An acid-base reaction is a chemical reaction that occurs between an acid and a base . It can be used to determine pH . Several theoretical frameworks provide alternative concepts of reaction mechanisms and their application in solving related problems; These are called acid-base theories , for example, the Brnsted–Lowry acid–base theory .
Their importance becomes clear in analyzing acid–base reactions for gaseous or liquid species, or when the acid or base character may be somewhat less pronounced. The first of these concepts was provided by the French chemist Antoine Lavoisier around 1776.
It is important to think of acid-base reaction models as principles that complement each other.  For example, the current Lewis model has a broad definition of what acids and bases are, with the Brnsted–Lowry theory of what acids and bases are, and the Arrhenius theory being the most restrictive.
The concept of an acid-base reaction was first proposed in 1754 by Guillaume-François Roule , who introduced the term “base” to chemistry, meaning a substance that reacts with an acid to give it a solid form. (as salt).
Lavoisier’s Oxygen Theory of Acids
acids and bases was provided by Voisier since Voisier’s knowledge was restricted to strong acids mainly oxoacids , such as HNO , around 1776. 3(nitric acid) and H2So
4(sulfuric acid), which has central atoms in higher oxidation states surrounded by oxygen , and since he was not aware of the exact structure of hydrohalic acids ( HF , HCl , HBr , and HI ), he defined acids as their containing oxygen , which he actually named from the Greek words “acid-former” ( from Greek ( oxy ) meaning “acid” or “sharp” and μαι ( genomai ) meaning ” engender “). The Lavoisier definition lasted more than 30 years, until an 1810 article and by Sir Humphrey Davy.He held until later lectures in which he proved the reduction of oxygen in H.
2S , H 2 T , and hydrohalic acid . However, Davy failed to develop a new theory, concluding that “acidity does not depend on any particular elementary substance, but on a strange arrangement of various substances”.  A notable revision of the oxygen theory was provided by Jöns Jacob Berzelius , who postulated that acids are oxides of nonmetals while alkalis are oxides of metals.
Liebig’s hydrogen theory of acids
In 1838, Justus von Liebig proposed that an acid is a hydrogen-containing compound whose hydrogen could be replaced by a metal.    This redefinition was based on his extensive work on the chemical structure of organic acids , ending the theoretical shift from oxygen-based acids to hydrogen-based acids introduced by Davy. Liebig’s definition, while completely empirical, remained in use for almost 50 years until the adoption of the Arrhenius definition.
The first modern definitions of acid and base in molecular terms were formulated by Svante Arrhenius .   A hydrogen theory of acids, it followed his 1884 work with Friedrich Wilhelm Ostwald establishing the presence of ions in aqueous solution , and Arrhenius received the Nobel Prize in Chemistry in 1903 .
As defined by Arrhenius:
An Arrhenius acid is a substance that dissociates in water to form hydrogen ions ( H+ ) ;  That is, an acid increases the concentration of H + ions in aqueous solution.
This causes the formation of water protons or the hydronium (H 3 O + ) ion. [note 1] Thus, in modern times, the symbol H + is interpreted as a shorthand for H3O + , as it is now known that a bare proton exists as a free species in aqueous solution . Not there. 
- An Arrhenius base is a substance that dissociates in water to form hydroxide (OH – ) ions; That is, an alkali increases the concentration of OH- ions in aqueous solution .
Arrhenius definitions of acidity and alkalinity are restricted to aqueous solutions, and refer to the concentration of solvent ions. Under this definition, pure H2SO4 and HCl dissolved in toluene are not acidic , and solutions of molten NaOH and calcium amide in liquid ammonia are not alkaline. This led to the development of the Brnsted–Lowry principle and later the Lewis theory to account for these non-aquatic exceptions. 
Overall, to qualify as an Arrhenius acid, upon introduction to water, the chemical must either directly or otherwise cause:
- an increase in aqueous hydronium concentration, or
- decrease in aqueous hydroxide concentration.
In contrast, to qualify as an Arrhenius base, upon introduction to water, the chemical must either directly or otherwise cause:
- a decrease in aqueous hydronium concentration, or
- Increase in aqueous hydroxide concentration.
The reaction of acid and base is called neutralization reaction. The products of this reaction are salt and water.Acid + Base → Salt + Water
In this conventional representation an acid-base neutralization reaction is modeled as a double-replacement reaction. For example, the reaction of hydrochloric acid, HCl, sodium hydroxide, NaOH, solutions forms a solution of sodium chloride, NaCl and some additional water molecules.HCl(aq) + NaOH(aq) → NaCl(aq) + H 2 O
Rather than explicitly including the modifier ( aq ) in this equation, it was implied by Arrhenius. This indicates that the substances have dissolved in water. Although all three substances, HCl, NaOH and NaCl, are able to exist as pure compounds, in aqueous solution they completely dissociate into the aqueous ions H + , Cl- , Na + and OH- .
The Brnsted–Lowry definition, independently formulated in 1923 by Johannes Nicolaus Brnstedt in Denmark and Martin Lowry in England,   is based on the idea of protonation of bases via the precipitate of acids—that is, the potential of The ability of acids to “donate” hydrogen ions (H + ) – otherwise known as protons – to bases, which “accept” them.  [note 2]
An acid-base reaction, thus, is the removal of a hydrogen ion from an acid and its addition to a base.  The removal of a hydrogen ion from an acid forms its conjugate base , which is an acid that removes a hydrogen ion. The reception of a proton by a base produces its conjugate acid , which is a base with a hydrogen ion.
Unlike the previous definitions, the Brnsted–Lowry definition does not refer to the formation of the salt and solvent, but to the formation of conjugate acids and conjugate bases produced by the transfer of a proton from acid to base .   In this approach, acids and bases are fundamentally different in behavior from salts, which are viewed as electrolytes, subject to the principles of Debbie, Onsagar, and others. An acid and a base react not to produce a salt and a solvent, but to form a new acid and a new base. Thus the concept of neutrality is absent. Brnsted–Lowry acid–base behavior is formally independent of any solvent, making it more comprehensive than the Arrhenius model. The calculation of pH under the Arrhenius model depends on the base dissolved in water (aqueous solution). The Brnsted–Lowry model allows pH testing to be performed using insoluble and soluble solutions (gas, liquid, solid).
The general formula for acid-base reactions according to the Brnsted–Lowry definition is:Ha + b → BH + + a −
Where HA represents the acid, B represents the base, BH + represents the conjugate acid of B, and A represents the conjugate base of HA .
For example, the Bronsted–Lowry model for the dissociation of hydrochloric acid (HCl) in aqueous solution would be the following:HCl + H 2 O H 3 O + + Cl –
Removal of H + from HCl makes the chloride ion, Cl−, the conjugate base of the acid. Adding H + to H 2 O (acting as a base) forms the conjugate acid of the hydronium ion, H 3 O + , the base.
Water is amphoteric i.e. it can act as both an acid and a base. The Bronsted–Lowry model explains this by showing the dissociation of water at low concentrations of hydronium and hydroxide ions:H 2 O + H 2 O H 3 O + + OH –
This equation is displayed in the image below:
Here, one molecule of water acts as an acid, donating an H + and forming the conjugate base, OH- , and the second molecule of water acts as a base, accepting the H + ion and forming the conjugate base of the acid. Formation, H 3 O + .
Consider an aqueous solution of pyridine, C 5 H 5 N, as an example of water acting as an acid.C 5 H 5 N + H 2 O O [C 5 H 5 NH] + + OH –
In this example, a water molecule is split into a hydrogen ion donated to a pyridine molecule, and a hydroxide ion.
In the Brnsted–Lowry model, the solvent does not have to be water, as required by the Arrhenius acid–base model. For example, consider what happens when acetic acid, CH3COOH , dissolves in liquid ammonia.Chieftain
3COOH + NH
An H + ion is removed from acetic acid, forming its conjugate base, the acetate ion , CH3COO− . Addition of H + ion to ammonia molecule of solvent forms its conjugate acid, ammonium ion, NH .+4.
The Brnsted-Lowry model calls hydrogen-containing substances (such as HCl) acids. Thus, some substances, which are considered by many chemists to be acids, such as SO 3 or BCl 3 , are out of this classification due to their lack of hydrogen. Gilbert N. Lewis wrote in 1938, “To limit the group of acids to substances that contain hydrogen interferes with the systematic understanding of chemistry as seriously as the restriction of the term oxidizing agent to oxygen-containing substances.” ”  In addition, KOH and KNH are not considered to be Brnsted bases, but rather OH- and NH2-based salts.
Arrhenius and Brnsted–Lowry’s hydrogen requirement was superseded by the Lewis definition of acid–base reactions, which was introduced in 1923 by Gilbert N. Lewis,  in the same year as Brnsted-Lowry, but it was not elaborated upon by him until 1938.  Acid–base reactions in terms of protons or other bonded substances Rather than being defined, the Lewis definition defines a base ( referred to as a Lewis base ) as a compound that can donate an electron pair , and an acid (a Lewis acid ) as being a compound. for which can acquire this electron pair. 
For example, boron trifluoride, BF3 is a typical Lewis acid. It can accept a pair of electrons because it has a vacancy in its octet. The fluoride ion is a complete octet and can donate a pair of electrons. thus bf 3 + f – → bf–4
A typical Lewis acid is a Lewis base reaction. All compounds of group 13 elements having the formula AX 3 can behave as Lewis acids. Similarly, compounds of group 15 elements with the formula DY3 , such as amines, NRs, and phosphines, PR3, can behave as Lewis bases. The adducts between them have the formula X 3 A←DY 3 with a basic covalent bond , shown symbolically as between atoms A (acceptor) and D (donor). Group 16 compounds with the formula DX 2 can also act as Lewis bases; Thus, ether, r 2 o, or thioether, r 2A compound like S can act as a Lewis base. Lewis’ definition is not limited to these examples. For example, carbon monoxide acts as a Lewis base when it forms a compound with boron trifluoride of the formula F3B←CO .
Adhesions involving metal ions are known as coordination compounds; Each ligand donates a pair of electrons to the metal ion.
[Ag (H 2 O) 4 ] + + 2NH 3 → [Ag (NH 3 ) 2 ] + + 4H 2 O
Can be viewed as an acid-base reaction in which a strong base (ammonia) replaces a weak one (water)
The Lewis and Brnsted–Lowry definitions are consistent with each other since the reaction H + + OH – H 2 O
There is an acid-base reaction in both theories.
Solvent System Definition
One of the limitations of the Arrhenius definition is its dependence on water solutions. Edward Curtis Franklin studied acid-base reactions in liquid ammonia in 1905 and pointed out similarities to the water-based Arrhenius theory. Albert FO German, working with liquid phosgene, COCl
2, formulated the solvent-based theory in 1925, generalizing the Arrhenius definition to cover aprotic solvents. 
German pointed out that many solutions contain ions in equilibrium with neutral solvent molecules:
- Solvonium ion: A common name for positive ions. (The term solvonium has replaced the older term lionium ions : positive ions formed by protonation of solvent molecules.)
- Solvate ions: A common name for negative ions. (the term solvate is replaced by the larger term lyate ions : negative ions formed by deprotonation of solvent molecules.)
For example, water and ammonia undergo such dissociation into hydronium and hydroxide, and ammonium and amide, respectively:
2 H2O ⇌ H3O++ OH−
2 NH3 ⇌ NH+4 + NH−2
Some aprotic systems also undergo dissociation, such as dinitrogen tetroxide in nitrosonium and nitrate, antimony trichloride in dichloroantimonium and tetrachloroantimonate, and chlorocarboxonium and in cyanogen in chloride:
N2O4 ⇌ NO++ NO−3
2 SbCl3 ⇌ SbCl+2 + SbCl−4
COCl2 ⇌ COCl++ Cl−
A solute that causes an increase in the concentration of solvonium ions and a decrease in the concentration of solvate ions is defined as an acid . A solute that causes an increase in the concentration of solvate ions and a decrease in the concentration of solvonium ions is defined as a base .
Thus, KNH in liquid ammonia2( Supply of NH-2) is a strong base, and NH4No3( Supply of NH+4) is a strong acid. In liquid sulfur dioxide ( SO.)2), thionyl compound ( SO . SUPPLIED )2+) behave as acids and sulfites ( supply SO. )2−3) as the basis.
The non-aqueous acid-base reactions in liquid ammonia are similar to those in water:
2 NaNH2(base)+ Zn(NH2)2(amphiphilic amide)→ Na2[Zn(NH2)4]
2 NH4I(acid)+Zn ( NH2)2(amphiphilic amide)→ [ zn ( nh3)4]I2
Nitric acid can be a base in liquid sulfuric acid:
HNO3(base)+ 2 H2SO4 → NO2 + H3O+ + 2HSO4
The unique strength of this definition shows in describing the reactions in aprotic solvents; for example, in liquid N2O4:
Because the definition of the solvent system depends on the solute as well as the solvent, a particular solute can be either an acid or a base, depending on the choice of solvent: HClO
4Water is a strong acid, acetic acid has a weak acid and fluorosulfonic acid has a weak base; This feature of the theory has been seen as both a strength and a weakness, as some substances (such as SO.)
3) itself has been observed to be acidic or basic. On the other hand, solvent system theory has been criticized as being too general to be useful. In addition, it is thought that there is something intrinsically acidic about hydrogen compounds, a property not shared by non-hydrogen solvonium salts. 
This acid-base theory was a revival of the oxygen theory of acids and bases, proposed in 1939 by German chemist Hermann Lux   , which was further improved by Höcken Flood around 1947  and Still used in modern geochemistry and electrochemistry. of molten salt. This definition describes an acid as an oxide ion ( O.)2-
) a base as the acceptor and an oxide ion donor. For example:
MgO (base) + CO2 (acid) → MgCO3
CaO (base) + SiO2 (acid) → CaSiO3
NO−3 (base) + S2O2−7 (acid) → NO+ O2 + 2 SO2−4
This principle is also useful in the systematization of reactions of noble gas compounds, especially xenon oxide, fluoride and oxofluoride. 
Mikhail Usanovich developed a general theory that does not limit acidity to hydrogen-containing compounds, but his approach, published in 1938, was even more general than Lewis’ theory.  Usanovich’s principle can be summarized as defining an acid that accepts negative species or donates positive ones, and a base as the reverse. It defined the concept of redox (oxidation-reduction) as a special case of acid-base reactions.
Some examples of Usanovich acid-base reactions include:
Na2o (base) + SO3 (acid) → 2 Na+ + SO2−4 (species exchanged: O2− anion)
3 (NH4)2S(base)+ Sb2S5 (acid) → 6 NH+ 4 + 2 B 3− 4(species exchanged: 3 S2− anions)
2 (base) + Cl 2 (a) → 2Na+ + 2Cl− (species exchanged: 2 electrons)
Rationalizing the Strength of Lewis Acid-Base Interactions
In 1963, Ralph Pearson proposed a qualitative concept known as hard and soft acid and base theory.  Later made quantitative in 1984 with the help of Robert Parr.   ‘Hard’ is applied to species that are smaller, have higher charge, and have weak polarization. ‘Soft’ applies to species that are larger, have fewer charge states and are strongly polarizable. Acids and bases interact, and the most stable interactions are hard-hard and soft-soft. This principle has found use in organic and inorganic chemistry.
The ECW model created by Russell S. Drago describes a quantitative model and predicts the strength of the Lewis acid-base interaction , -Δ H. The model gave the E and C parameters to several Lewis acids and bases. Each acid is characterized by E A and C A. Similarly each base is characterized by its own E B and C B. The E and C parameters, respectively, make up the electrostatic and covalent contributions to the bond strengths that will form acids and bases. the equation is−Δ H = E A E B + C A C B + W
The W term represents a constant energy contribution to an acid-base reaction such as the cleavage of a dimeric acid or base. The equation predicts the inversion of acid and base strengths. The graphical representation of the equation shows that there is no single sequence of Lewis base strengths or Lewis acid strengths. 
The reaction of a strong acid and a strong base is essentially a quantitative reaction. for example,
HCl (aq) + Na(OH) (aq) → H 2 O + NaCl (aq)
In this reaction both sodium and chloride ions are spectators as in the neutral reaction,
H + + OH – → H 2 O
does not include them. The addition of acids with weak bases is not quantitative because a solution of a weak base is a buffer solution. A solution of a weak acid is also a buffer solution. An equilibrium mixture is formed when a weak acid reacts with a weak base. For example, adenine, which is written as AH, is the hydrogen phosphate ion, HPO2-4. can react with
AH + HPO2−4 ⇌ A− + H2PO−4
The equilibrium constant for this reaction can be derived from the acid dissociation constant of adenine and the dihydrogen phosphate ion.
[A−] [H+] = Ka1[AH]
[HPO2−4] [H+] = Ka2[H2PO−4]
The notation [X] denotes “the concentration of X”. When these two equations are combined by eliminating the hydrogen ion concentration, the expression for the equilibrium constant, K , is obtained.
[A−] [H2PO−4] = K[AH] [HPO2−4]; K = Ka1/Ka2
The acid–base reaction is a special case of the acid–base reaction, where the base used is also an alkali. When an acid reacts with a base salt (metal hydroxide), the products are the metal salt and water. Acid-base reactions are also neutralization reactions.
In general, acid-base reactions can be simplified to
OH−(aq) + H+(aq) → H2O
Except for the spectator ions.
Acids are normally pure substances that contain hydrogen cations ( H+) or cause them to be produced in solution. Hydrochloric acid ( HCl ) and sulfuric acid ( H.)2So4) are common examples. In water, these break down into ions:
HCl → H+(aq) + cl–(aq)
h2So4→ h+(aq) + HSO–4(aq)
The alkali dissociates in water, giving the dissolved hydroxide ion:
NaOH → Na+(aq) + OH–(aq)