Acid-base concepts

Various concepts have developed around the terms acid and base in chemistry , which are based on different definitions of terms. The drive for this development is based on the one hand on the search for a definition that is as comprehensive and generally applicable as possible, and on the other hand on a specific area of ​​application of a concept.

The acid-base definition in the narrower sense usually used as a basis today is that according to Brønsted and Lowry , which, like the previous definition according to Svante Arrhenius , is based on proton transfer reactions in water, but also discusses those without the presence of water. Although this is a rather special definition, it is still widely used in chemistry.

In a broader sense, on the other hand, the Lewis definition is usually used today , in which case the term Lewis acids or Lewis bases is usually used explicitly .

Definition according to Arrhenius

The acid-base concept according to Svante Arrhenius was established in 1887 and is based on the ion theory , the experimentally determinable electrolytic conductivity of aqueous solutions containing salts , acids or bases. In addition to the salts, which are referred to as real electrolytes , acids are potential electrolytes , since as pure substances they do not conduct electricity. However, they are subject to electrolytic dissociation in water. The conductivity is based on the formation of freely moving positively charged particles, the cations , and negatively charged particles, the anions .

Acids

The characteristic feature of an acid is its dissociation into positively charged hydrogen ions (H ⁺ ions) and negatively charged anions in an aqueous solution. The anion of an acid is called an acid residue .

${\displaystyle \mathrm {S{\ddot {a}}uremolek{\ddot {u}}l\rightleftharpoons Wasserstoffion+S{\ddot {a}}urerestion} }$

As examples we mention the reactions of the acids hydrogen chloride , acetic acid and hydrogen cyanide when given in water:

${\displaystyle \mathrm {HCl\ \rightleftharpoons \ H_{(aq)}^{+}+Cl_{(aq)}^{-}} }$

${\displaystyle \mathrm {H_{3}C{-}COOH\ \rightleftharpoons \ H_{(aq)}^{+}\ +\ H_{3}C{-}COO_{(aq)}^{-}} }$

${\displaystyle \mathrm {HCN\ \rightleftharpoons \ H_{(aq)}^{+}+CN_{(aq)}^{-}} }$

These potential electrolytes are subject to a dissociation equilibrium (an equilibrium reaction ). The acids can be divided qualitatively into strong, medium-strength, and weak electrolytes. While strong acids such as hydrogen chloride largely dissociate when comparing one-molar solutions, weak acids such as hydrogen cyanide only lead to a low degree of dissociation and are only weak electrolytes.

Bases

Bases are compounds that dissociate into hydroxide ions (OH ^- ions) and cations in water . The cations of bases are called base residues . Metal hydroxides play an important role ; their cations are metal ions :

${\displaystyle \mathrm {Metallhydroxid\rightleftharpoons Basenrest+Hydroxidion} }$

As examples we mention the reactions of sodium hydroxide and calcium hydroxide when given in water:

${\displaystyle \mathrm {NaOH\ \rightleftharpoons \ Na_{(aq)}^{+}+OH_{(aq)}^{-}} }$

${\displaystyle \mathrm {Ca(OH)_{2}\ \rightleftharpoons \ Ca_{(aq)}^{2+}+2\ OH_{(aq)}^{-}} }$

Neutralization and the formation of salts

The reaction of a strong or moderately strong acid with an equivalent amount of a strong or moderately strong base is called neutralization . A solution of a salt is usually formed in the process.

${\displaystyle \mathrm {S{\ddot {a}}ure+Base\ {\xrightarrow[{}]{Neutralisation}}\ Salz_{(aq)}+Wasser} }$

${\displaystyle \mathrm {HCl+NaOH\ \rightleftharpoons \ NaCl+H_{2}O} }$

In the process, H ⁺ and OH ^- ions are converted into water, forming a neutral solution ( pH = 7):

${\displaystyle \mathrm {H^{+}+OH^{-}\longrightarrow \ H_{2}O} }$

Salts are compounds which dissociate into base and acid residues in water or in the melt. The reaction product of the above reaction corresponds to a solution formed by introducing sodium chloride in water:

${\displaystyle \mathrm {NaCl\ \rightleftharpoons \ Na_{(aq)}^{+}+Cl_{(aq)}^{-}} }$

Salts are also formed via other chemical reactions, see also salt formation reaction .

Hydrolysis by dissolving salts

The reverse direction of the reaction to neutralization is called hydrolysis (also: "salt hydrolysis") according to Arrhenius . Hydrolysis of a salt occurs when at least the acid residue or the base residue is derived from a weak acid or base. By dissolving such salts, basic (pH> 7) or acidic (pH <7) solutions are formed. Accordingly, these salts are called basic or acidic salts .

${\displaystyle \mathrm {Salz+Wasser\ {\xrightarrow[{}]{Hydrolyse}}\ S{\ddot {a}}ure+Base} }$

When sodium cyanide is dissolved, the very weak acid residue forms CN ^- molecular hydrogen cyanide . In the process, dissolved sodium hydroxide is formed, which leads to a basic solution:

${\displaystyle \mathrm {NaCN+H_{2}O\ \rightleftharpoons \ HCN+NaOH_{(aq)}} }$

Similarly, when ammonium chloride is dissolved , ammonia and dissociated hydrogen chloride are formed , which leads to an acidic solution:

${\displaystyle \mathrm {NH_{4}Cl+H_{2}O\ \rightleftharpoons \ NH_{3}+HCl_{(aq)}} }$

A hydrolysis _constant K _Hydr. Can be derived from the hydrolysis of a salt . In the case of a residue of a weak acid, the following applies:

${\displaystyle K_{\mathrm {Hydr.} }={\frac {c_{\mathrm {HA} }\cdot c_{\mathrm {OH^{-}} }}{c_{\mathrm {A^{-}} }}}\ {\Bigl (}={\frac {K\mathrm {_{W}} }{K\mathrm {_{S}} }}=K_{\mathrm {B} }{\Bigl )}}$

In the case of a remainder of a weak base:

${\displaystyle K_{\mathrm {Hydr.} }={\frac {c_{\mathrm {H^{+}} }\cdot c_{\mathrm {BOH} }}{c_{\mathrm {B^{+}} }}}\ {\Bigl (}={\frac {K\mathrm {_{W}} }{K\mathrm {_{B}} }}=K_{\mathrm {S} }{\Bigl )}}$

In these cases there is an equality of the salt hydrolysis with the definition according to Brønsted and Lowry with the base constant K _B , the acid constant K _S and the ion product , which are shown in the equations in brackets. ${\displaystyle K_{\text{W}}=c_{\mathrm {{H_{3}O}^{+}} }\cdot c_{\mathrm {OH^{-}} }}$

Limits of definition

In this definition, acids and bases are limited to water as a solvent. In the case of acids, the release of H ⁺ ions is postulated, which are hydrated. On closer inspection, however, oxonium ions (H ₃ O ⁺ ) are formed, as the solvent water is an important reaction partner.

Basic reactions of substances that cannot give off OH ^- ions are not described with this model. The basic reaction of ammonia or of element organic compounds in water cannot be described with this definition of a base. The focus of the reaction of an acid with a base is the neutralization reaction and the formation of salts, while according to the Brønsted model, which has largely replaced the Arrhenius model, the distinction between acidic, basic and neutral solutions is of secondary importance and buffer systems are better described to let. A distinction is made between salts, their ions in solution and molecules from the perspective of their conductivity. More modern considerations order these substances by their chemical reactivity, whereby non-reactive particles can be ignored.

Definition according to Brønsted and Lowry

Johannes Nicolaus Brønsted and Thomas Lowry described 1923 independently acid as a particle, the protons (H ⁺ - ion ) to a second reaction partners, the so-called base can transmit. In contrast to Arrhenius, however, bases and acids are no longer certain classes of substances, but particles that show certain properties in a reaction with H ⁺ ions:

• Particles that can donate protons are called proton donors or acids .
• On the other hand, particles that can take up protons are called proton acceptors or bases .

Acid-base reactions in which protons are transferred in the manner mentioned above are also called protolysis . Free protons (H ⁺ ), however, do not exist at any point in time: Every reaction of a partner as an acid necessarily requires the presence of a second partner as a base, to which the acid can transfer its protons :

${\displaystyle \mathrm {Y+HX\ \rightleftharpoons \ Y\cdots H^{+}\cdots X^{-}\rightleftharpoons \ HY^{+}+X^{-}} }$

Systems of this kind are also called conjugated or corresponding acid-base pairs , between which a chemical equilibrium is always established after a certain time . In the above reaction equation, HX and HY ^{+ are} the acids, Y and X ^- are the bases. What distinguishes both pairs is only their ability to release or absorb protons. As can also be seen, in a protolytic reaction, an acid (here HX) always produces its so-called conjugated or corresponding base (here X ^- ), and a base (here Y) always produces its so-called conjugated or corresponding acid (here HY ⁺ ) and vice versa.

If a chemical substance can both release and absorb protons, i.e. if it can act both as an acid and a base, one speaks of an ampholyte or the property of being amphoteric . The best known ampholyte is water, which allows the formation of both OH ^- and H ₃ O ⁺ :

${\displaystyle \mathrm {H_{2}O+H_{2}O\ \rightleftharpoons \ H_{3}O^{+}+OH^{-}} }$

The most important protolytic reactions in practice are accordingly reactions with water:

${\displaystyle \mathrm {HX+H_{2}O\ \rightleftharpoons \ H_{3}O^{+}+X^{-}} }$

The equilibrium position of this reaction is determined by the acid strength of HX, described numerically by its acid constant .

Examples of acid-base reactions according to Brønsted

• Reaction of hydrogen chloride (HCl) and water to form hydrochloric acid (HCl _(aq) ):

${\displaystyle \mathrm {HCl+H_{2}O\ \rightleftharpoons \ Cl^{-}+H_{3}O^{+}} }$

• Reaction of sulfuric acid and water:

${\displaystyle \mathrm {H_{2}SO_{4}+2\ H_{2}O\ \rightleftharpoons \ HSO_{4}^{-}+H_{3}O^{+}+H_{2}O\ \rightleftharpoons \ SO_{4}^{2-}+2\ H_{3}O^{+}} }$

• Reaction of ammonia and water:

${\displaystyle \mathrm {NH_{3}+H_{2}O\rightleftharpoons \ NH_{4}^{+}+OH^{-}} }$

• But Brønsted and Lowry's concept of acid also explains - in contrast to Arrhenius - the acid-base reaction of hydrogen chloride and ammonia gas to form ammonium chloride (NH ₄ Cl) despite the absence of water:

${\displaystyle \mathrm {HCl+NH_{3}\ \rightleftharpoons \ NH_{4}^{+}Cl^{-}} }$

Lewis definition

→ Main article : Lewis acid-base concept

Gilbert Newton Lewis published a treatise on his acid-base theory in 1923. Accordingly, a Lewis acid is an electrophilic electron pair acceptor and a Lewis base is an electron pair donor.

The Lewis acids include:

• Compounds with incomplete electron octet such as: B (CH ₃ ) ₃ , BF ₃ , AlCl ₃ , FeCl ₂
• Metal cations as central atoms in chemical complexes
• Molecules with polarized double bonds , such as CO ₂
• Halides with unsaturated coordination , for example SiCl ₄ or PF ₅

All Brønsted and Lowry bases are also Lewis bases.

Example of an acid-base reaction according to Lewis

${\displaystyle \mathrm {AlCl_{3}+Cl^{-}\ \rightleftharpoons \ AlCl_{4}^{-}} }$

The Lewis acid AlCl ₃ reacts with the Lewis base Cl ^- to form the Lewis acid-base adduct AlCl ₄ ^- .

Definition according to lux and flood

The focus of the concept established by Hermann Lux in 1939 and expanded by Håkon Flood in 1947 is the oxide ions instead of protons. This was set up in order to be able to describe acid-base reactions also in proton-free systems, as occurs in inorganic melts.

According to Lux and Flood, acids are oxide ion acceptors , bases are oxide ion donors . Non-metal oxides (e.g. SO ₂ , CO ₂ ) are considered to be acid anhydrides , since they react acidic in aqueous solution, and metal oxides (e.g. MgO, Fe ₂ O ₃ ) are base anhydrides , as they form hydroxide ions in aqueous solution.

Examples of acid-base reactions according to Lux and Flood

${\displaystyle \mathrm {MgO+CO_{2}\ \rightarrow \ MgCO_{3}} }$

The lux flood acid CO ₂ reacts with the lux flood base MgO.

${\displaystyle \mathrm {CaO+SiO_{2}\ \rightarrow \ CaSiO_{3}} }$

The lux flood acid SiO ₂ reacts with the lux flood base CaO.

Definition according to Uzhanovich

In 1939 the Russian scientist Michail Ussanowitsch established the following definition of the acid-base term:

“Acids are substances that can split off cations or accept anions or electrons.
Bases are substances that can split off anions or electrons or accept cations. "

- M. Ussanowitsch: J.allg.Chem. (USSR), 9 , 182 (1939)

This definition of terms includes the reactions according to the Lewis concept, but extends it again by the fact that the uptake or release of electrons is no longer restricted to electron pairs and thus also includes all traditional redox reactions in which a complete electron transfer takes place.

One point of criticism of this little-used theory is that it is too general and that the term acid-base reaction can therefore be applied to too many different types of reactions.

Pearson's concept of hard and soft acids and bases

→ Main article : HSAB concept

Ralph Pearson in 1963 developed the concept of hard and soft acids and bases ( Hard and Soft Acids and Bases , HSAB). It is said:

"Hard acids combine preferentially with hard bases and soft acids preferentially combine with soft bases."

The electronegativity and the polarizability of the observed particle are considered:

Type	properties	Examples
Hard acids	low electronegativity low polarizability	H + , Na + , K +
Hard bases	high electronegativity low polarizability	OH - , F - , SO 4 2−
Soft acids	low electronegativity high polarizability	Cu + , Ag + , I 2
Soft bases	high electronegativity high polarizability	I - , SCN - , R 2 S

The concept reflects tendencies, there are few absolutely hard or soft particles. However, it does help in estimating the stability of connections . So z. B. the softer Fe ²⁺ in nature as sulfide , while the harder Fe ^{3+ is present} as a hydroxide or oxide .

Comparison of the acid-base theories

Acids and bases are chemical counterparts whose opposite properties cancel each other out when they interact. The constitutional and functional characteristics of acids and bases are the subject of the acid-base definitions.

acid	base	author
contain oxygen	not further defined	(Lavoisier)
contain hydrogen	not further defined	(Davy)
contain hydrogen, by metal is replaceable	not further defined	(Liebig, 1838)
give off H + ions in water	give off OH - ions in water	(Arrhenius / Ostwald, 1884)
give off H + ions	take up H + ions	Brønsted
increase the concentration of the solvent's own cations	increase the concentration of lyations or decrease the concentration of lyonium ions	(Ebert / Konopik 1949)
cleaving cations or H + from or accept electrons to or anions	cleave electrons or anions from or take cations or H + to	Usanovich 1939
have electron vacancy, in which a pair of electrons to form a coordinate bond can be added	have a free electron pair that can be made available to form a coordinative bond	(Lewis, 1923)

Web links

Wikibooks: Inorganic Chemistry for Students / Acids - Bases - Reactions  - Learning and Teaching Materials

Wikibooks: General and Inorganic Chemistry / The Acid-Base Term  - Learning and Teaching Materials

Individual evidence

1. ^ ^{A b} Hans-Dieter Jakubke, Ruth Karcher (Ed.): Lexikon der Chemie. Spectrum Academic Publishing House, Heidelberg, 2001.
2. ↑ ^{a b c d} Karl-Heinz Lautenschläger, Werner Schröter, Joachim Teschner, Hildegard Bibrack, Taschenbuch der Chemie. 18th edition, Harri Deutsch, Frankfurt (Main), 2001.
3. ↑ Entry on salts. In: Römpp Online . Georg Thieme Verlag, accessed on June 20, 2014.
4. ^ AF Holleman , E. Wiberg , N. Wiberg : Textbook of Inorganic Chemistry . 102nd edition. Walter de Gruyter, Berlin 2007, ISBN 978-3-11-017770-1 , p. 208.
5. ↑ JN Brönsted: Some remarks on the concept of acids and bases . In: Recueil des Travaux Chimiques des Pays-Bas . tape 42 , no. 8 , January 1, 1923, ISSN  0165-0513 , p. 718–728 , doi : 10.1002 / recl.19230420815 ( wiley.com [accessed September 11, 2017]).
6. ↑ Lothar Kolditz (Ed.): Inorganikum. Deutscher Verlag der Wissenschaften, Berlin 1970, p. 423.
7. ↑ Gilbert Newton Lewis: Valence and the structure of atoms and molecules . Chemical Catalog Comp., New York 1923 ( gbv.de [accessed September 9, 2017]).
8. ↑ J. E, Huheey, E, A. Keiter, RL Keiter: Inorganic Chemistry , De Gruyter-Verlag 2014, ISBN 978-3-11-030433-6 , Chapter 9.1.2: Definition according to Lux and Flood
9. ↑ Lothar Kolditz (Ed.): Inorganikum. Deutscher Verlag der Wissenschaften, Berlin 1970, pp. 439–440.
10. ↑ G. Henrion: Science and Progress 16 (1966), p. 308.
11. ^ Acid-base theories. Comparison (PDF; 265 kB).