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Acid–base reaction
Acid–base reaction
An acid–base reaction is a chemical reaction that occurs between an acid and a base. Several concepts that provide
alternative definitions for the reaction mechanisms involved and their application in solving related problems exist.
Despite several differences in definitions, their importance becomes apparent as different methods of analysis when
applied to acid–base reactions for gaseous or liquid species, or when acid or base character may be somewhat less
apparent. The first of these scientific concepts of acids and bases was provided by the French chemist Antoine
Lavoisier, circa 1776.[1]
Historic acid–base theories
Lavoisier's oxygen theory of acids
The first scientific concept of acids and bases was provided by Antoine Lavoisier circa 1776. Since Lavoisier's
knowledge of strong acids was mainly restricted to oxoacids, such as HNO3 (nitric acid) and H2SO4 (sulfuric acid),
which tend to contain central atoms in high oxidation states surrounded by oxygen, and since he was not aware of the
true composition of the hydrohalic acids (HF, HCl, HBr, and HI), he defined acids in terms of their containing
oxygen, which in fact he named from Greek words meaning "acid-former" (from the Greek οξυς (oxys) meaning
"acid" or "sharp" and γεινομαι (geinomai) meaning "engender"). The Lavoisier definition was held as absolute truth
for over 30 years, until the 1810 article and subsequent lectures by Sir Humphry Davy in which he proved the lack of
oxygen in H2S, H2Te, and the hydrohalic acids. However, Davy failed to develop a new theory, concluding that
"acidity does not depend upon any particular elementary substance, but upon peculiar arrangement of various
substances".[2] One notable modification of oxygen theory was provided by Berzelius, who stated that acids are
oxides of nonmetals while bases are oxides of metals.
Liebig's hydrogen theory of acids
This definition was proposed by Justus von Liebig circa 1838,[3] based on his extensive works on the chemical
composition of organic acids. This finished the doctrinal shift from oxygen-based acids to hydrogen-based acids,
started by Davy. According to Liebig, an acid is a hydrogen-containing substance in which the hydrogen could be
replaced by a metal.[4] Liebig's definition, while completely empirical, remained in use for almost 50 years until the
adoption of the Arrhenius definition.[5]
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Acid–base reaction
2
Common acid–base theories
Arrhenius definition
The Arrhenius definition of acid–base reactions is a development of
the hydrogen theory of acids, devised by Svante Arrhenius, which was
used to provide a modern definition of acids and bases that followed
from his work with Friedrich Wilhelm Ostwald in establishing the
presence of ions in aqueous solution in 1884, and led to Arrhenius
receiving the Nobel Prize in Chemistry in 1903 for "recognition of the
extraordinary services... rendered to the advancement of chemistry by
his electrolytic theory of dissociation".[6]
As defined by Arrhenius, acid–base reactions are characterized by
Arrhenius acids, which dissociate in aqueous solution to form
hydrogen ions (H+),[6] and Arrhenius bases, which form hydroxide ()
ions. More recent IUPAC recommendations now suggest the newer
term "hydronium"[7] be used in favor of the older accepted term
"oxonium"[8] to illustrate reaction mechanisms such as those defined in
Svante Arrhenius
the Brønsted–Lowry and solvent system definitions more clearly, with
the Arrhenius definition serving as a simple general outline of
acid–base character.[6] The Arrhenius definition can be summarised as "Arrhenius acids form hydrogen ions in
aqueous solution with Arrhenius bases forming hydroxide ions."
The universal aqueous acid–base definition of the Arrhenius concept is described as the formation of water from
hydrogen and hydroxide ions, or hydrogen ions and hydroxide ions from the dissociation of an acid and base in
aqueous solution:
H+ (aq) + (aq)
H2O
(In modern times, the use of H+ is regarded as a shorthand for H3O+, since it is now known that the bare proton H+
does not exist as a free species in solution.)
This leads to the definition that in Arrhenius acid–base reactions, a salt and water is formed from the reaction
between an acid and a base.[6] In other words, this is a neutralization reaction.
acid+ + base− → salt + water
The positive ion from a base forms a salt with the negative ion from an acid. For example, two moles of the base
sodium hydroxide (NaOH) can combine with one mole of sulfuric acid (H2SO4) to form two moles of water and one
mole of sodium sulfate.
2 NaOH + H2SO4 → 2 H2O + Na2SO4
The Arrhenius definitions of acidity and alkalinity are restricted to aqueous solutions, and refer to the concentration
of the solvent ions. Under this definition, pure H2SO4 or HCl dissolved in toluene are not acidic, and molten KOH
and solutions of sodium amide in liquid ammonia are not alkaline.
Acid–base reaction
3
Solvent system definition
One of the limitations of Arrhenius definition was its reliance on water solutions. Edward C. Franklin studied the
acid–base reactions in liquid ammonia in 1905 and pointed out the similarities to water-based Arrhenius theory, and
Albert F. O. Germann, working with liquid COCl2, generalized Arrhenius definition to cover aprotic solvents and
formulated the solvent system theory in 1925.[9]
Germann pointed out that in many solvents there is a certain concentration of a positive species, solvonium (earlier
lyonium) cations and negative species, solvate (earlier lyate) anions, in equilibrium with the neutral solvent
molecules. For example, water and ammonia undergo such dissociation into hydronium and hydroxide, and
ammonium and amide, respectively:
2 H2O
H3O+ +
2 NH3
NH
+ NH2−
Some aprotic systems also undergo such dissociation, such as dinitrogen tetroxide into nitrosonium and nitrate,
antimony trichloride into dichloroantimonium and tetrachloroantimonate, and phosgene into chlorocarboxonium and
chloride.
N2O4
2 SbCl3
COCl2
NO+ + NO3−
SbCl
+
+ SbCl4−
COCl +
A solute causing an increase in the concentration of the solvonium ions and a decrease in the solvate ions is defined
as an acid and one causing the reverse is defined as a base. Thus, in liquid ammonia, KNH2 (supplying NH2−) is a
strong base, and NH4NO3 (supplying NH ) is a strong acid. In liquid sulfur dioxide (SO2), thionyl compounds
(supplying SO2+) behave as acids, and sulfites (supplying SO32−) behave as bases.
The non-aqueous acid–base reactions in liquid ammonia are similar to the reactions 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 → NO
+ H3O+ + 2 HSO4−
The unique strength of this definition shows in describing the reactions in aprotic solvents, for example in liquid
N2O4:
AgNO3 (base) + NOCl (acid) → N2O4 (solvent) + AgCl (salt)
Since solvent-system definition depends on the solvent as well as on the compound itself, the same compound can
change its role depending on the choice of the solvent. Thus, HClO4 is a strong acid in water, a weak acid in acetic
acid, and a weak base in fluorosulfonic acid. This was seen as both a strength and a weakness, since some
substances, such as SO3 and NH3 were felt to be acidic or basic on their own right. On the other hand, solvent system
theory was criticized as too general to be useful; it was felt that there was something intrinsically acidic about
hydrogen compounds, not shared by non-hydrogenic solvonium salts.[2]
Brønsted–Lowry definition
The Brønsted–Lowry definition, formulated in 1923, independently by Johannes Nicolaus Brønsted in Denmark and
Martin Lowry in England, is based upon the idea of protonation of bases through the de-protonation of acids — that
is, the ability of acids to "donate" hydrogen ions (H+) or protons to bases, which "accept" them.[10] Unlike the
previous definitions, the Brønsted–Lowry definition does not refer to the formation of salt and solvent, but instead to
the formation of conjugate acids and conjugate bases, produced by the transfer of a proton from the acid to the
base.[6] [10] In this approach, acids and bases are fundamentally different in behavior from salts, which are seen as
Acid–base reaction
4
electrolytes, subject to the theories of Debye, Onsager, 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. The concept of neutralization is thus absent.[2]
According to Brønsted–Lowry definition, an acid is a compound that can donate a proton, and a base is a compound
that can receive a proton. An acid–base reaction is, thus, the removal of a hydrogen ion from the acid and its addition
to the base.[11] This does not refer to the removal of a proton from the nucleus of an atom, which would require
levels of energy not attainable through the simple dissociation of acids, but to removal of a hydrogen ion (H+).
The removal of a proton (hydrogen ion) from an acid produces its conjugate base, which is the acid with a hydrogen
ion removed, and the reception of a proton by a base produces its conjugate acid, which is the base with a hydrogen
ion added.
For example, the removal of H+ from hydrochloric acid (HCl) produces the chloride ion (), the conjugate base of the
acid:
HCl → H+ +
The addition of H+ to the hydroxide ion (), a base, produces water (H2O), its conjugate acid:
H+ + → H2O
Although Brønsted–Lowry acid–base behavior is formally independent of any solvent, it encompasses Arrhenius
and solvent system definitions in an unenforced way. For example, protonation of ammonia, a base, gives
ammonium ion, its conjugate acid:
H+ + NH3 → NH
The reaction of ammonia, a base, with acetic acid in absence of water can be described to give ammonium cation, an
acid, and acetate anion, a base:
CH3COOH + NH3 → NH
+ CH3COO
This definition also explains the dissociation of water into low concentrations of hydronium and hydroxide ions:
H2O + H2O
H3O+ +
Water, being amphoteric, can act as both an acid and a base; here, one molecule of water acts as an acid, donating a
H+ ion and forming the conjugate base, , and a second molecule of water acts as a base, accepting the H+ ion and
forming the conjugate acid, H3O+.
Acid dissociation and acid hydrolysis are seen to be entirely similar phenomena:
HCl (acid) + H2O (base)
NH
(acid) + H2O (base)
H3O+ (acid) + (base)
H3O+ (acid) + NH3 (base)
as are basic dissociation and basic hydrolysis:
NH3 (base) + H2O (acid)
NH
CH3COO (base) + H2O (acid)
(acid) + (base)
CH3COOH (acid) + (base)
Acid–base reaction
Thus, the general formula for acid–base reactions according to the Brønsted–Lowry definition is:
AH + B → BH+ +
where AH represents the acid, B represents the base, BH+ represents the conjugate acid of B, and represents the
conjugate base of AH.
Although Brønsted–Lowry calls hydrogen-containing substances like HCl acids, KOH and KNH2 are not bases but
salts containing the bases and NH2−. Also, some substances, which many chemists considered to be acids, such as
SO3 or BCl3, are excluded from this classification due to lack of hydrogen. Gilbert Lewis wrote in 1938, "To restrict
the group of acids to those substances which contain hydrogen interferes as seriously with the systematic
understanding of chemistry as would the restriction of the term oxidizing agent to substances containing oxygen."[2]
Lewis definition
The hydrogen requirement of Arrhenius and Brønsted–Lowry was removed by the Lewis definition of acid–base
reactions, devised by Gilbert N. Lewis in 1923,[12] in the same year as Brønsted–Lowry, but it was not elaborated by
him until 1938.[2] Instead of defining acid–base reactions in terms of protons or other bonded substances, the Lewis
definition defines a base (referred to as a Lewis base) to be a compound that can donate an electron pair, and an acid
(a Lewis acid) to be a compound that can receive this electron pair.[13]
In this system, an acid does not exchange atoms with a base, but combines with it. For example, consider this
classical aqueous acid–base reaction:
HCl (aq) + NaOH (aq) → H2O (l) + NaCl (aq)
The Lewis definition does not regard this reaction as the formation of salt and water or the transfer of H+ from HCl
to . Instead, it regards the acid to be the H+ ion itself, and the base to be the ion, which has an unshared electron pair.
Therefore, the acid–base reaction here, according to the Lewis definition, is the donation of the electron pair from to
the H+ ion. This forms a covalent bond between H+ and , thus producing water (H2O).
By treating acid–base reactions in terms of electron pairs instead of specific substances, the Lewis definition can be
applied to reactions that do not fall under other definitions of acid–base reactions. For example, a silver cation
behaves as an acid with respect to ammonia, which behaves as a base, in the following reaction:
Ag+ + 2 :NH3 → [H3N:Ag:NH3]+
The result of this reaction is the formation of an ammonia–silver adduct.
In reactions between Lewis acids and bases, there is the formation of an adduct[13] when the highest occupied
molecular orbital (HOMO) of a molecule, such as NH3 with available lone electron pair(s) donates lone pairs of
electrons to the electron-deficient molecule's lowest unoccupied molecular orbital (LUMO) through a co-ordinate
covalent bond; in such a reaction, the HOMO-interacting molecule acts as a base, and the LUMO-interacting
molecule acts as an acid.[13] In highly-polar molecules, such as boron trifluoride (BF3),[13] the most electronegative
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Acid–base reaction
6
element pulls electrons towards its own orbitals, providing a more positive charge on the less-electronegative
element and a difference in its electronic structure due to the axial or equatorial orbiting positions of its electrons,
causing repulsive effects from lone pair – bonding pair (Lp–Bp) interactions between bonded atoms in excess of
those already provided by bonding pair – bonding pair (Bp–Bp) interactions.[13] Adducts involving metal ions are
referred to as co-ordination compounds.[13]
Other acid–base theories
Usanovich definition
Simultaneously with Lewis, a Soviet chemist Mikhail Usanovich from Tashkent, developed a general theory that
does not restrict acidity to hydrogen-containing compounds, but his approach, published in 1938, was even more
general than Lewis theory.[2] Usanovich's theory can be summarized as defining an acid as anything that accepts
negative species or donates positive ones, and a base as the reverse. This pushed the concept of acid–base reactions
to its logical limits, and even redefined the concept of redox (oxidation-reduction) as a special case of acid-base
reactions, and so did not become widespread, despite being easier to understand than Lewis theory, which required
detailed familiarity with atomic structure. Some examples of Usanovich acid-base reactions include:
Na2O (base) + SO3 (acid) → 2 Na+ + SO42− (species exchanged: anion )
3 (NH4)2S (base) + Sb2S3 (acid) → 6 NH
+
+ 2 SbS42− (species exchanged: anion )
Na (base) + Cl (acid) → Na + (species exchanged: electron)
Lux–Flood definition
This acid–base theory was a revival of oxygen theory of acids and bases, proposed by German chemist Hermann
Lux[14] [15] in 1939, further improved by Håkon Flood circa 1947[16] and is still used in modern geochemistry and
electrochemistry of molten salts. This definition describes an acid as an oxide ion () acceptor and a base as an oxide
ion donor. For example:[17]
MgO (base) + CO2 (acid) → MgCO3
CaO (base) + SiO2 (acid) → CaSiO3
NO3− (base) + S2O72− (acid) → NO
+ 2 SO42−
Pearson definition
In 1963[18] Ralph Pearson proposed an advanced qualitative concept known as Hard Soft Acid Base principle, later
made quantitative with help of Robert Parr in 1984. 'Hard' applies to species that are small, have high charge states,
and are weakly polarizable. 'Soft' applies to species that are large, have low charge states and are strongly
polarizable. Acids and bases interact, and the most stable interactions are hard–hard and soft–soft. This theory has
found use in organic and inorganic chemistry.
Acid–base reaction
Acid–alkali reaction
An acid–alkali reaction is a special case of an acid–base reaction, where the base used is also an alkali. When an acid
reacts with an alkali it forms a metal salt and water. Acid–alkali reactions are also a type of neutralization reaction.
In general acid–alkali reactions can be simplified to
(aq) + H+(aq) → H2O
by omitting spectator ions.
Acids are generally pure substances which contain hydrogen ions (H+) or cause them to be produced in solutions.
Hydrochloric acid (HCl) and sulfuric acid (H2SO4) are common examples. In water, these break apart into ions:
HCl → H+(aq) + (aq)
H2SO4 → H+(aq) + HSO4−(aq)
An alkali is a base, more precisely a base which contains a metal from column 1 or 2 of the periodic table (the alkali
metals or the alkaline earth metals). Alkalis may be defined as soluble bases, which means they must be able to
dissolve in water. Bases generally are defined as substances which contain hydroxide ion () or produce it in solution.
Therefore, one may also speak of hydroxide bases which dissolve in water, and thus these would also be alkalis.
Some examples, then, of alkalis would be sodium hydroxide (NaOH), potassium hydroxide (KOH), magnesium
hydroxide (Mg(OH)2), and calcium hydroxide (Ca(OH)2). Note that only hydroxides with an alkali metal—column
1—are very soluble in water; hydroxides with an alkaline earth metal—column 2—are not as soluble. Some sources
will even say the alkaline earth metal hydroxides are insoluble.
To produce hydroxide ions in water, the alkali breaks apart into ions as below:
NaOH → Na+(aq) + (aq)
However, alkalies may also have a broader definition which includes carbonates (CO32−) bonded to a column 1
metal, an ammonium ion (NH ), or an amine (NHx radical) as the positive ion. Examples of alkalis would then
also include Li2CO3, Na2CO3, and (NH4)2CO3.
There seems to be conflicting information on acid-base reactions being neutralization reactions. Some sources
define a neutralization reaction as the reaction between an acid and a base which produces a salt and water. Yet in
the book Chemical Misconceptions: Prevention, Diagnosis and Cure by K. Tabor (2002), it is noted that “the term
neutralization is usually reserved for acid–alkali reactions.” Thus this does not make acid–alkali a type of
neutralization reaction, but the only kind of neutralization reaction.
There are many uses of neutralization reactions which are acid-alkali reactions. A very common use is antacid
tablets. These are designed to neutralize excess stomach acid (HCl) which may be causing discomfort in the stomach
or lower esophagus. Also in the digestive tract, neutralization reactions are used when food is moved from the
stomach to the intestines. In order for the nutrients to be absorbed through the intestinal wall, an alkaline
environment is needed, so the pancreas produce an antacid bicarbonate to cause this transformation to occur.[19]
Another common use, though perhaps not as widely known, is in fertilizers and control of soil pH. Slaked lime
(calcium hydroxide) or limestone (calcium carbonate) may be worked into soil that is too acidic for plant growth.[20]
Fertilizers which improve plant growth are made by neutralizing sulfuric acid (H2SO4) or nitric acid (HNO3) with
ammonia gas (NH3) making ammonium sulfate or ammonium nitrate. These are salts utilized in the fertilizer.[21]
Industrially, a by-product of the burning of coal, sulfur dioxide gas may combine with water vapor in the air to
eventually produce sulfuric acid, which falls as acid rain. To prevent the sulfur dioxide from being released, a device
known as a scrubber gleans the gas from smoke stacks. This device first blows calcium carbonate into the
combustion chamber where it decomposes into calcium oxide (lime) and carbon dioxide. This lime then reacts with
the sulfur dioxide produced forming calcium sulfite. A suspension of lime is then injected into the mixture to
produce a slurry, which removes the calcium sulfite and any remaining unreacted sulfur dioxide.[22]
7
Acid–base reaction
References
[1] Miessler, L. M., Tar, D. A., (1991) p. 166 – Table of discoveries attributes Antoine Lavoisier as the first to posit a scientific theory in relation
to oxyacids.
[2] Hall, Norris F. (March 1940). "Systems of Acids and Bases". J. Chem. Educ. 17 (3): 124–128. doi:10.1021/ed017p124.
[3] Miessler, L. M., Tar, D. A., (1991) "Inorganic Chemistry" 2nd ed. Pearson Prentice-Hall p. 166 – table of discoveries attributes Justus von
Liebig's publication as 1838
[4] Meyers, R. (2003). The Basics of Chemistry. Greenwood Press. p. 156.
[5] H. L. Finston and A. C. Rychtman, A New View of Current Acid-Base Theories, John Wiley & Sons, New York, 1983, pp. 140–146.
[6] Miessler, L. M., Tar, D. A., (1991) "Inorganic Chemistry" 2nd ed. Pearson Prentice-Hall p. 165
[7] Murray, K. K., Boyd, R. K., et al. (2006) "Standard definition of terms relating to mass spectrometry recommendations" International Union
of Pure and Applied Chemistry. – Please note that, in this document, there is no reference to deprecation of "oxonium", which is also still
accepted as it remains in the IUPAC Gold book, but rather reveals preference for the term "Hydronium".
[8] International Union of Pure and Applied Chemistry (2006) IUPAC Compendium of Chemical Terminology, Electronic version Retrieved from
International Union of Pure and Applied Chemistry on 9 May 2007 on URL http:/ / goldbook. iupac. org/ O04379. html "Oxonium Ions"
[9] Germann, Albert F. O. (6 October 1925). "A General Theory of Solvent Systems". J.Am.Chem.Soc. 47 (10): 2461–2468.
doi:10.1021/ja01687a006.
[10] Miessler, L. M., Tar, D. A., (1991) "Inorganic Chemistry" 2nd ed. Pearson Prentice-Hall pp. 167–169 – According to this page, the original
definition was that "acids have a tendency to lose a proton"
[11] Clayden, J., Warren, S., et al. (2000) "Organic Chemistry" Oxford University Press pp. 182–184
[12] Miessler, L. M., Tar, D. A., (1991) "Inorganic Chemistry" 2nd ed. Pearson Prentice-Hall p. 166 – Table of discoveries attributes the date of
publication/release for the Lewis theory as 1923.
[13] Miessler, L. M., Tar, D. A., (1991) "Inorganic Chemistry" 2nd ed. Pearson Prentice-Hall pp. 170–172
[14] Franz, H. (1966). "Solubility of Water Vapor in Alkali Borate Melts". J. Am. Ceram. Soc. 49 (9): 473–477.
doi:10.1111/j.1151-2916.1966.tb13302.x.
[15] Lux, Hermann (1939). ""Säuren" und "Basen" im Schmelzfluss: die Bestimmung. der Sauerstoffionen-Konzentration". Ztschr. Elektrochem
45 (4): 303–309.
[16] Flood, H.; Forland, T. (1947). "The Acidic and Basic Properties of Oxides". Acta Chem. Scand. 1 (6): 592.
doi:10.3891/acta.chem.scand.01-0592. PMID 18907702.
[17] Drago, Russel S.; Whitten, Kenneth W. (1966). "The Synthesis of Oxyhalides Utilizing Fused-Salt Media". Inorg. Chem. 5 (4): 677–682.
doi:10.1021/ic50038a038.
[18] Pearson, Ralph G. (1963). "Hard and Soft Acids and Bases". J. Am. Chem. Soc. 85 (22): 3533–3539. doi:10.1021/ja00905a001.
[19] Print this page (http:/ / www. wddty. com/ UtilityPages/ Print. aspx?nodeId=-3363800369331166395). WDDTY. Retrieved on 2010-12-10.
[20] Neutralisation – ‘curing acidity’ - Acids, alkalis and salts - Intermediate - Experiments (http:/ / www. practicalchemistry. org/ experiments/
intermediate/ acids-alkalis-and-salts/ neutralisation-curing-acidity,103,EX. html). Practical Chemistry. Retrieved on 2010-12-10.
[21] Reversible Reactions, Chemical Equilibrium, Ammonia & Nitric acid, their Manufacture and Uses e.g. in Artificial Fertilisers (http:/ / www.
docbrown. info/ page04/ 4_74revNH3. htm)
[22] Zumdahl, 2000, pp. 226, 228
External links
• Acid-base Physiology: an on-line text (http://www.anaesthesiamcq.com/AcidBaseBook/ABindex.php)
• John W. Kimball's online Biology book section of acid and bases. (http://users.rcn.com/jkimball.ma.ultranet/
BiologyPages/A/Acids_Bases.html)
8
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