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Chemical characteristics

In water, the following equilibrium occurs between an acid (HA) and the water, which acts as a base:

HA(aq) + H2O(l) ⇌ H3O+(aq) + A-(aq)

The acidity constant (or acid dissociation constant) is the equilibrium constant that indicates the degree to which hydrogen ions dissociate from an acid.

Strong acids are those that almost completely dissociate in water. They have large Ka values; hence, the acid dissociation equilibrium lies 100% to the right, which means that there are mostly H3O+ and A- ions in solution with a very minute amount of undissociated HA molecules.

Common strong acids are perchloric acid (HClO4), hydrochloric acid (HCl), hydrobromic acid (HBr), hydroiodic acid (HI), nitric acid (HNO3), and sulphuric acid (H2SO4). For example, the Ka value for hydrochloric acid (HCl) is 107.

Weak acids are those that partially dissociate in water. They have small Ka values; therefore, only a small percent of protons are donated to water, keeping the acid dissociation equilibrium to the left. The solution mainly contains undissociated HA molecules with very little H3O+ and A- ions. Common weak acids are nitrous acid (HNO2), hydrofluoric acid (HF), and acetic acid (CH3CO2H). For example, the Ka value for acetic acid is 1.8 x 10-5.

Note on terms used:

  • The terms "hydrogen ion" and "proton" are used interchangeably; both refer to H+.
  • In aqueous solution, the water is protonated to form hydronium ion, H3O+(aq). This is often abbreviated as H+(aq) even though the symbol is not chemically correct.
  • The term "hydroxide ion" (OH-) is also called hydroxyl ion.
  • The strength of an acid is measured by its acid dissociation constant (Ka) or equivalently its pKa (pKa= - log(Ka)).
  • The pH of a solution is a measurement of the concentration of hydronium ions. This will depend of the concentration and nature of acids and bases in solution.

Polyprotic acids

Polyprotic acids are able to donate more than one proton per acid molecule, in contrast to monoprotic acids that only donate one proton per molecule. Specific types of polyprotic acids have more specific names, such as diprotic acid (two potential protons to donate) and triprotic acid (three potential protons to donate).

A monoprotic acid can undergo one dissociation (sometimes called ionization) as follows and simply has one acid dissociation constant as shown above:

HA(aq) + H2O(l) ⇌ H3O+(aq) + A(aq) Ka

A diprotic acid (here symbolized by H2A) can undergo one or two dissociations depending on the pH. Each dissociation has its own dissociation constant, Ka1 and Ka2.

H2A(aq) + H2O(l) ⇌ H3O+(aq) + HA(aq) Ka1HA(aq) + H2O(l) ⇌ H3O+(aq) + A2−(aq) Ka2

The first dissociation constant is typically greater than the second; i.e., Ka1 > Ka2 . For example, sulfuric acid (H2SO4) can donate one proton to form the bisulfate anion (HSO4), for which Ka1 is very large; then it can donate a second proton to form the sulfate anion (SO42−), wherein the Ka2 is intermediate strength. The large Ka1 for the first dissociation makes sulfuric a strong acid. In a similar manner, the weak unstable carbonic acid (H2CO3) can lose one proton to form bicarbonate anion (HCO3) and lose a second to form carbonate anion (CO32−). Both Ka values are small, but Ka1 > Ka2 .

A triprotic acid (H3A) can undergo one, two, or three dissociations and has three dissociation constants, where Ka1 > Ka2 > Ka3 .

H3A(aq) + H2O(l) ⇌ H3O+(aq) + H2A(aq) Ka1H2A(aq) + H2O(l) ⇌ H3O+(aq) + HA2−(aq) Ka2HA2−(aq) + H2O(l) ⇌ H3O+(aq) + A3−(aq) Ka3

An inorganic example of a triprotic acid is orthophosphoric acid (H3PO4), usually just called phosphoric acid. All three protons can be successively lost to yield H2PO4, then HPO42−, and finally PO43− , the orthophosphate ion, usually just called phosphate. An organic example of a triprotic acid is citric acid, which can successively lose three protons to finally form the citrate ion. Even though the positions of the protons on the original molecule may be equivalent, the successive Ka values will differ since it is energetically less favorable to lose a proton if the conjugate base is more negatively charged.

Neutralization

Neutralization is the reaction between an acid and a base, producing a salt and water; for example, hydrochloric acid and sodium hydroxide form sodium chloride and water:

HCl(aq) + NaOH(aq) → H2O(l) + NaCl(aq)

Neutralization is the basis of titration, where a pH indicator shows an equivalence point when the same number of moles of a base have been added to an acid.

Weak acid/weak base equilibria

In order to lose a proton, it is necessary that the pH of the system rise above the pKa of the protonated acid. The decreased concentration of H+ in that basic solution shifts the equilibrium towards the conjugate base form (the deprotonated form of the acid). In lower-pH (more acidic) solutions, there is a high enough H+ concentration in the solution to cause the acid to remain in its protonated form, or to protonate its conjugate base (the deprotonated form).

Acidification of the environment

Acidification is the process whereby a compound is added to a solution, leading to a drop in the pH of the solution. One example is when the pollution of air-mainly sulfur dioxide and nitrogen oxides-is converted into acidic substances.

This "acid rain" is best known for the damage it causes to forests and lakes. It also damages freshwater and coastal ecosystems, soils, and even ancient historical monuments.

Sulfur dioxide and the nitrogen oxides are mainly emitted by burning fossil fuels. The 1990s saw these emissions drop substantially, thanks to a combination of European Directives forcing the installation of desulfurisation systems, the move away from coal as a fossil fuel, and major economic restructuring in the new German Lander.

Acidification is nevertheless still a major environmental problem in Europe. It is a cross-border issue, requiring coordinated initiatives across countries and sectors. This section brings together the EEA's reports on the scale of the problem and the effectiveness of the solutions tried to date.1

Footnotes

  1. ↑ European Environment Agency. Acidification. Retrieved May 17, 2007.

References

  • Brown, Theodore E., H. Eugene LeMay, and Bruce E Bursten. 2005. Chemistry: The Central Science (10th Edition). Upper Saddle River, NJ: Prentice Hall. ISBN 0131096869
  • Corwin, C. H., 2001. Introductory Chemistry Concepts & Connections (3rd ed.). Upper Saddle River, NJ: Prentice Hall. ISBN 0130874701
  • McMurry, J. and R. C. Fay. 2004. Chemistry (4th ed.). Upper Saddle River, NJ: Prentice Hall. ISBN 0131402080
  • Moore, J. W., C. L. Stanitski, and P. C. Jurs,. 2002. Chemistry The Molecular Science. New York: Harcourt College.
  • Oxlade, Chris. 2002. Acids and Bases. Heinemann Library. ISBN 1588101940

External links

All links retrieved November 3, 2019.

  • Curtipot: Acid-Base equilibria diagrams, pH calculation and titration curves simulation and analysis (freeware).

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