Acid–base reaction

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An acid–base reaction is a chemical reaction that occurs between an acid and a base. Several theoretical frameworks provide alternative conceptions of the reaction mechanisms and their application in solving related problems. 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, circa 1776.[1]

Common acid–base theories[edit]

Arrhenius definition[edit]

Svante Arrhenius

The first modern definition of acids and bases was devised by Svante Arrhenius. A hydrogen theory of acids, it followed from his 1884 work with Friedrich Wilhelm Ostwald in establishing the presence of ions in aqueous solution and led to Arrhenius receiving 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+);[2] that is, an acid increases the concentration of H+ ions in an aqueous solution.

This causes the protonation of water, or the creation of the hydronium (H3O+) ion.[5] Thus, in modern times, the use of H+ is regarded as a shorthand for H3O+, because it is now known that a bare proton does not exist as a free species in aqueous solution.[6]

  • an Arrhenius base is a substance that dissociates in water to form hydroxide (OH) ions; that is, a base increases the concentration of OH ions in an aqueous solution.

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 and HCl dissolved in toluene are not acidic, and molten NaOH and solutions of calcium amide in liquid ammonia are not alkaline.

Overall, to qualify as an Arrhenius acid, upon the introduction to water, the chemical must either cause, directly or otherwise:

  • an increase in the aqueous hydronium concentration, or
  • a decrease in the aqueous hydroxide concentration.

Conversely, to qualify as an Arrhenius base, upon the introduction to water, the chemical must either cause, directly or otherwise:

  • a decrease in the aqueous hydronium concentration, or
  • an increase in the aqueous hydroxide concentration.

The universal aqueous acid–base definition of the Arrhenius concept is described as the formation of a water molecule from a proton and hydroxide ion. This leads to the definition that in Arrhenius acid–base reactions, a salt and water are formed from the reaction between an acid and a base.[2] This is a neutralization reaction - the acid and base properties of H+ and OH are neutralized, for they combine to form H2O, the water molecule. The acid-base neutralization reaction can be put into a word equation:

acid + base → salt + water

The positive ion from a base and the negative ion from an acid form a salt together - in other words, an acid-base neutralization reaction is a double-replacement reaction. For example, when a neutralization reaction takes place between hydrochloric acid (HCl) and sodium hydroxide (NaOH), the products are sodium chloride (common table salt) and water.

HCl(aq) + NaOH(aq) → NaCl + H2O


Notice how the cations and the anions merely switched places: the Na+ from the NaOH combined with the Cl- from the HCl to form NaCl, while the OH from the NaOH combined with the H+ from the HCl to form H2O.

Solvent system definition[edit]

One of the limitations of the Arrhenius definition is its reliance on water solutions. Edward Curtis Franklin studied the acid–base reactions in liquid ammonia in 1905 and pointed out the similarities to the water-based Arrhenius theory. Albert F. O. Germann, working with liquid COCl
2
, formulated the solvent-based theory in 1925, thereby generalizing the Arrhenius definition to cover aprotic solvents.[7]

Germann pointed out that in many solutions, there are ions in equilibrium with the neutral solvent molecules:

  • solvonium:[8] A generic name for a positive ion.
  • solvate:[9] A generic name for a negative ion.

For example, water and ammonia undergo such dissociation into hydronium and hydroxide, and ammonium and amide, respectively:

2 H
2
O
is in equilibrium with H
3
O+
+ OH
2 NH
3
is in equilibrium with NH+
4
+ NH
2

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:

N
2
O
4
is in equilibrium with NO+
+ NO
3
2 SbCl
3
is in equilibrium with SbCl+
2
+ SbCl
4
COCl
2
is in equilibrium with COCl+
+ Cl

A solute that causes an increase in the concentration of the 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 the solvate ions and a decrease in the concentration of the solvonium ions is defined as a base.

Thus, in liquid ammonia, KNH
2
(supplying NH
2
) is a strong base, and NH
4
NO
3
(supplying NH+
4
) is a strong acid. In liquid sulfur dioxide (SO
2
), thionyl compounds (supplying SO2+
) behave as acids, and sulfites (supplying SO2−
3
) behave as bases.

The non-aqueous acid–base reactions in liquid ammonia are similar to the reactions in water:

2 NaNH
2
(base) + Zn(NH
2
)
2
(amphiphilic amide) → Na
2
[Zn(NH
2
)
4
]
2 NH
4
I
(acid) + Zn(NH
2
)
2
(amphiphilic amide) → [Zn(NH
3
)
4
)]I
2

Nitric acid can be a base in liquid sulfuric acid:

HNO
3
(base) + 2 H
2
SO
4
NO+
2
+ H
3
O+
+ 2 HSO
4

The unique strength of this definition shows in describing the reactions in aprotic solvents; for example, in liquid N
2
O
4
:

AgNO
3
(base) + NOCl (acid) → N
2
O
4
(solvent) + AgCl (salt)

Because the solvent system definition depends on the solute as well as on the solvent itself, a particular solute can be either an acid or a base depending on the choice of the solvent: HClO
4
is a strong acid in water, a weak acid in acetic acid, and a weak base in fluorosulfonic acid; this characteristic of the theory has been seen as both a strength and a weakness, because some substances (such as SO
3
and NH
3
) have been seen to be acidic or basic on their own right. On the other hand, solvent system theory has been criticized as being too general to be useful. Also, it has been thought that there is something intrinsically acidic about hydrogen compounds, a property not shared by non-hydrogenic solvonium salts.[10]

Brønsted–Lowry definition[edit]

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+)—otherwise known as protons—to bases, which "accept" them.[11] It is important to note that the theory 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+).

An acid–base reaction is, thus, the removal of a hydrogen ion from the acid and its addition to the base.[12] The removal of a hydrogen ion from an acid produces its conjugate base, which is the acid with a hydrogen ion removed. The reception of a proton by a base produces its conjugate acid, which is the base with a hydrogen ion added.

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.[2][11] In this approach, acids and bases are fundamentally different in behavior from salts, which are seen as 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.[10] Brønsted–Lowry acid–base behavior is formally independent of any solvent, making it more all-encompassing than the Arrhenius model.

The general formula for acid–base reactions according to the Brønsted–Lowry definition is:

HA + B → BH+ + A

where AH represents the acid, B represents the base, BH+ represents the conjugate acid of B, and A represents the conjugate base of HA.

For example, a Brønsted-Lowry model for the dissociation of hydrochloric acid (HCl) in aqueous solution would be the following:

HCl + H2O is in equilibrium with H3O+ + Cl

The removal of H+ from the HCl produces the chloride ion, Cl, the conjugate base of the acid. The addition of H+ to the H2O (acting as a base) forms the hydronium ion, H3O+, the conjugate acid of the base.

Water is amphoteric—that is, it can act as both an acid and a base. The Brønsted-Lowry model explains this, showing the dissociation of water into low concentrations of hydronium and hydroxide ions:

H2O + H2O is in equilibrium with H3O+ + OH

This equation is demonstrated in the image below:

Bronsted-lowry-3d-explanation-diagram.png

Here, one molecule of water acts as an acid, donating an H+ and forming the conjugate base, OH, and a second molecule of water acts as a base, accepting the H+ ion and forming the conjugate acid, H3O+.

As an example of water acting as an acid, consider an aqueous solution of ammonia (NH3)

NH3 + H2O is in equilibrium with NH4+ + OH

Ammonia acts as a base, accepting an H+ to form the ammonium ion (NH4+). Water acts an acid, donating a H+ to form the hydroxide ion (OH).

Furthermore, in the Brønsted-Lowry model, the solvent does not necessarily have to be water. For example, consider the reaction of ammonia, a base, with acetic acid (CH3COOH) in absence of water.

CH
3
COOH
+ NH
3
NH+
4
+ CH
3
COO

The removal of an H+ from acetic acid forms its conjugate base, the acetate ion (CH3COO). The addition of the H+ to the ammonia forms its conjugate acid, the ammonium ion (NH4+).

The Brønsted–Lowry model calls hydrogen-containing substances (like HCl) acids. Thus, 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 N. Lewis wrote in 1938, "To restrict the group of acids to those substances that contain hydrogen interferes as seriously with the systematic understanding of chemistry as would the restriction of the term oxidizing agent to substances containing oxygen."[10] Furthermore, KOH and KNH2 are not considered Brønsted bases, but rather salts containing the bases OH and NH2.

Lewis definition[edit]

Further information: Lewis acids and bases

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,[13] in the same year as Brønsted–Lowry, but it was not elaborated by him until 1938.[10] 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.[14]

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) → H
2
O
(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 OH
. Instead, it regards the acid to be the H+
ion itself, and the base to be the OH
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 OH
to the H+
ion. This forms a covalent bond between H+
and OH
, thus producing water (H
2
O
).

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 :NH
3
[H
3
N
:Ag:NH
3
]+

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[14] when the highest occupied molecular orbital (HOMO) of a molecule, such as NH
3
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.[14] In highly-polar molecules, such as boron trifluoride (BF
3
),[14] the most electronegative 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.[14] Adducts involving metal ions are referred to as co-ordination compounds.[14]

Other acid–base theories[edit]

Lux–Flood definition[edit]

This acid–base theory was a revival of oxygen theory of acids and bases, proposed by German chemist Hermann Lux[15][16] in 1939, further improved by Håkon Flood circa 1947[17] and is still used in modern geochemistry and electrochemistry of molten salts. This definition describes an acid as an oxide ion (O2−
) acceptor and a base as an oxide ion donor. For example:[18]

MgO (base) + CO
2
(acid) → MgCO
3
CaO (base) + SiO
2
(acid) → CaSiO
3
NO
3
(base) + S
2
O2−
7
(acid) → NO+
2
+ 2 SO2−
4

Pearson definition[edit]

Main article: HSAB theory

In 1963,[19] 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.

Usanovich definition[edit]

Mikhail Usanovich 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.[10] 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 defined the concept of redox (oxidation-reduction) as a special case of acid-base reactions

Some examples of Usanovich acid-base reactions include:

Na
2
O
(base) + SO
3
(acid) → 2 Na+
+ SO2−
4
(species exchanged: anion O2−
)
3 (NH
4
)
2
S
(base) + Sb
2
S
3
(acid) → 6 NH+
4
+ 2 SbS3−
4
(species exchanged: anion S2−
)
Na (base) + Cl (acid) → Na+
+ Cl
(species exchanged: electron)

Acid–alkali reaction[edit]

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

OH
(aq)
+ H+
(aq) → H
2
O

by omitting spectator ions.

Acids are in general pure substances that contain hydrogen ions (H+
) or cause them to be produced in solutions. Hydrochloric acid (HCl) and sulfuric acid (H
2
SO
4
) are common examples. In water, these break apart into ions:

HClH+
(aq) + Cl
(aq)
H
2
SO
4
H+
(aq) + HSO
4
(aq)

To produce hydroxide ions in water, the alkali breaks apart into ions as below:

NaOHNa+
(aq) + OH
(aq)

Historic acid–base theories[edit]

Lavoisier's oxygen theory of acids[edit]

The first scientific concept of acids and bases was provided by Lavoisier circa 1776. Since Lavoisier's knowledge of strong acids was mainly restricted to oxoacids, such as HNO
3
(nitric acid) and H
2
SO
4
(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".[10] 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[edit]

Circa 1838 Justus von Liebig proposed[20] that an acid is a hydrogen-containing substance in which the hydrogen could be replaced by a metal.[21] This redefinition was based on his extensive work on the chemical composition of organic acids, finishing the doctrinal shift from oxygen-based acids to hydrogen-based acids started by Davy. Liebig's definition, while completely empirical, remained in use for almost 50 years until the adoption of the Arrhenius definition.[22]

See also[edit]

References[edit]

  1. ^ Miessler, G.L., Tarr, D. A., "Inorganic Chemistry" (1991) p. 166 – Table of discoveries attributes Antoine Lavoisier as the first to posit a scientific theory in relation to oxyacids.
  2. ^ a b c d Miessler, G. L., Tarr, D. A., (1991) "Inorganic Chemistry" 2nd ed. Pearson Prentice-Hall p. 165
  3. ^ 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".
  4. ^ 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"
  5. ^ More recent IUPAC recommendations now suggest the newer term "hydronium"[3] be used in favor of the older accepted term "oxonium"[4] to illustrate reaction mechanisms such as those defined in the Brønsted–Lowry and solvent system definitions more clearly, with the Arrhenius definition serving as a simple general outline of acid–base character.[2]
  6. ^ LeMay, Eugene (2002). Chemistry. Upper Saddle River, New Jersey: Prentice-Hall. p. 602. ISBN 0-13-054383-7. 
  7. ^ 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. 
  8. ^ The term solvonium has replaced the older term lyonium.
  9. ^ the term solvate has replaced the older term lyate.
  10. ^ a b c d e f Hall, Norris F. (March 1940). "Systems of Acids and Bases". J. Chem. Educ. 17 (3): 124–128. Bibcode:1940JChEd..17..124H. doi:10.1021/ed017p124. 
  11. ^ a b Miessler, G. L., Tarr, 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"
  12. ^ Clayden, J., Warren, S., et al. (2000) "Organic Chemistry" Oxford University Press pp. 182–184
  13. ^ 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.
  14. ^ a b c d e f Miessler, G. L., Tarr, D. A., (1991) "Inorganic Chemistry" 2nd ed. Pearson Prentice-Hall pp. 170–172
  15. ^ 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. 
  16. ^ Lux, Hermann (1939). ""Säuren" und "Basen" im Schmelzfluss: die Bestimmung. der Sauerstoffionen-Konzentration". Ztschr. Elektrochem 45 (4): 303–309. 
  17. ^ Flood, H.; Forland, T. (1947). "The Acidic and Basic Properties of Oxides". Acta Chem. Scand. 1 (6): 592–604. doi:10.3891/acta.chem.scand.01-0592. PMID 18907702. 
  18. ^ 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. 
  19. ^ Pearson, Ralph G. (1963). "Hard and Soft Acids and Bases". J. Am. Chem. Soc. 85 (22): 3533–3539. doi:10.1021/ja00905a001. 
  20. ^ Miessler, G. L., Tarr, D. A., (1991)
  21. ^ Meyers, R. (2003). The Basics of Chemistry. Greenwood Press. p. 156.  "Inorganic Chemistry" 2nd ed. Pearson Prentice-Hall p. 166 – table of discoveries attributes Justus von Liebig's publication as 1838
  22. ^ H. L. Finston and A. C. Rychtman, A New View of Current Acid-Base Theories, John Wiley & Sons, New York, 1983, pp. 140–146.

External links[edit]