An acid is a molecule or ion capable of donating a hydron (proton or
hydrogen ion H+), or, alternatively, capable of forming a covalent
bond with an electron pair (a Lewis acid).
The first category of acids is the proton donors or Brønsted acids.
In the special case of aqueous solutions, proton donors form the
hydronium ion H3O+ and are known as Arrhenius acids. Brønsted and
Lowry generalized the Arrhenius theory to include non-aqueous
solvents. A Brønsted or Arrhenius acid usually contains a hydrogen
atom bonded to a chemical structure that is still energetically
favorable after loss of H+.
Aqueous Arrhenius acids have characteristic properties which provide a
practical description of an acid. Acids form aqueous solutions with
a sour taste, can turn blue litmus red, and react with bases and
certain metals (like calcium) to form salts. The word acid is derived
Latin acidus/acēre meaning sour. An aqueous solution of
an acid has a pH less than 7 and is colloquially also referred to as
'acid' (as in 'dissolved in acid'), while the strict definition refers
only to the solute. A lower pH means a higher acidity, and thus a
higher concentration of positive hydrogen ions in the solution.
Chemicals or substances having the property of an acid are said to be
Common aqueous acids include hydrochloric acid (a solution of hydrogen
chloride which is found in gastric acid in the stomach and activates
digestive enzymes), acetic acid (vinegar is a dilute aqueous solution
of this liquid), sulfuric acid (used in car batteries), and citric
acid (found in citrus fruits). As these examples show, acids (in the
colloquial sense) can be solutions or pure substances, and can be
derived from acids (in the strict sense) that are solids, liquids,
or gases. Strong acids and some concentrated weak acids are corrosive,
but there are exceptions such as carboranes and boric acid.
The second category of acids are Lewis acids, which form a covalent
bond with an electron pair. An example is boron trifluoride (BF3),
whose boron atom has a vacant orbital which can form a covalent bond
by sharing a lone pair of electrons on an atom in a base, for example
the nitrogen atom in ammonia (NH3). Lewis considered this as a
generalization of the Brønsted definition, so that an acid is a
chemical species that accepts electron pairs either directly or by
releasing protons (H+) into the solution, which then accept electron
pairs. However, hydrogen chloride, acetic acid, and most other
Brønsted-Lowry acids cannot form a covalent bond with an electron
pair and are therefore not Lewis acids. Conversely, many Lewis
acids are not Arrhenius or Brønsted-Lowry acids. In modern
terminology, an acid is implicitly a Brønsted acid and not a Lewis
acid, since chemists almost always refer to a Lewis acid explicitly as
a Lewis acid.
1 Definitions and concepts
1.1 Arrhenius acids
1.2 Brønsted–Lowry acids
1.3 Lewis acids
2 Dissociation and equilibrium
5 Chemical characteristics
5.1 Monoprotic acids
5.2 Polyprotic acids
5.4 Weak acid–weak base equilibrium
6 Applications of acids
7 Biological occurrence
8 Common acids
8.1 Mineral acids (inorganic acids)
8.2 Sulfonic acids
8.3 Carboxylic acids
8.4 Halogenated carboxylic acids
Vinylogous carboxylic acids
8.6 Nucleic acids
10 External links
Definitions and concepts
Main article: Acid–base reaction
Modern definitions are concerned with the fundamental chemical
reactions common to all acids.
Most acids encountered in everyday life are aqueous solutions, or can
be dissolved in water, so the Arrhenius and Brønsted-Lowry
definitions are the most relevant.
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.
Hydronium ions are acids according to all three definitions.
Interestingly, although alcohols and amines can be Brønsted-Lowry
acids, they can also function as Lewis bases due to the lone pairs of
electrons on their oxygen and nitrogen atoms.
The Swedish chemist
Svante Arrhenius attributed the properties of
acidity to hydrogen ions (H+) or protons in 1884. An Arrhenius acid is
a substance that, when added to water, increases the concentration of
H+ ions in the 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+. Thus, an Arrhenius acid can also be
described as a substance that increases the concentration of hydronium
ions when added to water. Examples include molecular substances such
as HCl and acetic acid.
An Arrhenius base, on the other hand, is a substance which increases
the concentration of hydroxide (OH−) ions when dissolved in water.
This decreases the concentration of hydronium because the ions react
to form H2O molecules:
(aq) + OH−
(aq) ⇌ H2O(l) + H2O(l)
Due to this equilibrium, any increase in the concentration of
hydronium is accompanied by a decrease in the concentration of
hydroxide. Thus, an Arrhenius acid could also be said to be one that
decreases hydroxide concentration, while an Arrhenius base increases
In an acidic solution, 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, acidic solutions
thus have a pH of less than 7.
Main article: Brønsted–Lowry acid–base theory
Acetic acid, a weak acid, donates a proton (hydrogen ion, highlighted
in green) to water in an equilibrium reaction to give the acetate ion
and the hydronium ion. Red: oxygen, black: carbon, white: hydrogen.
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
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. 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:
3COOH + H
2O ⇌ CH
3COO− + H
3COOH + NH
3 ⇌ CH
3COO− + NH+
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 does not relate to
the Arrhenius definition of an acid because the reaction does not
produce hydronium. Nevertheless, CH3COOH is both an Arrhenius and a
Brønsted-Lowry theory can be used to describe reactions of molecular
compounds in nonaqueous solution or the gas phase. 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
(aq) + Cl−
(aq) + NH3 → Cl−
(aq) + NH+
4(aq) + H2O
HCl(benzene) + NH3(benzene) → NH4Cl(s)
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 by the
HCl solute. 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.
Main article: Lewis acids and bases
A third, only marginally related 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. Brønsted acid-base
reactions are proton transfer reactions while Lewis acid-base
reactions are electron pair transfers. Many Lewis acids are not
Brønsted-Lowry acids. Contrast how the following reactions are
described in terms of acid-base chemistry:
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. Organic Brønsted
acids, such as acetic, citric, or oxalic acid, are not Lewis acids.
They dissociate in water to produce a Lewis acid, H+, but at the same
time also yield an equal amount of a
Lewis base (acetate, citrate, or
oxalate, respectively, for the acids mentioned). Few, if any, of the
acids discussed in the following are Lewis acids.
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. This
reaction is referred to as protolysis. The protonated form (HA) of an
acid is also sometimes referred to as the free acid.
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 product of the concentrations 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+.
displaystyle K_ a = frac ce [H+] [A^-] ce [HA]
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.
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,
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
Classical naming system:
perchloric acid (HClO4)
chloric acid (HClO3)
chlorous acid (HClO2)
hypochlorous acid (HClO)
hydrochloric acid (HCl)
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
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
While Ka measures the strength of an acid compound, the strength of an
aqueous acid solution is measured by pH, which is an indication of the
concentration of hydronium in the solution. The pH of a simple
solution of an acid compound in water is determined by the dilution of
the compound and the compound's Ka.
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−
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).
Acid dissociation constant
Acid dissociation constant § Monoprotic 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) + H2O(l) ⇌ H3O+
(aq) + A2−
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 (HSO−
4), for which Ka1 is very large; then it can donate a second proton to
form the sulfate anion (SO2−
4), 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 (HCO−
3) and lose a second to form carbonate anion (CO2−
3). 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) + H2O(l) ⇌ H3O+
(aq) + HA2−
(aq) + H2O(l) ⇌ H3O+
(aq) + A3−
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 H2PO−
4, then HPO2−
4, and finally PO3−
4, the orthophosphate ion, usually just called phosphate. Even though
the positions of the three protons on the original phosphoric acid
molecule are equivalent, the successive Ka values differ since it is
energetically less favorable to lose a proton if the conjugate base is
more negatively charged. An organic example of a triprotic acid is
citric acid, which can successively lose three protons to finally form
the citrate ion.
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:
displaystyle begin aligned alpha _ ce H2A &= frac ce
[H+]^ 2 ce [H+]^ 2 +[ ce H+ ]K_ 1 +K_ 1 K_ 2 = frac ce
[H2A] ce [H2A] + [HA^ - ] + [A^ 2- ] \alpha _ ce HA^ -
&= frac [ ce H+ ]K_ 1 ce [H+]^ 2 +[ ce H+ ]K_ 1 +K_ 1
K_ 2 = frac ce [HA^ - ] ce [H2A] + [HA^ - ] + [A^ 2- ]
\alpha _ ce A^ 2- &= frac K_ 1 K_ 2 ce [H+]^ 2 +[ ce H+
]K_ 1 +K_ 1 K_ 2 = frac ce [A^ 2- ] ce [H2A] + [HA^ - ] + [A^
2- ] end aligned
A plot of these fractional concentrations against pH, for given K1 and
K2, is known as a Bjerrum plot. A pattern is observed in the above
equations and can be expanded to the general n -protic acid that has
been deprotonated i -times:
displaystyle alpha _ ce H _ n-i A^ i- = [ ce H+ ]^ n-i
displaystyle prod _ j=0 ^ i K_ j over displaystyle sum _ i=0 ^ n
Big [ [ ce H+ ]^ n-i displaystyle prod _ j=0 ^ i K_ j Big ]
where K0 = 1 and the other K-terms are the dissociation constants for
Acid dissociation constant
Acid dissociation constant § Polyprotic acids
Hydrochloric acid (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
Weak acid–weak base equilibrium
Main article: Henderson–Hasselbalch equation
In order for a protonated acid to lose a proton, the pH of the system
must rise above the pKa of the 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.
Solutions of weak acids and salts of their conjugate bases form buffer
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
Carbonic acid is an important part of some cola drinks and
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
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 for the human body
and is present in such foods as amla (Indian gooseberry), 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 important roles in the human body. The hydrochloric acid
present in the stomach aids digestion by breaking down large and
complex food molecules. Amino acids are required for synthesis of
proteins required for growth and repair of body tissues. Fatty acids
are also required for growth and repair of body tissues. Nucleic acids
are important for the manufacturing of
RNA and transmitting of
traits to offspring through genes.
Carbonic acid is important for
maintenance of pH equilibrium in the body.
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. Some acids, such as sulfuric,
phosphoric, and hydrochloric acids, also effect dehydration and
condensation reactions. In biochemistry, many enzymes employ acid
Basic structure of an amino acid.
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.
the chemical blueprint for the synthesis of proteins which are made up
of amino acid subunits. Cell membranes contain fatty acid esters such
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
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 (-NH+
3). 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
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
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+ + HCO−
It is the decrease in pH that signals the brain to breathe faster and
deeper, expelling the excess CO2 and resupplying the cells with O2.
Aspirin (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.
Mineral acids (inorganic acids)
Hydrogen halides and their solutions: hydrofluoric acid (HF),
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
analogs for bromine and iodine
Hypofluorous acid (HFO), the only known oxoacid for fluorine.
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)
A sulfonic acid has the general formula RS(=O)2–OH, where R is an
Methanesulfonic acid (or mesylic acid, CH3SO3H)
Ethanesulfonic acid (or esylic acid, CH3CH2SO3H)
Benzenesulfonic acid (or besylic acid, C6H5SO3H)
Toluenesulfonic acid (or tosylic acid, CH3C6H4SO3H)
Trifluoromethanesulfonic acid (or triflic acid, CF3SO3H)
Polystyrene sulfonic acid (sulfonated polystyrene, [CH2CH(C6H4)SO3H]n)
A carboxylic acid has the general formula R-C(O)OH, where R is an
organic radical. The carboxyl group -C(O)OH contains a carbonyl group,
C=O, and a hydroxyl group, O-H.
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)
Halogenated carboxylic acids
Halogenation at alpha position increases acid strength, so that the
following acids are all stronger than acetic acid.
Vinylogous carboxylic acids
Normal carboxylic acids are the direct union of a carbonyl group and a
hydroxyl group. In vinylogous carboxylic acids, a carbon-carbon double
bond separates the carbonyl and hydroxyl groups.
Deoxyribonucleic acid (DNA)
Ribonucleic acid (RNA)
^ a b c
IUPAC Gold Book - acid
^ Petrucci R.H., Harwood, R.S. and Herring, F.G. General Chemistry
(8th ed., Prentice-Hall 2002) p.146 ISBN 0-13-014329-4
^ Merriam-Webster's Online Dictionary: acid
^ a b c d Oxtoby, D. W; Gillis, H.P., Butler, L. J. (2015).Principles
of Modern Chemistry, Brooks Cole. p. 617. ISBN 978-1305079113
^ a b c Ebbing, D.D., & Gammon, S. D. (2005). General chemistry
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curves simulation and analysis – freeware