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An acid (often represented by the generic formula HA) is traditionally considered any chemical compound that when dissolved in water, gives a solution with a pH of less than 7.0. That approximates the modern definition of Brønsted and Lowry, who defined an acid as a compound which donates a hydrogen ion (H+) to another compound (called a base). Common examples include acetic acid (in vinegar) and sulfuric acid (used in car batteries).
Definitions of acids and bases The word "acid" comes from the Latin acidus meaning "sour," but in chemistry the term acid has a more specific meaning. There are three common ways to define an acid, namely, the Arrhenius, the Brønsted-Lowry and the Lewis definitions, in order of increasing generality. Although not the most general theory, the Brnsted-Lowry definition is the most widely used definition. The strength of an acid may be understood by this definition by the stability of hydronium and the solvated conjugate base upon dissociation. Increasing stability of the conjugate base will increase the acidity of a compound. This concept of acidity is used frequently for organic acids such as carboxylic acid. The molecular orbital description, where the unfilled proton orbital overlaps with a lone pair, is connected to the Lewis definition. Solutions of weak acids and salts of their conjugate bases form buffer solutions. Acid/base systems are different from redox reactions in that there is no change in oxidation state. Properties Generally, acids have the following properties: Strong acids and most concentrated acids are dangerous, causing severe burns for even minor contact. Generally, acid burns are treated by rinsing the affected area abundantly with running water (15 minutes) and followed up with immediate medical attention. In addition to dangers from the acidity, even dilute solutions of weak acids may also be dangerous, due to toxic or other effects of the ions involved. Nomenclature Acids are named according to their anions. That ionic suffix is dropped and replaced with a new suffix (and sometimes prefix), according to the table below. For example, HCl has chloride as its anion, so the -ide suffix makes it take the form hydrochloric acid. Chemical characteristics In water the following equilibrium occurs between an acid (HA) and water, which acts as a base: HA(aq) H3O+(aq) + A-(aq) The acidity constant (or acid dissociation constant) is the equilibrium constant for the reaction of HA with water: Strong acids have large Ka values (i.e. the reaction equilibrium lies far to the right; the acid is almost completely dissociated to H3O+ and A-). Strong acids include the heavier hydrohalic acids: hydrochloric acid (HCl), hydrobromic acid (HBr), and hydroiodic acid (HI). (However, hydrofluoric acid, HF, is relatively weak.) For example, the Ka value for hydrochloric acid (HCl) is 107. Weak acids have small Ka values (i.e. at equilibrium significant amounts of HA and A− exist together in solution; modest levels of H3O+ are present; the acid is only partially dissociated). For example, the Ka value for acetic acid is 1.8 x 10-5. Most organic acids are weak acids. Oxoacids, which tend to contain central atoms in high oxidation states surrounded by oxygen may be quite strong or weak. Nitric acid, sulfuric acid, and perchloric acid are all strong acids, whereas nitrous acid, sulfurous acid and hypochlorous acid are all weak. Note on terms used: 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) Ka1 HA−(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) Ka1 H2A−(aq) + H2O(l) H3O+(aq) + HA2−(aq) Ka2 HA2−(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 equal amounts of 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 equivalence point when the equivalent 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 Acidification is the process whereby air pollution – mainly ammonia, sulphur dioxide and nitrogen oxides – is converted into acid substances. This ‘acid rain’ is best known for the damage it causes to forests and lakes. Less well known are the many ways it damages freshwater and coastal ecosystems, soils and even ancient historical monuments, or the heavy metals these acids help release into groundwater. Sulphur dioxide and the nitrogen oxides are mainly emitted by burning fossil fuels. As some of the reports in this section show, the 1990s saw these emissions drop substantially, thanks to a combination of European Directives forcing the installation of desulphurisation 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. Chemistry Environment | ||||||||
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