Acid–base reaction

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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 (hydrogen fluroide), and HI) (hydrogen iodide), 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]

Common acid-base theories

Arrhenius definition

File:Arrhenius2.jpg
Svante Arrhenius

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 (OH−) 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 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) + OH− (aq) 15px 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.

Solvent system definition

One of the limitations of Arrhenius definition was its reliance on water solutions. E. 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 15px H3O+ + OH−
2 NH3 15px 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.

N2O4 15px NO+ + NO−3
2 SbCl3 15px SbCl+2 + SbCl−4
COCl2 15px COCl+ + Cl−

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 NH−2) is a strong base, and NH4NO3 (supplying NH+4) is a strong acid. In liquid sulfur dioxide (SO2), 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 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+2 + H3O+ + 2 HSO−4

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

Main article: Brønsted–Lowry acid-base theory

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 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 (Cl−), the conjugate base of the acid:

HCl → H+ + Cl−

The addition of H+ to the hydroxide ion (OH−), a base, produces water (H2O), its conjugate acid:

H+ + OH− → 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+4

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+4 + CH3COO−

This definition also explains the dissociation of water into low concentrations of hydronium and hydroxide ions:

H2O + H2O 15px H3O+ + OH−

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, OH−, 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) 15px H3O+ (acid) + Cl− (base)
NH+4 (acid) + H2O (base) 15px H3O+ (acid) + NH3 (base)

as are basic dissociation and basic hydrolysis:

NH3 (base) + H2O (acid) 15px NH+4 (acid) + OH− (base)
CH3COO− (base) + H2O (acid) 15px CH3COOH (acid) + OH− (base)
600px

Thus, the general formula for acid-base reactions according to the Brønsted-Lowry definition is:

AH + 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 AH.

Although Brønsted-Lowry calls hydrogen-containing substances like HCl acids, KOH and KNH2 are not bases but salts containing the bases OH− and NH−2. 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

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,[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 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 (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 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 wide spread, 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+ + SO2−4 (species exchanged: anion O2−)
3 (NH4)2S (base) + Sb2S3 (acid) → 6 NH4+ + 2 SbS2−4 (species exchanged: anion S2−)
Na (base) + Cl (acid) → Na+ + Cl− (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 (O2−) acceptor and a base as an oxide ion donor. For example:[17]

MgO (base) + CO2 (acid) → MgCO3
CaO (base) + SiO2 (acid) → CaSiO3
NO−3 (base) + S2O2−7 (acid) → NO+2 + 2 SO2−4

Pearson definition

Main article: HSAB theory

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

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) → 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) + Cl−(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). The Science Module 3 within The Digital Brain [1] (Science KS3 SU3 Module 3, Acids and Bases) and Doc Brown (http://www.docbrown.info/, 2000-2008, p. 3) define alkalis as soluble bases, which means they must be able to dissolve in water. Bases generally are defined as substances which contain hydroxide ion (OH-) or produce it in solution. Therefore, we 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) + OH-(aq)

However, alkalies may also have a broader definition which includes carbonates (CO32-) bonded to a column 1 metal, an ammonium ion (NH4+), 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. (http://www.wddty.com/UtilityPages/Print.aspx?nodeId=-3363800369331166395)

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. (http://www.practicalchemistry.org/experiments/intermediate/acids-alkalis-and-salts/neutralisation-curing-acidity,103,EX.html) 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. (http://www.docbrown.info, 2000-2008, p. 3)

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. (Zumdahl, 2000, p. 226, 228)

See also

References

  1. ↑ Miessler, L. M., Tar, D. A., (1991) p166 - Table of discoveries attributes Antoine Lavoisier as the first to posit a scientific theory in relation to oxyacids.
  2. ↑ 2.0 2.1 2.2 2.3 2.4 2.5 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 p166 - 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, 1982, pp. 140-146.
  6. ↑ 6.0 6.1 6.2 6.3 6.4 Miessler, L. M., Tar, D. A., (1991) "Inorganic Chemistry" 2nd ed. Pearson Prentice-Hall p165
  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 Oct 1925). "A General Theory of Solvent Systems". J.Am.Chem.Soc. 47 (10): 2461–2468. doi:10.1021/ja01687a006. 
  10. ↑ 10.0 10.1 Miessler, L. M., Tar, D. A., (1991) "Inorganic Chemistry" 2nd ed. Pearson Prentice-Hall p167-169 -- 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 p182-184
  12. ↑ Miessler, L. M., Tar, D. A., (1991) "Inorganic Chemistry" 2nd ed. Pearson Prentice-Hall p166 - Table of discoveries attributes the date of publication/release for the Lewis theory as 1923.
  13. ↑ 13.0 13.1 13.2 13.3 13.4 13.5 Miessler, L. M., Tar, D. A., (1991) "Inorganic Chemistry" 2nd ed. Pearson Prentice-Hall p170-172
  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. ↑ Lua error in package.lua at line 80: module 'Module:Citation/CS1/Suggestions' not found.
  17. ↑ Lua error in package.lua at line 80: module 'Module:Citation/CS1/Suggestions' not found.
  18. ↑ Pearson, Ralph G. (1963). "Hard and Soft Acids and Bases". J. Am. Chem. Soc. 85 (22): 3533–3539. doi:10.1021/ja00905a001. 

External links

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