, a typical metal, reacting with
hydrochloric acid, a typical acid]]
An
acid (from the
Latin acidus/acēre meaning
sour) is a substance which reacts with a
base. Commonly, acids can be identified as tasting sour, reacting with metals such as
calcium, and bases like
sodium carbonate. Aqueous acids have a pH of less than 7, where an acid of lower pH is typically stronger, and turn blue
litmus paper red. Chemicals or substances having the property of an acid are said to be
acidic.
Common examples of acids include acetic acid (in vinegar), sulfuric acid (used in car batteries), and tartaric acid (used in baking). As these three examples show, acids can be solutions, liquids, or solids. Gases such as hydrogen chloride can be acids as well. Strong acids and some concentrated weak acids are corrosive, but there are exceptions such as carboranes and boric acid.
There are three common definitions for acids: the Arrhenius definition, the Brønsted-Lowry definition, and the Lewis definition. The Arrhenius definition states that acids are substances which increase the concentration of hydronium ions (H3O+) in solution. The Brønsted-Lowry definition is an expansion: an acid is a substance which can act as a proton donor. Most acids encountered in everyday life are aqueous solutions, or can be dissolved in water, and these two definitions are most relevant. The reason why pHs of acids are less than 7 is that the concentration of hydronium ions is greater than 10−7 moles per liter. Since pH is defined as the negative logarithm of the concentration of hydronium ions, acids thus have pHs of less than 7. By the Brønsted-Lowry definition, any compound which can easily be deprotonated can be considered an acid. Examples include alcohols and amines which contain O-H or N-H fragments.
In chemistry, the Lewis definition of acidity is frequently encountered. Lewis acids are electron-pair acceptors. Examples of Lewis acids include all metal cations, and electron-deficient molecules such as boron trifluoride and aluminium trichloride. Hydronium ions are acids according to all three definitions. Interestingly, although alcohols and amines can be Brønsted-Lowry acids as mentioned above, they can also function as Lewis bases due to the lone pairs of electrons on their oxygen and nitrogen atoms.
Definitions and concepts
Modern definitions are concerned with the fundamental chemical reactions common to all acids.
Arrhenius acids
The Swedish chemist
Svante Arrhenius attributed the properties of acidity to
hydrogen in 1884. An
Arrhenius acid is a substance that increases the concentration of the
hydronium ion, H
3O
+, when dissolved in water. This definition stems from the equilibrium dissociation of water into hydronium and
hydroxide (OH
−) ions:
: H
2O(
l) + H
2O(
l) H
3O
+(
aq) + OH
−(
aq)
In pure water the majority of molecules exist as H
2O, but a small number of molecules are constantly dissociating and re-associating. Pure water is neutral with respect to acidity or basicity because the concentration of hydroxide ions is always equal to the concentration of hydronium ions. An Arrhenius
base is a molecule which increases the concentration of the hydroxide ion when dissolved in water. Note that chemists often write H
+(
aq) and refer to the
hydrogen ion when describing acid-base reactions but the free hydrogen nucleus, a
proton, does not exist alone in water, it exists as the hydronium ion, H
3O
+.
Brønsted-Lowry acids
While the Arrhenius concept is useful for describing many reactions, it is also quite limited in its scope. In 1923 chemists
Johannes Nicolaus Brønsted and
Thomas Martin Lowry independently recognized that acid-base reactions involve the transfer of a proton. A
Brønsted-Lowry acid (or simply Brønsted acid) is a species that donates a proton to a Brønsted-Lowry base.
|-
|HF
|
hydrofluoric acid
|3.17
|-
|H
2O
|water
|15.7
|-
|NH
3
|
ammonia
|38
|-
|CH
4
|
methane
|48
|}
The electronegative element need not be directly bonded to the acidic hydrogen to increase its acidity. An electronegative atom can pull electron density out of an acidic bond through the
inductive effect. The electron-withdrawing ability diminishes quickly as the electronegative atom moves away from the acidic bond. The effect is illustrated by the following series of
halogenated butanoic acids.
Chlorine is more electronegative than
bromine and therefore has a stronger effect. The hydrogen atom bonded to the oxygen is the acidic hydrogen. Butanoic acid is a
carboxylic acid.
{|class="wikitable" border="1"
|-
!Structure
!Name
!p
Ka
|-
|
|
butanoic acid or
butyric acid
|≈4.8
|-
|
|4-chlorobutanoic acid
|4.5
|-
|
|3-chlorobutanoic acid
|≈4.0
|-
|
|2-bromobutanoic acid
|2.93
|-
|
|2-chlorobutanoic acid
|2.86
|}
As the chlorine atom moves further away from the acidic O—H bond, its effect diminishes. When the chlorine atom is just one carbon removed from the carboxylic acid group the acidity of the compound increases significantly, compared to butanoic acid (a.k.a.
butyric acid). However, when the chlorine atom is separated by several bonds the effect is much smaller. Bromine is much more electronegative than either carbon or hydrogen, but not as electronegative as chlorine, so the p
Ka of 2-bromobutanoic acid is slightly greater than the p
Ka of 2-chlorobutanoic acid.
(HClO4) is an oxoacid and a strong acid.]]The number of electronegative atoms adjacent an acidic bond also affects acid strength. Oxoacids have the general formula HOX where X can be any atom and may or may not share bonds to other atoms. Increasing the number of electronegative atoms or groups on atom X decreases the electron density in the acidic bond, making the loss of the proton easier. Perchloric acid is a very strong acid (pKa ≈ -8) and completely dissociates in water. Its chemical formula is HClO4 and it comprises a central chlorine atom with three chlorine-oxygen double bonds (Cl=O) and one chlorine-oxygen single bond (Cl—O). The singly bonded oxygen bears an extremely acidic hydrogen atom which is easily abstracted. In contrast, chloric acid (HClO3) is a weaker acid, though still quite strong (pKa = -1.0), while chlorous acid (HClO2, pKa = +2.0) and hypochlorous acid (HClO, pKa = +7.53) acids are weak acids.
Carboxylic acids are organic acids that contain an acidic hydroxyl group and a carbonyl (C=O bond). Carboxylic acids can be reduced to the corresponding alcohol; the replacement of an electronegative oxygen atom with two electropositive hydrogens yields a product which is essentially non-acidic. The reduction of acetic acid to ethanol using LiAlH4 (lithium aluminium hydride or LAH) and ether is an example of such a reaction.
:
The pKa for ethanol is 16, compared to 4.76 for acetic acid.
Atomic radius and bond strength
Another factor that contributes to the ability of an acid to lose a proton is the strength of the
bond between the acidic hydrogen and the atom that bears it. This, in turn, is dependent on the size of the atoms sharing the bond. For an acid HA, as the size of atom A increases, the strength of the bond decreases, meaning that it is more easily broken, and the strength of the acid increases. Bond strength is a measure of how much
energy it takes to break a bond. In other words, it takes less energy to break the bond as atom A grows larger, and the proton is more easily removed by a base. This partially explains why hydrofluoric acid is considered a weak acid while the other hydrohalic acids (HCl, HBr, HI) are strong acids. Although fluorine is more electronegative than the other halogens, its
atomic radius is also much smaller, so it shares a stronger bond with hydrogen. Moving down a column on the periodic table atoms become less electronegative but also significantly larger, and the size of the atom tends to dominate its acidity when sharing a bond to hydrogen.
Hydrogen sulfide, H
2S, is a stronger acid than water, even though oxygen is more electronegative than sulfur. Just as with the halogens, this is because sulfur is larger than oxygen and the H—S bond is more easily broken than the H—O bond.
Chemical characteristics
Monoprotic acids
Monoprotic acids are those acids that are able to donate one
proton per molecule during the process of
dissociation (sometimes called ionization) as shown below (symbolized by HA):
:::::HA(aq) + H2O(l) H3O+(aq) + A−(aq) Ka
Common examples of monoprotic acids in mineral acids include hydrochloric acid (HCl) and nitric acid (HNO3). On the other hand, for organic acids the term mainly indicates the presence of one carboxylic acid group and sometimes these acids are known as monocarboxylic acid. Examples in organic acids include formic acid (HCOOH), acetic acid (CH3COOH) and benzoic acid (C6H5COOH).
Polyprotic acids
Polyprotic acids, also known as polybasic 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 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.
Although the subsequent loss of each hydrogen ion is less favorable, all of the conjugate bases are present in solution. The fractional concentration, α (alpha), for each species can be calculated. For example, a generic diprotic acid will generate 3 species in solution: H2A, HA-, and A2-. The fractional concentrations can be calculated as below when given either the pH (which can be converted to the [H+]) or the concentrations of the acid with all its conjugate bases:
A pattern is observed in the above equations and can be expanded to the general n -protic acid that has been deprotonated i -times:
where K0 = 1 and the other K-terms are the dissociation constants for the acid.
Neutralization
(in
beaker) reacting with
ammonia fumes to produce
ammonium chloride (white smoke).]]
Neutralization is the reaction between an acid and a base, producing a
salt and neutralized base; 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. It is often wrongly assumed that neutralization should result in a solution with pH 7.0, which is only the case with similar acid and base strengths during a reaction.
Neutralization with a base weaker than the acid results in a weakly acidic salt. An example is the weakly acidic ammonium chloride, which is produced from the strong acid hydrogen chloride and the weak base ammonia. Conversely, neutralizing a weak acid with a strong base gives a weakly basic salt, e.g. sodium fluoride from hydrogen fluoride and sodium hydroxide.
Weak acid/weak base equilibria
In order to lose a proton, it is necessary that the pH of the system rise above the p
Ka 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).
Solutions of weak acids and salts of their conjugate bases form buffer solutions.
Applications of acids
There are numerous uses for acids. Acids are often used to remove rust and other corrosion from metals in a process known as
pickling. They may be used as an electrolyte in a
wet cell battery, such as
sulfuric acid in a
car battery.
Strong acids, sulfuric acid in particular, are widely used in mineral processing. For example, phosphate minerals react with sulfuric acid to produce phosphoric acid for the production of phosphate fertilizers, and zinc is produced by dissolving zinc oxide into sulfuric acid, purifying the solution and electrowinning.
In the chemical industry, acids react in neutralization reactions to produce salts. For example, nitric acid reacts with ammonia to produce ammonium nitrate, a fertilizer. Additionally, carboxylic acids can be esterified with alcohols, to produce esters.
Acids are used as additives to drinks and foods, as they alter their taste and serve as preservatives. Phosphoric acid, for example, is a component of cola drinks. Acetic acid is used in day to day life as vinegar. Carbonic acid is an important part of some cola drinks and soda.
Citric acid is used as a preservative in sauces and pickles.
Tartaric acid is an important component of some commonly used foods like unripened mangoes and tamarind. Natural fruits and vegetables also contain acids. Citric acid is present in oranges, lemon and other citrus fruits. Oxalic acid is present in tomatoes, spinach, and especially in carambola and rhubarb; rhubarb leaves and unripe carambolas are toxic because of high concentrations of oxalic acid.
Ascorbic acid (Vitamin C) is an essential vitamin required in our body and is present in such foods as amla, lemon, citrus fruits, and guava.
Certain acids are used as drugs. Acetylsalicylic acid (Aspirin) is used as a pain killer and for bringing down fevers.
Acids play very important roles in the human body. The hydrochloric acid present in our stomach aids in digestion by breaking down large and complex food molecules. Amino acids are required for synthesis of proteins required for growth and repair of our body tissues. Fatty acids are also required for growth and repair of body tissues. Nucleic acids are important for the manufacturing of DNA, RNA and transmission of characters to offspring through genes. Carbonic acid is important for maintenance of pH equilibrium in the body.
Acid catalysis
Acids are used as
catalysts in industrial and organic chemistry; for example,
sulfuric acid is used in very large quantities in the
alkylation process to produce gasoline. Strong acids, such as sulfuric, phosphoric and hydrochloric acids also effect
dehydration and
condensation reactions. In biochemistry, many
enzymes employ acid catalysis.
Biological occurrence
.]]Many biologically important molecules are acids.
Nucleic acids, which contain acidic
phosphate groups, include
DNA and
RNA. Nucleic acids contain the genetic code that determines many of an organism's characteristics, and is passed from parents to offspring. DNA contains the chemical blueprint for the synthesis of
proteins which are made up of
amino acid subunits.
Cell membranes contain
fatty acid esters such as
phospholipids.
An α-amino acid has a central carbon (the α or alpha carbon) which is covalently bonded to a carboxyl group (thus they are carboxylic acids), an amino group, a hydrogen atom and a variable group. The variable group, also called the R group or side chain, determines the identity and many of the properties of a specific amino acid. In glycine, the simplest amino acid, the R group is a hydrogen atom, but in all other amino acids it is contains one or more carbon atoms bonded to hydrogens, and may contain other elements such as sulfur, oxygen or nitrogen. With the exception of glycine, naturally occurring amino acids are chiral and almost invariably occur in the . Peptidoglycan, found in some bacterial cell walls contains some D-amino acids. At physiological pH, typically around 7, free amino acids exist in a charged form, where the acidic carboxyl group (-COOH) loses a proton (-COO−) and the basic amine group (-NH2) gains a proton (-NH3+). The entire molecule has a net neutral charge and is a zwitterion, with the exception of amino acids with basic or acidic side chains. Aspartic acid, for example, possesses one protonated amine and two deprotonated carboxyl groups, for a net charge of -1 at physiological pH.
Fatty acids and fatty acid derivatives are another group of carboxylic acids that play a significant role in biology. These contain long hydrocarbon chains and a carboxylic acid group on one end. The cell membrane of nearly all organisms is primarily made up of a phospholipid bilayer, a micelle of hydrophobic fatty acid esters with polar, hydrophilic phosphate "head" groups. Membranes contain additional components, some of which can participate in acid-base reactions.
In humans and many other animals, hydrochloric acid is a part of the gastric acid secreted within the stomach to help hydrolyze proteins and polysaccharides, as well as converting the inactive pro-enzyme, pepsinogen into the enzyme, pepsin. Some organisms produce acids for defense; for example, ants produce formic acid.
Acid-base equilibrium plays a critical role in regulating mammalian breathing. Oxygen gas (O2) drives cellular respiration, the process by which animals release the chemical potential energy stored in food, producing carbon dioxide (CO2) as a byproduct. Oxygen and carbon dioxide are exchanged in the lungs, and the body responds to changing energy demands by adjusting the rate of ventilation. For example, during periods of exertion the body rapidly breaks down stored carbohydrates and fat, releasing CO2 into the blood stream. In aqueous solutions such as blood CO2 exists in equilibrium with carbonic acid and bicarbonate ion.
: CO2 + H2O H2CO3 H+ + HCO3−
It is the decrease in pH that signals the brain to breath faster and deeper, expelling the excess CO2 and resupplying the cells with O2.
(acetylsalicylic acid) is a carboxylic acid.]] Cell membranes are generally impermeable to charged or large, polar molecules because of the lipophilic fatty acyl chains comprising their interior. Many biologically important molecules, including a number of pharmaceutical agents, are organic weak acids which can cross the membrane in their protonated, uncharged form but not in their charged form (i.e. as the conjugate base). For this reason the activity of many drugs can be enhanced or inhibited by the use of antacids or acidic foods. The charged form, however, is often more soluble in blood and cytosol, both aqueous environments. When the extracellular environment is more acidic than the neutral pH within the cell, certain acids will exist in their neutral form and will be membrane soluble, allowing them to cross the phospholipid bilayer. Acids that lose a proton at the intracellular pH will exist in their soluble, charged form and are thus able to diffuse through the cytosol to their target. Ibuprofen, aspirin and penicillin are examples of drugs that are weak acids.
Common acids
Mineral acids (inorganic acids)
Hydrogen halides and their solutions: hydrochloric acid (HCl), hydrobromic acid (HBr), hydroiodic acid (HI)
Halogen oxoacids: hypochlorous acid (HClO), chlorous acid (HClO2), chloric acid (HClO3), perchloric acid (HClO4), and corresponding compounds for bromine and iodine
Sulfuric acid (H2SO4)
Fluorosulfuric acid (HSO3F)
Nitric acid (HNO3)
Phosphoric acid (H3PO4)
Fluoroantimonic acid (HSbF6)
Fluoroboric acid (HBF4)
Hexafluorophosphoric acid (HPF6)
Chromic acid (H2CrO4)
Boric acid (H3BO3)
Sulfonic acids
Methanesulfonic acid (or mesylic acid, CH3SO3H)
Ethanesulfonic acid (or esylic acid, CH3CH2SO3H)
Benzenesulfonic acid (or besylic acid, C6H5SO3H)
p-Toluenesulfonic acid (or tosylic acid, CH3C6H4SO3H)
Trifluoromethanesulfonic acid (or triflic acid, CF3SO3H)
Polystyrene sulfonic acid (sulfonated polystyrene, [CH2CH(C6H4)SO3H]n)
Carboxylic acids
Acetic acid (CH3COOH)
Citric acid (C6H8O7)
Formic acid (HCOOH)
Gluconic acid HOCH2-(CHOH)4-COOH
Lactic acid (CH3-CHOH-COOH)
Oxalic acid (HOOC-COOH)
Tartaric acid (HOOC-CHOH-CHOH-COOH)
Vinylogous carboxylic acids
Ascorbic acid
Meldrum's acid
Nucleic acids
Deoxyribonucleic acid (DNA)
Ribonucleic acid (RNA)
See also
;Chemistry
Acid value
Acid salt
Base
Basic salt
Binary acid
Hard and soft acids and bases (HSAB theory)
Vitriol
Acid-base extraction
;Environment
Acid rain
Ocean acidification
Acid sulfate soil
References
Listing of strengths of common acids and bases
Zumdahl, Chemistry, 4th Edition.
Ebbing, D.D., & Gammon, S. D. (2005). General chemistry (8th ed.). Boston, MA: Houghton Mifflin. ISBN 0-618-51177-6
Pavia, D.L., Lampman, G.M., & Kriz, G.S. (2004). Organic chemistry volume 1: Organic chemistry 351. Mason, OH: Cenage Learning. ISBN 9780759342724
External links
Science Aid: Acids and Bases Information for High School students
Curtipot: Acid-Base equilibria diagrams, pH calculation and titration curves simulation and analysis - freeware
A summary of the Properties of Acids for the beginning chemistry student
The UN ECE Convention on Long-Range Transboundary Air Pollution
Chem 106 - Acidity Concepts
Category:Acid-base chemistry
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