases are the hydroxides of the alkali metals and the alkaline earth metals (NaOH, Ca(OH)2, etc.—see alkali hydroxide and alkaline earth hydroxide).

These particular substances produce hydroxide ions (OH) in aqueous solutions, and are thus classified as Arrhenius bases. For a substance to be classified as an Arrhenius base, it must produce hydroxide ions in an aqueous solution. Arrhenius believed that in order to do so, the base must contain hydroxide in the formula. This makes the Arrhenius model limited, as it cannot explain the basic properties of aqueous solutions of ammonia (NH3) or its organic derivatives (amines).[2] There are also bases that do not contain a hydroxide ion but nevertheless react with water, resulting in an increase in the concentration of the hydroxide ion.[3] An example of this is the reaction between ammonia and water to produce ammonium and hydroxide.[3] In this reaction ammonia is the base because it accepts a proton from the water molecule.[3] Ammonia and other bases similar to it usually have the ability to form a bond with a proton due to the unshared pair of electrons that they possess.[3] In the more general Brønsted–Lowry acid–base theory, a base is a substance that can accept hydrogen cations (H+)—otherwise known as protons. In the Lewis model, a base is an electron pair donor.[4]

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. However, it is important to realize that basicity is not the same as alkalinity. Metal oxides, hydroxides, and especially alkoxides are basic, and counteranions of weak acids are weak bases.

Bases can be thought of as the chemical opposite of acids. However, some strong acids are able to act as bases.[5] Bases and acids are seen as opposites because the effect of an acid is to increase the hydronium (H3O+) concentration in water, whereas bases reduce this concentration. A reaction between an acid and base is called neutralization. In a neutralization reaction, an aqueous solution of a base reacts with an aqueous solution of an acid to produce a solution of water and 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.

The notion of a base as a concept in chemistry was first introduced by the French chemist Guillaume François Rouelle in 1754. He noted that acids, which at that time were mostly volatile liquids (like acetic acid), turned into solid salts only when combined with specific substances. Rouelle considered that such a substance serves as a "base" for the salt, giving the salt a "concrete or solid form".[6]

Decorative Soaps
Soaps are weak bases formed by the reaction of fatty acids with sodium hydroxide or potassium hydroxide.


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 in taste.[7]

Reactions between bases and water

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]

B(aq) + H2O(l) ⇌ BH+(aq) + OH(aq)

The equilibrium constant, Kb, for this reaction can be found using the following general equation:[3]

Kb = [BH+][OH]/[B]

In this equation, both the base (B) and the extremely strong base (the conjugate base) compete with one another for the proton.[8] As a result, bases that react with water have relatively small equilibrium constant values.[8] The base is weaker when it has a lower equilibrium constant value.[3]

Neutralization of acids

Hydrochloric acid ammonia
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.[5] When dissolved in water, the strong base sodium hydroxide ionizes into hydroxide and sodium ions:

NaOH → Na+
+ OH

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

HCl + H
+ Cl

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

+ OH
→ 2 H

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

Bases are generally compounds that can neutralize an amount of acids. 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:

Na2CO3 + H2O → 2 Na+ + HCO3 + OH
NH3 + H2O → NH4+ + OH

From this, a pH, or acidity, can be calculated for aqueous solutions of bases. Bases also directly act as electron-pair donors themselves:

CO32− + H+ → HCO3
NH3 + H+ → NH4+

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.[5] There are a limited number of elements that have atoms with the ability to provide a molecule with basic properties.[5] Carbon can act as a base as well as nitrogen and oxygen. Fluorine and sometimes rare gases possess this ability as well.[5] 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

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 NaOH and Ca(OH)
, 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.[9] One advantage of this low solubility is that "many antacids were suspensions of metal hydroxides such as aluminum hydroxide and magnesium hydroxide."[10] These compounds have low solubility and have 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.[10] As the reaction continues and the salts dissolve, the stomach acid reacts with the hydroxide produced by the suspensions.[10] Strong bases hydrolyze in water almost completely, resulting in the leveling effect."[5] In this process, the water molecule combines with a strong base, due to the water's amphoteric ability; and, a hydroxide ion is released.[5] 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:

The cations of these strong bases appear in the first and second groups of the periodic table (alkali and earth alkali metals).

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


Group 1 salts of carbanions, amides, 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 water solution because they are stronger bases than the hydroxide ion. As such, they deprotonate the conjugate acid water. For example, the ethoxide ion (conjugate base of ethanol) in the presence of water undergoes this reaction.

+ H
+ OH

Examples of superbases are:

Neutral bases

When a neutral base forms a bond with a neutral acid, a condition of electric stress occurs.[5] The acid and the base share the electron pair that formerly only belonged to the base.[5] As a result, a high dipole moment is created, which can only be destroyed by rearranging the molecules.[5]

Weak bases

A weak base is one which does not fully ionize in an aqueous solution, or in which protonation is incomplete.

Solid bases

Examples of solid bases include:

  • Oxide mixtures: SiO2, Al2O3; MgO, SiO2; CaO, SiO2[11]
  • Mounted bases: LiCO3 on silica; NR3, NH3, KNH2 on alumina; NaOH, KOH mounted on silica on alumina[11]
  • Inorganic chemicals: BaO, KNaCO3, BeO, MgO, CaO, KCN[11]
  • Anion exchange resins[11]
  • Charcoal that has been treated at 900 degrees Celsius or activates with N2O, NH3, ZnCl2-NH4Cl-CO2[11]

Depending on a solid surface's ability to successfully form a conjugate base by absorbing an electrically neutral acid, the basic strength of the surface is determined.[12] "The number of basic sites per unit surface area of the solid" is used to express how much base is found on a solid base catalyst.[12] Scientists have developed two methods to measure the amount of basic sites: titration with benzoic acid using indicators and gaseous acid adsorption.[12] A solid with enough basic strength will absorb an electrically neutral acid indicator and cause the acid indicator's color to change to the color of its conjugate base.[12] When performing the gaseous acid adsorption method, nitric oxide is used.[12] The basic sites are then determined using the amount of carbon dioxide than is absorbed.[12]

Bases as catalysts

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 have been used for hydrogenations, the migration of double bonds, in the Meerwein-Ponndorf-Verley reduction, the Michael reaction, and many other reactions. Both CaO and BaO can be highly active catalysts if they are treated with high temperature heat.[12]

Uses of bases

  • Sodium hydroxide is used in 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 exhaust, that is found in power plants and factories.[10]
  • 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 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

Acidity of bases

The number of ionizable hydroxide (OH-) ions present in one molecule of base is called the acidity of bases.[13] On the basis of acidity bases can be classified into three types: monoacidic, diacidic and triacidic.

Monoacidic bases

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

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

Diacidic bases

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

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

Triacidic bases

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

Aluminium hydroxide, ferrous hydroxide, Gold Trihydroxide,

Etymology of the term

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. ... 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 Jensen, The origin of the term "base"[6]

See also


  1. ^ Johll, Matthew E. (2009). Investigating chemistry: a forensic science perspective (2nd ed.). New York: W. H. Freeman and Co. ISBN 1429209895. OCLC 392223218.
  2. ^ Whitten et al. (2009), p. 363.
  3. ^ a b c d e f g Zumdahl & DeCoste (2013), p. 257.
  4. ^ Whitten et al. (2009), p. 349.
  5. ^ a b c d e f g h i j Lewis, Gilbert N. (1938). "Acids and Bases" (PDF). Journal of the Franklin Institute. pp. 293–313. Retrieved 19 February 2015.
  6. ^ a b Jensen, William B. (2006). "The origin of the term 'base'" (PDF). The Journal of Chemical Education. 83 (8): 1130. Bibcode:2006JChEd..83.1130J. doi:10.1021/ed083p1130. Archived from the original (PDF) on 4 March 2016.
  7. ^ "Definition of BASE". Archived from the original on 21 March 2018. Retrieved 3 May 2018.
  8. ^ a b Zumdahl & DeCoste (2013), p. 258.
  9. ^ Zumdahl & DeCoste (2013), p. 255.
  10. ^ a b c d Zumdahl & DeCoste (2013), p. 256.
  11. ^ a b c d e Tanabe, Kozo (1970). Solid Acids and Bases: their catalytic properties. Academic Press. p. 2. Retrieved 19 February 2015.
  12. ^ a b c d e f g Tanabe, K.; Misono, M.; Ono, Y.; Hattori, H. (1990). New Solid Acids and Bases: their catalytic properties. Elsevier. p. 14. Retrieved 19 February 2015.
  13. ^ "Electrophile - Nucleophile - Basicity - Acidity - pH Scale". City Collegiate. Archived from the original on 30 June 2010. Retrieved 20 June 2016.
  • Whitten, Kenneth W.; Peck, Larry; Davis, Raymond E.; Lockwood, Lisa; Stanley, George G. (2009). Chemistry (9th ed.). ISBN 0-495-39163-8.
  • Zumdahl, Steven; DeCoste, Donald (2013). Chemical Principles (7th ed.). Mary Finch.
Acid–base reaction

An acid–base reaction is a chemical reaction that occurs between an acid and a base, which can be used to determine pH. Several theoretical frameworks provide alternative conceptions of the reaction mechanisms and their application in solving related problems; these are called the acid–base theories, for example, Brønsted–Lowry acid–base theory.

Their importance becomes apparent in analyzing acid–base reactions for gaseous or liquid species, or when acid or base character may be somewhat less apparent. The first of these concepts was provided by the French chemist Antoine Lavoisier, around 1776.

It is important to think of the acid-base reaction models as theories that complement each other. For example the current Lewis model has the broadest definition of what an acid and base are, with the Bronsted-Lowry theory being a subset of what acids and bases are, and the Arrhenius theory being the most restrictive.


Alkalinity (from Arabic "al-qalī") is the capacity of water to resist changes in pH that would make the water more acidic. (It should not be confused with basicity which is an absolute measurement on the pH scale.) Alkalinity is the strength of a buffer solution composed of weak acids and their conjugate bases. It is measured by titrating the solution with a monoprotic acid such as HCl until its pH changes abruptly, or it reaches a known endpoint where that happens. Alkalinity is expressed in units of meq/L (milliequivalents per liter), which corresponds to the amount of monoprotic acid added as a titrant in millimoles per liter.

Although alkalinity is primarily a term invented by oceanographers, it is also used by hydrologists to describe temporary hardness. Moreover, measuring alkalinity is important in determining a stream's ability to neutralize acidic pollution from rainfall or wastewater. It is one of the best measures of the sensitivity of the stream to acid inputs. There can be long-term changes in the alkalinity of streams and rivers in response to human disturbances.


In chemistry, an amphoteric compound is a molecule or ion that can react both as an acid and as a base. Many metals (such as copper, zinc, tin, lead, aluminium, and beryllium) form amphoteric oxides or hydroxides. Amphoterism depends on the oxidation states of the oxide. Al2O3 is an example of an amphoteric oxide.

The prefix of the word 'amphoteric' is derived from a Greek prefix amphi-, which means both. In chemistry, an amphoteric substance is a substance that has the ability to act either as an acid or a base. Remember that acids donate protons (or accept electron pairs) and bases accept protons. Amphoteric substances can do either.

Metal oxides which react with both acids as well as bases to produce salts and water are known as amphoteric oxides. Amphoteric oxides include lead oxide and zinc oxide, among many others.

One type of amphoteric species are amphiprotic molecules, which can either donate or accept a proton (H+). Examples include amino acids and proteins, which have amine and carboxylic acid groups, and self-ionizable compounds such as water.

Ampholytes are amphoteric molecules that contain both acidic and basic groups and will exist mostly as zwitterions in a certain range of pH. The pH at which the average charge is zero is known as the molecule's isoelectric point.

Ampholytes are used to establish a stable pH gradient for use in isoelectric focusing.

Buffer solution

A buffer solution (more precisely, pH buffer or hydrogen ion buffer) is an aqueous solution consisting of a mixture of a weak acid and its conjugate base, or vice versa. Its pH changes very little when a small amount of strong acid or base is added to it. Buffer solutions are used as a means of keeping pH at a nearly constant value in a wide variety of chemical applications. In nature, there are many systems that use buffering for pH regulation. For example, the bicarbonate buffering system is used to regulate the pH of blood.

Buffering agent

A buffering agent is a weak acid or base used to maintain the acidity (pH) of a solution near a chosen value after the addition of another acid or base. That is, the function of a buffering agent is to prevent a rapid change in pH when acids or bases are added to the solution. Buffering agents have variable properties—some are more soluble than others; some are acidic while others are basic. As pH managers, they are important in many chemical applications, including agriculture, food processing, biochemistry, medicine and photography.

Chemical field-effect transistor

See also ISFET

A ChemFET is a chemically-sensitive field-effect transistor, that is a field-effect transistor used as a sensor for measuring chemical concentrations in solution. When the target analyte concentration changes, the current through the transistor will change accordingly. Here, the analyte solution separates the source and gate electrodes. A concentration gradient between the solution and the gate electrode arises due to a semi-permeable membrane on the FET surface containing receptor moieties that preferentially bind the target analyte. This concentration gradient of charged analyte ions creates a chemical potential between the source and gate, which is in turn measured by the FET.

Conjugate acid

A conjugate acid, within the Brønsted–Lowry acid–base theory, is a species formed by the reception of a proton (H+) by a base—in other words, it is a base with a hydrogen ion added to it. On the other hand, a conjugate base is what is left over after an acid has donated a proton during a chemical reaction. Hence, a conjugate base is a species formed by the removal of a proton from an acid. Because some acids are capable of releasing multiple protons, the conjugate base of an acid may itself be acidic.

In summary, this can be represented as the following chemical reaction:

Acid + Base ⇌ Conjugate Base + Conjugate Acid

Johannes Nicolaus Brønsted and Martin Lowry introduced the Brønsted–Lowry theory,

which proposed that any compound that can transfer a proton to any other compound is an acid, and the compound that accepts the proton is a base. A proton is a nuclear particle with a unit positive electrical charge; it is represented by the symbol H+ because it constitutes the nucleus of a hydrogen atom, that is, a hydrogen cation.

A cation can be a conjugate acid, and an anion can be a conjugate base, depending on which substance is involved and which acid–base theory is the viewpoint. The simplest anion which can be a conjugate base is the solvated electron whose conjugate acid is the atomic hydrogen.

Coordinate covalent bond

A coordinate covalent bond, also known as a dative bond or coordinate bond is a kind of 2-center, 2-electron covalent bond in which the two electrons derive from the same atom. The bonding of metal ions to ligands involves this kind of interaction. This type of interaction is central to Lewis acid-base theory.


Deprotonation is the removal (transfer) of a proton (a hydrogen cation, H+) from a Brønsted–Lowry acid in an acid-base reaction. 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. 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 things, 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.

Electron donor

An electron donor is a chemical entity that donates electrons to another compound. It is a reducing agent that, by virtue of its donating electrons, is itself oxidized in the process.

Typical reducing agents undergo permanent chemical alteration through covalent or ionic reaction chemistry. This results in the complete and irreversible transfer of one or more electrons. In many chemical circumstances, however, the transfer of electronic charge to an electron acceptor may be only fractional, meaning an electron is not completely transferred, but results in an electron resonance between the donor and acceptor. This leads to the formation of charge transfer complexes in which the components largely retain their chemical identities.

The electron donating power of a donor molecule is measured by its ionization potential which is the energy required to remove an electron from the highest occupied molecular orbital.

The overall energy balance (ΔE), i.e., energy gained or lost, in an electron donor-acceptor transfer is determined by the difference between the acceptor's electron affinity (A) and the ionization potential (I):

In chemistry, the class of electron donors that donate not just one, but a set of two paired electrons that form a covalent bond with an electron acceptor molecule, is known as a Lewis base. This phenomenon gives rise to the wide field of Lewis acid-base chemistry. The driving forces for electron donor and acceptor behavior in chemistry is based on the concepts of electropositivity (for donors) and electronegativity (for acceptors) of atomic or molecular entities.

Free base

Free base (freebase, free-base) is the conjugate base (deprotonated) form of an amine, as opposed to its conjugate acid (protonated) form. The amine is often an alkaloid, such as nicotine, cocaine, morphine, and ephedrine, or derivatives thereof.

Hammett acidity function

The Hammett acidity function (H0) is a measure of acidity that is used for very concentrated solutions of strong acids, including superacids. It was proposed by the physical organic chemist Louis Plack Hammett and is the best-known acidity function used to extend the measure of Brønsted–Lowry acidity beyond the dilute aqueous solutions for which the pH scale is useful.

In highly concentrated solutions, simple approximations such as the Henderson–Hasselbalch equation are no longer valid due to the variations of the activity coefficients. The Hammett acidity function is used in fields such as physical organic chemistry for the study of acid-catalyzed reactions, because some of these reactions use acids in very high concentrations, or even neat (pure).

Henderson–Hasselbalch equation

In chemistry and biochemistry, the Henderson–Hasselbalch equation can be used to estimate the pH of a buffer solution containing given concentrations of an acid and its conjugate base (or a base and its conjugate acid). The numerical value of the acid dissociation constant of the acid must also be known.

Lewis acids and bases

A Lewis acid is a chemical species that contains an empty orbital which is capable of accepting an electron pair from a Lewis base to form a Lewis adduct. A Lewis base, then, is any species that has a filled orbital containing an electron pair which is not involved in bonding but may form a dative bond with a Lewis acid to form a Lewis adduct. For example, NH3 is a Lewis base, because it can donate its lone pair of electrons. Trimethylborane (Me3B) is a Lewis acid as it is capable of accepting a lone pair. In a Lewis adduct, the Lewis acid and base share an electron pair furnished by the Lewis base, forming a dative bond. In the context of a specific chemical reaction between NH3 and Me3B, the lone pair from NH3 will form a dative bond with the empty orbital of Me3B to form an adduct NH3•BMe3. The terminology refers to the contributions of Gilbert N. Lewis.

Neutralization (chemistry)

In chemistry, neutralization or neutralisation (see spelling differences) is a chemical reaction in which an acid and a base react quantitatively with each other. In a reaction in water, neutralization results in there being no excess of hydrogen or hydroxide ions present in the solution. The pH of the neutralized solution depends on the acid strength of the reactants. Neutralization is used in many applications.


The pCO2, PCO2, or is the partial pressure of carbon dioxide (CO2), often used in reference to blood, but also used in oceanography to describe the partial pressure of CO2 in the Ocean, and in life support systems engineering and underwater diving to describe the partial pressure in a breathing gas. Usually the arterial blood is the relevant context; the symbol for in arterial blood is . Measurement of in the systemic circulation indicates the effectiveness of ventilation at the lungs' alveoli, given the diffusing capacity of the gas. It is a good indicator of respiratory function and the closely related factor of acid–base homeostasis, reflecting the amount of acid in the blood (without lactic acid).

PH meter

A pH meter is a scientific instrument that measures the hydrogen-ion activity in water-based solutions, indicating its acidity or alkalinity expressed as pH. The pH meter measures the difference in electrical potential between a pH electrode and a reference electrode, and so the pH meter is sometimes referred to as a "potentiometric pH meter". The difference in electrical potential relates to the acidity or pH of the solution. The pH meter is used in many applications ranging from laboratory experimentation to quality control.

Polyatomic ion

A polyatomic ion, also known as a molecular ion, is a charged chemical species (ion) composed of two or more atoms covalently bonded or of a metal complex that can be considered to be acting as a single unit. The prefix poly- means "many," in Greek, but even ions of two atoms are commonly referred to as polyatomic. In older literature, a polyatomic ion is also referred to as a radical, and less commonly, as a radical group. In contemporary usage, the term radical refers to free radicals that are (not necessarily charged) species with an unpaired electron.

An example of a polyatomic ion is the hydroxide ion; consisting of one oxygen atom and one hydrogen atom, hydroxide has a charge of −1. Its chemical formula is OH−. An ammonium ion is made up of one nitrogen atom and four hydrogen atoms: it has a charge of +1, and its chemical formula is NH+4.

Polyatomic ions are often useful in the context of acid-base chemistry or in the formation of salts. A polyatomic ion can often be considered as the conjugate acid/base of a neutral molecule. For example, the conjugate base of sulfuric acid (H2SO4) is the polyatomic hydrogen sulfate anion (HSO−4). The removal of another hydrogen ion yields the sulfate anion (SO2−4).

Solid acid

Solid acids are acids that do not dissolve in the reaction medium. They are often used in heterogeneous catalysts.

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