An acid is a chemical substance whose aqueous
solutions are characterized by a sour taste,
the ability to turn blue litmus red, and the
ability to react with bases and certain metals
to form salts. Aqueous solutions of acids
have a pH of less than 7. A lower pH means
a higher acidity, and thus a higher concentration
of hydrogen ions in the solution. Chemicals
or substances having the property of an acid
are said to be acidic.
Common examples of acids include hydrochloric
acid, acetic acid, sulfuric acid, and tartaric
acid. As these examples show, acids can be
solutions or pure substances, and can be derived
from solids, liquids, or gases. 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 defines acids as substances
which increase the concentration of hydrogen
ions, or more accurately, hydronium ions,
when dissolved in water. The Brønsted-Lowry
definition is an expansion: an acid is a substance
which can act as a proton donor. By this 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. A Lewis acid is a substance that
can accept a pair of electrons to form a covalent
bond. Examples of Lewis acids include all
metal cations, and electron-deficient molecules
such as boron trifluoride and aluminium trichloride.
Definitions and concepts
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 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.
Arrhenius acids
The Swedish chemist Svante Arrhenius attributed
the properties of acidity to hydrogen ions
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. This
definition stems from the equilibrium dissociation
of water into hydronium and hydroxide ions:
H2O(l) + H2O(l) H3O+(aq) + OH−(aq)
In pure water the majority of molecules are
H2O, but the molecules are constantly dissociating
and re-associating, and at any time a small
number of the molecules are hydronium and
an equal number are hydroxide. Because the
numbers are equal, pure water is neutral.
An Arrhenius base, on the other hand, is a
substance which increases the concentration
of hydroxide ions when dissolved in water,
hence decreasing the concentration of hydronium.
The constant association and disassociation
of H2O molecules forms an equilibrium in which
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, with an Arrhenius base increasing
it.
The reason that 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.
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 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, the organic acid that gives vinegar
its characteristic taste:
CH
3COOH + H
2O CH
3COO− + H
3O+
CH
3COOH + NH
3 CH
3COO− + NH+
4
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,
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
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:
H3O+(aq) + Cl−(aq) + NH3 → Cl−(aq) +
NH4+(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 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.
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.
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.
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. 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
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.
Nomenclature
In the classical naming system, acids are
named according to their anions. That ionic
suffix is dropped and replaced with a new
suffix, 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:
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, hydroiodic acid, hydrobromic
acid, perchloric acid, nitric acid and sulfuric
acid. 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. 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.
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.
Chemical characteristics
Monoprotic acids
Monoprotic acids are those acids that are
able to donate one proton per molecule during
the process of dissociation as shown below:
HA(aq) + H2O(l) H3O+(aq) + A−(aq)         Ka
Common examples of monoprotic acids in mineral
acids include hydrochloric acid and nitric
acid. 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, acetic
acid and benzoic acid.
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 and
triprotic acid.
A diprotic acid 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 can donate one
proton to form the bisulfate anion, for which
Ka1 is very large; then it can donate a second
proton to form the sulfate anion, 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 can lose one proton
to form bicarbonate anion and lose a second
to form carbonate anion. Both Ka values are
small, but Ka1 > Ka2 .
A triprotic acid 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, 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, 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 or the concentrations of the acid with
all its conjugate bases:
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:
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.
Weak acid–weak base equilibrium
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. In lower-pH
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 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 is an essential vitamin for
the human body and is present in such foods
as amla, lemon, citrus fruits, and guava.
Certain acids are used as drugs. Acetylsalicylic
acid 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 in 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 DNA and RNA and transmitting
of traits 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 which
is covalently bonded to a carboxyl group,
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 loses a proton and the basic amine group
gains a proton. 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 drives cellular respiration, the process
by which animals release the chemical potential
energy stored in food, producing carbon dioxide
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 breathe 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. 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
Hydrogen halides and their solutions: hydrofluoric
acid, hydrochloric acid, hydrobromic acid,
hydroiodic acid
Halogen oxoacids: hypochlorous acid, chlorous
acid, chloric acid, perchloric acid, and corresponding
compounds for bromine and iodine
Sulfuric acid
Fluorosulfuric acid
Nitric acid
Phosphoric acid
Fluoroantimonic acid
Fluoroboric acid
Hexafluorophosphoric acid
Chromic acid
Boric acid
Sulfonic acids
Methanesulfonic acid
Ethanesulfonic acid
Benzenesulfonic acid
p-Toluenesulfonic acid
Trifluoromethanesulfonic acid
Polystyrene sulfonic acidSO3H]n)
Carboxylic acids
Acetic acid
Citric acid
Formic acid
Gluconic acid HOCH2-(CHOH)4-COOH
Lactic acid
Oxalic acid
Tartaric acid
Halogenated carboxylic acids
Halogenation at alpha position increases acid
strength, so that the following acids are
all stronger than acetic acid.
Fluoroacetic acid
Trifluoroacetic acid
Chloroacetic acid
Dichloroacetic acid
Trichloroacetic acid
Vinylogous carboxylic acids
Normal carboxylic acids are the direct union
of a carbonyl group and a hydroxy group. In
vinylogous carboxylic acids, a carbon-carbon
double bond separates the carbonyl and hydroxyl
groups.
Ascorbic acid
Nucleic acids
Deoxyribonucleic acid
Ribonucleic acid
References
Listing of strengths of common acids and bases
IUPAC Gold Book - acid
Zumdahl, Chemistry, 4th Edition.
Ebbing, D.D., & Gammon, S. D.. General chemistry.
Boston, MA: Houghton Mifflin. ISBN 0-618-51177-6
Pavia, D.L., Lampman, G.M., & Kriz, G.S..
Organic chemistry volume 1: Organic chemistry
351. Mason, OH: Cenage Learning. ISBN 0-7593-4727-1
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
