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Acid
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Acid
An acid (from the Latin acidus/acēre meaning sour[1] ) 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
Zinc, a typical metal, reacting with hydrochloric
three examples show, acids can be solutions, liquids, or solids. Gases
acid, a typical acid
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, H3O+, when dissolved in water. This definition
stems from the equilibrium dissociation of water into hydronium and hydroxide (OH−) ions:[2]
H2O(l) + H2O(l)
H3O+(aq) + OH−(aq)
In pure water the majority of molecules exist as H2O, 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, H3O+.
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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.[2] Brønsted-Lowry acid-base theory has several advantages over Arrhenius theory.
Consider the following reactions of acetic acid (CH3COOH), the organic acid that gives vinegar its characteristic
taste:
Both theories easily describe the first reaction: CH3COOH acts as an Arrhenius acid because it acts as a source of
H3O+ when dissolved in water, and it acts as a Brønsted acid by donating a proton to water. In the second example
CH3COOH undergoes the same transformation, in this case donating a proton to ammonia (NH3), but cannot be
described using the Arrhenius definition of an acid because the reaction does not produce hydronium.
Brønsted-Lowry theory can also be used to describe molecular compounds, whereas Arrhenius acids must be ionic
compounds. Hydrogen chloride (HCl) and ammonia combine under several different conditions to form ammonium
chloride, NH4Cl. In aqueous solution HCl behaves as hydrochloric acid and exists as hydronium and chloride ions.
The following reactions illustrate the limitations of Arrhenius's definition:
1. H3O+(aq) + Cl−(aq) + NH3 → Cl−(aq) + NH4+(aq)
2. HCl(benzene) + NH3(benzene) → NH4Cl(s)
3. HCl(g) + NH3(g) → NH4Cl(s)
As with the acetic acid reactions, both definitions work for the first example, where water is the solvent and
hydronium ion is formed. The next two reactions do not involve the formation of ions but are still proton transfer
reactions. In the second reaction hydrogen chloride and ammonia (dissolved in benzene) react to form solid
ammonium chloride in a benzene solvent and in the third gaseous HCl and NH3 combine to form the solid.
Lewis acids
A third concept was proposed in 1923 by Gilbert N. Lewis which includes reactions with acid-base characteristics
that do not involve a proton transfer. A Lewis acid is a species that accepts a pair of electrons from another species;
in other words, it is an electron pair acceptor.[2] Brønsted acid-base reactions are proton transfer reactions while
Lewis acid-base reactions are electron pair transfers. All Brønsted acids are also Lewis acids, but not all Lewis acids
are Brønsted acids. Contrast the following reactions which could be described in terms of acid-base chemistry.
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In the first reaction a fluoride ion, F−, gives up an electron pair to boron trifluoride to form the product
tetrafluoroborate. Fluoride "loses" a pair of valence electrons because the electrons shared in the B—F bond are
located in the region of space between the two atomic nuclei and are therefore more distant from the fluoride nucleus
than they are in the lone fluoride ion. BF3 is a Lewis acid because it accepts the electron pair from fluoride. This
reaction cannot be described in terms of Brønsted theory because there is no proton transfer. The second reaction can
be described using either theory. A proton is transferred from an unspecified Brønsted acid to ammonia, a Brønsted
base; alternatively, ammonia acts as a Lewis base and transfers a lone pair of electrons to form a bond with a
hydrogen ion. The species that gains the electron pair is the Lewis acid; for example, the oxygen atom in H3O+ gains
a pair of electrons when one of the H—O bonds is broken and the electrons shared in the bond become localized on
oxygen. Depending on the context, a Lewis acid may also be described as an oxidizer or an electrophile.
The Brønsted-Lowry definition is the most widely used definition; unless otherwise specified acid-base reactions are
assumed to involve the transfer of a proton (H+) from an acid to a base.
Dissociation and equilibrium
Reactions of acids are often generalized in the form HA
H+ + A−, where HA represents the acid and A− is the
conjugate base. Acid-base conjugate pairs differ by one proton, and can be interconverted by the addition or removal
of a proton (protonation and deprotonation, respectively). Note that the acid can be the charged species and the
conjugate base can be neutral in which case the generalized reaction scheme could be written as HA+
H+ + A. In
solution there exists an equilibrium between the acid and its conjugate base. The equilibrium constant K is an
expression of the equilibrium concentrations of the molecules or the ions in solution. Brackets indicate
concentration, such that [H2O] means the concentration of H2O. The acid dissociation constant Ka is generally used
in the context of acid-base reactions. The numerical value of Ka is equal to the concentration of the products divided
by the concentration of the reactants, where the reactant is the acid (HA) and the products are the conjugate base and
H+.
The stronger of two acids will have a higher Ka than the weaker acid; the ratio of hydrogen ions to acid will be higher
for the stronger acid as the stronger acid has a greater tendency to lose its proton. Because the range of possible
values for Ka spans many orders of magnitude, a more manageable constant, pKa is more frequently used, where pKa
= -log10 Ka. Stronger acids have a smaller pKa than weaker acids. Experimentally determined pKa at 25°C in aqueous
solution are often quoted in textbooks and reference material.
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Nomenclature
In the classical naming system, 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. In the IUPAC naming system, "aqueous" is simply added
to the name of the ionic compound. Thus, for hydrogen chloride, the IUPAC name would be aqueous hydrogen
chloride. The prefix "hydro-" is added only if the acid is made up of just hydrogen and one other element.
Classical naming system:
Anion prefix
Anion suffix
Acid prefix
Acid suffix
per
ate
per
ic acid
perchloric acid (HClO4)
ate
ic acid
chloric acid (HClO3)
ite
ous acid
chlorous acid (HClO2)
hypo
Example
ite
hypo
ous acid
hypochlorous acid (HClO)
ide
hydro
ic acid
hydrochloric acid (HCl)
Acid strength
The strength of an acid refers to its ability or tendency to lose a proton. A strong acid is one that completely
dissociates in water; in other words, one mole of a strong acid HA dissolves in water yielding one mole of H+ and
one mole of the conjugate base, A−, and none of the protonated acid HA. In contrast a weak acid only partially
dissociates and at equilibrium both the acid and the conjugate base are in solution. Examples of strong acids are
hydrochloric acid (HCl), hydroiodic acid (HI), hydrobromic acid (HBr), perchloric acid (HClO4), nitric acid (HNO3)
and sulfuric acid (H2SO4). In water each of these essentially ionizes 100%. The stronger an acid is, the more easily it
loses a proton, H+. Two key factors that contribute to the ease of deprotonation are the polarity of the H—A bond
and the size of atom A, which determines the strength of the H—A bond. Acid strengths are also often discussed in
terms of the stability of the conjugate base.
Stronger acids have a larger Ka and a more negative pKa than weaker acids.
Sulfonic acids, which are organic oxyacids, are a class of strong acids. A common example is toluenesulfonic acid
(tosylic acid). Unlike sulfuric acid itself, sulfonic acids can be solids. In fact, polystyrene functionalized into
polystyrene sulfonate is a solid strongly acidic plastic that is filterable.
Superacids are acids stronger than 100% sulfuric acid. Examples of superacids are fluoroantimonic acid, magic acid
and perchloric acid. Superacids can permanently protonate water to give ionic, crystalline hydronium "salts". They
can also quantitatively stabilize carbocations.
Polarity and the inductive effect
Polarity refers to the distribution of electrons in a bond, the region of space between two atomic nuclei where a pair
of electrons is shared. When two atoms have roughly the same electronegativity (ability to attract electrons) the
electrons are shared evenly and spend equal time on either end of the bond. When there is a significant difference in
electronegativities of two bonded atoms, the electrons spend more time near the nucleus of the more electronegative
element and an electrical dipole, or separation of charges, occurs, such that there is a partial negative charge
localized on the electronegative element and a partial positive charge on the electropositive element. Hydrogen is an
electropositive element and accumulates a slightly positive charge when it is bonded to an electronegative element
such as oxygen or bromine. As the electron density on hydrogen decreases it is more easily abstracted, in other
words, it is more acidic. Moving from left to right across a row on the periodic table elements become more
electronegative (excluding the noble gases), and the strength of the binary acid formed by the element increases
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accordingly:
Formula
Name
[3]
pKa
HF
hydrofluoric acid 3.17
H2O
water
15.7
NH3
ammonia
38
CH4
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.
Structure
Name
[4]
pKa
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 pKa of 2-bromobutanoic acid is slightly greater than the pKa of 2-chlorobutanoic
acid.
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Perchloric 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.[5]
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.[4] [6]
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, H2S, 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.
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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)
HA−(aq) + H2O(l)
H3O+(aq) + HA−(aq) Ka1
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
HA2−(aq) + H2O(l)
H3O+(aq) + A3−(aq) Ka3
H2A−(aq) + H2O(l)
H3O+(aq) + HA2−(aq) Ka2
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:
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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
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.
Hydrochloric acid (in beaker) reacting with
ammonia fumes to produce ammonium chloride
(white smoke).
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).
Solutions of weak acids and salts of their conjugate bases form buffer solutions.
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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.[7]
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.
Basic structure of an amino acid.
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
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elements such as sulfur, oxygen or nitrogen. With the exception of glycine, naturally occurring amino acids are
chiral and almost invariably occur in the L-configuration. 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.
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
Aspirin (acetylsalicylic acid) is a carboxylic acid.
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.
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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)
References
[1] Merriam-Webster's Online Dictionary: acid (http:/ / www. merriam-webster. com/ dictionary/ acid)
[2] Ebbing, D.D., & Gammon, S. D. (2005). General chemistry (8th ed.). Boston, MA: Houghton Mifflin. ISBN 0-618-51177-6
[3] pKa's of Inorganic and Oxo-Acids (http:/ / www2. lsdiv. harvard. edu/ labs/ evans/ pdf/ evans_pKa_table. pdf)
[4] Section 8: Electrolytes, Electromotive forces and Chemical Equilibrium (http:/ / www. scribd. com/ doc/ 6792576/ 638478)
Acid
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[5] pKa values for HClOn from Housecroft, C. E.; Sharpe, A. G. (2004). Inorganic Chemistry (2nd ed.). Prentice Hall. ISBN 978-0130399137.
[6] pKa Data Compiled by R. Williams (http:/ / research. chem. psu. edu/ brpgroup/ pKa_compilation. pdf)
[7] Voet, Judith G.; Voet, Donald (2004). Biochemistry. New York: J. Wiley & Sons. pp. 496–500. ISBN 9780471193500.
• Listing of strengths of common acids and bases (http://www.csudh.edu/oliver/chemdata/data-ka.htm)
• 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 (http://scienceaid.co.uk/chemistry/physical/acidbases.html) Information for
High School students
• Curtipot (http://www2.iq.usp.br/docente/gutz/Curtipot_.html): 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 (http://canadaconnects.ca/chemistry/
10081/)
• The UN ECE Convention on Long-Range Transboundary Air Pollution (http://www.unece.org/env/lrtap/)
• Chem 106 - Acidity Concepts (http://isites.harvard.edu/fs/docs/icb.topic776365.files/lecture 17.pdf)
Article Sources and Contributors
Article Sources and Contributors
Acid Source: http://en.wikipedia.org/w/index.php?oldid=438233241 Contributors: -Midorihana-, 0612, 1nt2, 217.35.151.xxx, 24fan24, 2D, 64.26.98.xxx, 66.156.135.xxx, AJR, AVM, Adashiel,
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Image Sources, Licenses and Contributors
File:Zn reaction with HCl.JPG Source: http://en.wikipedia.org/w/index.php?title=File:Zn_reaction_with_HCl.JPG License: Public Domain Contributors: Chemicalinterest
Image:Equilibrium.svg Source: http://en.wikipedia.org/w/index.php?title=File:Equilibrium.svg License: Public Domain Contributors: L'Aquatique
File:Acid-base.png Source: http://en.wikipedia.org/w/index.php?title=File:Acid-base.png License: Public Domain Contributors: Cedric M. Johnson
File:LewisAcid.png Source: http://en.wikipedia.org/w/index.php?title=File:LewisAcid.png License: Public Domain Contributors: Myceteae
File:butanoic.png Source: http://en.wikipedia.org/w/index.php?title=File:Butanoic.png License: Public Domain Contributors: Myceteae
File:4chlorobutanoic.png Source: http://en.wikipedia.org/w/index.php?title=File:4chlorobutanoic.png License: Public Domain Contributors: Myceteae
File:3chlorobutanoic.png Source: http://en.wikipedia.org/w/index.php?title=File:3chlorobutanoic.png License: Public Domain Contributors: Myceteae
File:2bromobutanoic.png Source: http://en.wikipedia.org/w/index.php?title=File:2bromobutanoic.png License: Public Domain Contributors: Myceteae
File:2chlorobutanoic.png Source: http://en.wikipedia.org/w/index.php?title=File:2chlorobutanoic.png License: Public Domain Contributors: Myceteae
Image:Perchloricacid.png Source: http://en.wikipedia.org/w/index.php?title=File:Perchloricacid.png License: Public Domain Contributors: Cedric M. Johnson
File:Reduction.png Source: http://en.wikipedia.org/w/index.php?title=File:Reduction.png License: Public Domain Contributors: Cedric M. Johnson
Image:Hydrochloric acid ammonia.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Hydrochloric_acid_ammonia.jpg License: Public Domain Contributors: User:Walkerma
Image:Aminoacid.png Source: http://en.wikipedia.org/w/index.php?title=File:Aminoacid.png License: Public Domain Contributors: Myceteae
Image:Aspirin-skeletal.svg Source: http://en.wikipedia.org/w/index.php?title=File:Aspirin-skeletal.svg License: Public Domain Contributors: Originally User:Benjah-bmm27
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