Acids

Acids are, in the narrower sense, all chemical compounds that are able to transfer protons (H ⁺ ) to a reaction partner - they can act as proton donors. In aqueous solution, the reactant is essentially water. There form oxonium ions (H ₃ O ⁺ ) and the pH of the solution is lowered with it. Acids react with so-called bases to form water and salts . A base is therefore the counterpart to an acid and is able to neutralize it .

In a broader sense, various acid-base concepts describe a much broader range of chemical reactions that can extend far beyond the reactions mentioned above.

history

Probably the oldest known acid (lat. Acidum ) is vinegar (lat. Acetum ), an approximately five percent aqueous solution of the compound acetic acid . Acids were something like vinegar ( oxos or acidus ). In alchemy , acid was considered a primal substance . The acidic properties were based on a single ursic acid that had an effect in various substances. Acids cause lime or carbonates to foam up, have a caustic effect and a sour taste . Until the end of the 13th century, no acids besides vinegar and other vegetable juices were known. Extended knowledge of acids was ascribed to the giver , who lived in the 8th century, but they are based on writings from the much younger, so-called pseudo-givers . Knowledge of various inorganic acids developed, all of which were ascribed to the varieties of the element water . From the 18th century onwards, they were called mineral acids .

A production of nitric acid (HNO ₃ ) was described by the pseudo-donors probably after the 13th century as "dissolving water" (aqua dissolutiva) or "strong water" (aqua fortis) . It was vitriol (but also see Chalkanthit with) Salpeter and alums heated to red heat. Nitrous gases escape and form nitric acid with water. This acid was also called septic fluid because it dissolves silver in a chemical reaction, but not gold . In 1648 J. R. Glauber described the production of concentrated (smoking) nitric acid (spiritus acidus nitri) by reaction with sulfuric acid.

Aqua regia , from today's point of view a mixture of nitric and hydrochloric acid , is likely to have been known at a similarly early stage, as it can be obtained relatively easily by converting nitric acid with the well-known salmiac . Aqua regia was the queen of all acids , which even gold , the king of metals , could not withstand.

Sulphurous acid (H ₂ SO ₃ ) has been known for a long time, as it can be obtained by burning elemental sulfur . Sulfuric acid was certainly involved in many processes. More precise processes for the preparation of sulfuric acid (H ₂ SO ₄ ) were only described in more detail by A. Libavius in 1597 : a) Annealing of iron or copper sulfate , the so-called iron or copper vitriol - see also the vitriol process - for oleum vitrioli ( fuming sulfuric acid) and b) combustion of sulfur and subsequent oxidation, which led to dilute solutions (oleum sulphuris) . A technical implementation succeeded in the 19th century with the lead chamber process .

Hydrochloric acid (HCl) was not known until the end of the 16th century. By igniting a mixture of salt and clay described A. Libavius a synthesis. B. Valentinus described the conversion of vitriol and table salt to aqua caustica , the caustic water. Intensive investigations by J. R. Glauber in the 17th century led to Glauber's salt spirit (Spiritus salis Glauberianus) , the very concentrated, so-called fuming hydrochloric acid or the gas hydrogen chloride .

The important term base as a phenomenological counterpart to acid was used in the 17th century by alchemists and chemists such as G. E. Stahl , R. Boyle and G. F. Rouelle , because "basic" substances formed the non-volatile basis for fixing volatile acids and the (caustic) effect of Could pick up acids. Basic steps in chemistry were taken by A. L. Lavoisier in the 18th century, who assigned certain properties to certain chemical compounds. He thought that acids were always made from non-metal oxides and water, and bases were made from metal oxides and water. Humphry Davy found a counterexample ( hydrogen chloride ) in 1808 . J. von Liebig saw acids as hydrogen compounds that can be converted into salts by metals .

In 1887 S. Arrhenius defined acids as substances that dissociate when dissolving in water, releasing protons (H ⁺ ), and bases as substances that dissociate when dissolving in water, releasing hydroxide ions (OH ^- ). If acids and bases are combined, they neutralize each other and form water. However, the theory was still insufficient, since compounds without oxygen were not included: Ammonia also neutralizes an acid. In 1923, JN Brønsted and Martin Lowry independently described what is still the most important definition of acids and bases today. They form the basis of the acidic explanations below.

What are acids

Without going into more detail about the various acid-base concepts , a possible and common approach will be described here as an introduction. The presence and certain properties of the water are usually closely related to acids, often without explicit mention . Pure water is subject to what is known as autoprotolysis . Oxonium ions (H ₃ O ⁺ ) and hydroxide ions (OH ^- ) are formed from the water in very small and equal amounts :

{\ displaystyle \ mathrm {1. \ H_ {2} O + H_ {2} O \ \ rightleftharpoons \ H_ {3} O ^ {+} + OH ^ {-}}}

This reaction equation for water shows the property of an acid, namely the ability to form H ₃ O ⁺ ions in water. At the same time, OH ^- ions are formed in water - one of the properties that a base can have. However, water is neither called a base nor an acid and its behavior is called neutral . This relates to the pH value , which indicates the concentration of H ₃ O ⁺ ions in water. Pure water has a pH value of 7, which is a very small concentration. Like all the reactions described in this section, this reaction is an equilibrium reaction: The formation of the ions and their combination to form water takes place continuously and with the same frequency. So neutral does not mean that nothing happens.

Organic acids
Carboxylic acid
Sulfonic acid
R is an organyl group , e.g. B. a methyl or phenyl group. The functional groups are marked in blue .

Acids can be used to describe chemical compounds that interact in a certain way with water. They have hydrogen atoms that are bound like ions (ionogenic). Pure acetic acid (H ₃ C-COOH) reacts with water and forms further H ₃ O ⁺ ions. If such a reaction occurs, a compound can be called an acid. In addition to the oxonium ion, the acetate anion H ₃ C-COO ^{- is also formed} :

{\ displaystyle \ mathrm {2. \ H_ {3} C {-} COOH + H_ {2} O \ \ rightleftharpoons \ H_ {3} C {-} COO ^ {-} + H_ {3} O ^ {+ }}}

Acetic acid is a carboxylic acid and is a rather weak acid compared to inorganic acids such as hydrogen chloride . In aqueous solution, a good part of the molecules are undissociated as H ₃ C-COOH. Here, too, an equilibrium is quickly established . For this reason, the above reaction equation (2) can be read from right to left with equal justification. An acetate anion reacts with an oxonium ion to form water and acetic acid. In this reading direction a basic reaction takes place: the conversion of hydroxonium ions into water molecules. If a suitable amount of acetate anions is added to an acetic acid solution , for example in the form of the readily soluble sodium acetate , the acidic property of acetic acid can be completely compensated for by the basic property of the acetate ion. The aqueous solution is neutralized. Here, too, neutral does not mean that nothing happens in the solution. Only the concentration of the H ₃ O ⁺ ions is as low as in pure water.

In addition, the basic reaction that occurs when sodium acetate is dissolved in pure water should now be considered (the Na ⁺ cation is omitted in the reaction equation):

{\ displaystyle \ mathrm {3. \ H_ {3} C {-} COO ^ {-} + H_ {2} O \ \ rightleftharpoons \ OH ^ {-} + H_ {3} C {-} COOH}}

Here form hydroxide ions (OH ^- ). If a suitable amount of aqueous acetic acid solution is added to this acetate solution, the solution becomes neutral . The equilibrium (1) is established between H ₃ O ⁺ and OH ^- , which was initially presented as the basic property of water and is shown here as equation (1a) the other way around.

{\ displaystyle \ mathrm {1a) \ H_ {3} O ^ {+} + OH ^ {-} \ \ rightleftharpoons \ H_ {2} O + H_ {2} O}}

An acid can be understood as a chemical compound that has the ability to form H ₃ O ⁺ ions in aqueous solution or to convert OH ^- ions to water molecules. A base, on the other hand, has the ability in aqueous solution to form OH ^- ions or to convert H ₃ O ⁺ to H ₂ O.

Many substances referred to as acids are aqueous solutions from the outset and cannot simply be understood as chemical compounds that have ionically bonded hydrogen atoms. Hydrochloric acid is an aqueous solution of the gas hydrogen chloride (HCl) and is considered a strong acid. In this solution - before any practical use of the acid - the equilibrium reaction (4) is already present, in which the equilibrium is almost completely on the right side.

{\ displaystyle \ mathrm {4. \ HCl + H_ {2} O \ \ rightleftharpoons \ H_ {3} O ^ {+} + Cl ^ {-}}}

Hydrogen chloride has long since played out its potential to be an acid, and H ₃ O ⁺ ions have formed. The chemical effects that occur when hydrochloric acid is used in practice are due to reactions of the H ₃ O ⁺ ions. The acid is the H ₃ O ⁺ ion. The measure for the acid content ( concentration ) is the pH value , while with weaker acids such as acetic acid, the measure for the acid strength , the pKa value, is in the foreground. Strong and weak acids differ in their tendency to “like” or “not like” H ₃ O ⁺ ions in water. These tendencies are described in more detail in the section Acid-Base Balance.

The chemical effect when an acid is used in practice is mostly due to H ₃ O ⁺ ions. One can consider the oxonium ions as the most important acid.

Acid-base balance

During protolysis , a reaction partner (usually water) picks up the proton given off by the acid. This must be distinguished from the redox reactions in which electron transfers take place.

The general equilibrium reaction of an acid HA in aqueous solution is:

{\ displaystyle \ mathrm {5. \ HA + H_ {2} O \ \ rightleftharpoons \ H_ {3} O ^ {+} + A ^ {-}}}

The acids differ in their tendency to transfer H ⁺ ions to water. This is known as the acid strength K _s and indicates the equilibrium constant ( acid constant ) of the acid reaction. The acid constant is often given in the form of the p K _s value, which is defined as the negative decadic logarithm of the acid constant.

{\ displaystyle K _ {\ mathrm {S}} = {\ frac {c (\ mathrm {H} _ {3} \ mathrm {O} ^ {+}) \ cdot c (\ mathrm {A} ^ {-} )} {c (\ mathrm {HA})}}}

{\ displaystyle \ mathrm {p} K _ {\ mathrm {S}} = - \ log K _ {\ mathrm {S}}}

Acids with a high K _s value (small p K _s value) are strong acids. If the pH of a solution that contains an acid is two units below the p K _s value, only one hundredth of the H ₃ O ⁺ ions are formed.

Multi-protonic acids

Acids that can split off several protons are called polyprotonic acids or polybasic acids . Sulfuric acid (H ₂ SO ₄ ) is a biprotonic (also diprotonic) acid , phosphoric acid (H ₃ PO ₄ ) is a three- protonic (also triprotonic) acid. The effort to release the individual protons (protolysis) varies in size and can be described by the acid constant (K _s ). The following generally applies to the individual protolysis steps: K _s (I)> K _s (II)> K _s (III) (or p K _s (I) <p K _s (II) <p K _s (III)).

The following applies to phosphoric acid:

${\ displaystyle \ mathrm {H} _ {3} \ mathrm {PO_ {4} + H_ {2} O \ \ rightleftharpoons \ H_ {2} PO_ {4} ^ {-} + H_ {3} O ^ {+ }}}$	${\ displaystyle K _ {\ mathrm {S}} = 7 {,} 4 \ cdot 10 ^ {- 3}}$	${\ displaystyle \ mathrm {p} K _ {\ mathrm {S}} = 2 {,} 13 \}$
${\ displaystyle \ mathrm {H_ {2} PO_ {4} ^ {-} + H_ {2} O \ \ rightleftharpoons \ HPO_ {4} ^ {2 -} + H_ {3} O ^ {+}}}$	${\ displaystyle K _ {\ mathrm {S}} = 6 {,} 3 \ cdot 10 ^ {- 8}}$	${\ displaystyle \ mathrm {p} K _ {\ mathrm {S}} = 7 {,} 20 \}$
${\ displaystyle \ mathrm {HPO_ {4} ^ {2 -} + H_ {2} O \ \ rightleftharpoons \ PO_ {4} ^ {3 -} + H_ {3} O ^ {+}}}$	${\ displaystyle K _ {\ mathrm {S}} = 4 {,} 4 \ cdot 10 ^ {- 13}}$	${\ displaystyle \ mathrm {p} K _ {\ mathrm {S}} = 12 {,} 36 \}$

pK _S values important acids

See p K _s and p K _b values of some compounds .

properties

The properties of acids, especially the dangers they pose, vary widely. Examples are nitric acid (main risk: corrosive), hydrocyanic acid (strong poison) and picric acid (an explosive).

Acids particularly attack base metals and lime . But clothing, skin and eyes (generally all organic materials) run the risk of being destroyed by the acid on contact.
There are strong and weak acids. Hydrogen chloride is a strong acid and completely dissociates in water. The aqueous solution is called hydrochloric acid . Acetic acid is a weaker acid and only partially dissociates in water.
Acids can be diluted with water; depending on the dilution, their effect is significantly weaker. The dilution of concentrated acids is an exothermic reaction . So there is heat. The acid solution can splash away in an uncontrolled manner, especially when diluting concentrated sulfuric acid. Therefore, when diluting, the rule to add the acid to the water does not apply the other way around: "First the water, then the acid, otherwise the monster will happen." But even with correct mixing, make sure that the concentrated acid is slow and careful is added to the water.
It is a common misconception that acids are always liquids. Well-known representatives of acids that are purely solid are vitamin C and citric acid , a gaseous acid is, for example, hydrogen chloride .
Aqueous solutions of acids cause indicators to change color , for example they turn blue litmus paper red.
The "opponents of the acids" are the bases (base solution = lye). You can neutralize acids. Bases are also corrosive and attack many other substances that do not necessarily react with acids.
When dissolved in water, acids conduct electricity. Here, a is carried out electrolysis , at the at the cathode (the negative pole), and hydrogen at the anode form (the positive pole) of the neutralized substance of the acid anion, wherein the hydrochloric acid z. B. chlorine . A reduction (electron uptake) takes place at the cathode and oxidation (electron release) takes place at the anode .

Acid-base reactions without water

Analogous to the acid-base reactions that take place in aqueous solutions and with the participation of water, there are reactions in other media. In anhydrous ethanol , a reaction takes place with hydrogen chloride , in which ethanol takes on the role of a base:

{\ displaystyle \ mathrm {H_ {3} C {-} CH_ {2} {-} OH + HCl \ \ rightleftharpoons \ H_ {3} C {-} CH_ {2} {-} OH_ {2} ^ {+ } + Cl ^ {-}}}

In the gas phase, the gases ammonia and hydrogen chloride react to form the salt ammonium chloride .

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

In addition to water, other sufficiently polar solvents can also act as reactants in acid-base reactions. A good example is the autoprotolysis of liquid ammonia:

{\ displaystyle \ mathrm {NH_ {3} + NH_ {3} \ \ rightleftharpoons \ NH_ {4} ^ {+} + NH_ {2} ^ {-}}}

Examples of acids

Important acids are:

Sulfuric acid : H ₂ SO ₄ (industrial use, acid rain )
Hydrochloric acid : HCl (industrial use)
Silicic acid : H ₄ SiO ₄
Phosphoric acid : H ₃ PO ₄ (food industry, including cola , DNA )
Carbon dioxide : H ₂ CO ₃ (food industry, technology, atmosphere )
Acetic acid : CH ₃ COOH (salad preparation in the kitchen, food industry)
Benzoic acid ( preservatives for food)
Hydrofluoric acid : HF (computer chip manufacture)
Nitric acid : HNO ₃ (industrial use)

Salts of multi-protonic acids can also act as acids (“acidic salts”), for example

literature

Historical development of acids:

Claus Priesner , Karin Figala : Alchemy: Lexicon of a hermetic science. Beck, Munich 1998, ISBN 3-406-44106-8
V. Karpenko, JA Norris: Vitriol in the History of Chemistry. Chem. Listy, Volume 96, 2002, pages: 997-1005, PDF
Donors in Britannica

Web links

Wiktionary: acid - explanations of meanings, word origins, synonyms, translations

Individual evidence

↑ G. Jander, E. Blasius: Introduction to the inorganic-chemical practical course . 12th, revised edition. Hirtzel Verlag, Stuttgart 1987, ISBN 3-7776-0433-X , p. 5 .

[1] G. Jander, E. Blasius: Introduction to the inorganic-chemical practical course . 12th, revised edition. Hirtzel Verlag, Stuttgart 1987, ISBN 3-7776-0433-X , p. 5 .

Acids

contents

history