ACID
An acid (from the Latin acidus/acēre meaning souris
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. 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, 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)
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:
- H3O+(aq)
+ Cl−(aq) + NH3 → Cl−(aq) +
NH4+(aq)
- 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.
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.
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.
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.
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
|
Example
|
per
|
ate
|
per
|
ic
acid
|
perchloric
acid
(HClO4)
|
ate
|
ic
acid
|
chloric
acid
(HClO3)
|
||
ite
|
ous
acid
|
chlorous
acid
(HClO2)
|
||
hypo
|
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.
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.
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.
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:
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 sodium hydroxide.
Weak acid–weak base equilibrium
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).
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.[3]
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 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.
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
(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)
- 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
Nucleic acids
- Deoxyribonucleic acid (DNA)
- Ribonucleic acid (RNA)
See also
Chemistry
- Acid-base extraction
- Acid value
- Acid salt
- Base
- Basic salt
- Binary acid
- Hard and soft acids and bases (HSAB theory)
- Titration
- Vitriol
Environment
·
A base in chemistry is a substance that
can accept hydrogen ions (protons) or more generally,
donate electron pairs. A soluble base is referred to as an alkali if it contains and releases hydroxide
ions (OH−)
quantitatively. The Brønsted-Lowry
theory defines bases as proton (hydrogen ion) acceptors, while the more
general Lewis theory defines bases as electron pair donors, allowing
other Lewis acids than
protons to be included.[1] The oldest Arrhenius theory
defines bases as hydroxide anions,[2] which is strictly applicable only to alkali. In
water, by altering the autoionization equilibrium, bases
give solutions with a hydrogen ion activity lower
than that of pure water, i.e. a pH higher
than 7.0 at standard conditions. Examples of common bases are sodium hydroxide and ammonia. Metal oxides, hydroxides and especially alkoxides are basic, and
counteranions of weak acids are weak bases.
·
Bases can be thought of as the chemical
opposite of acids. A
reaction between an acid and base is called neutralization. Bases
and acids are seen as opposites because the effect of an acid is to increase
the hydronium ion (H3O+) concentration in water, whereas
bases reduce this concentration. Bases and acids are typically found in aqueous solution forms.
Aqueous solutions of bases react with aqueous solutions of acids to produce water and salts in aqueous solutions in which the salts
separate into their component ions. If the aqueous solution is a saturated
solution with respect to a given salt solute any additional such salt present in
the solution will result in formation of a precipitate of the salt
Definitions
A strong base is a base which hydrolyzes completely, raising the pH
of the solution toward 14. Concentrated bases, like concentrated acids, attack
living tissue and cause serious burns. The reaction of bases upon contact with
skin is different from that of acids. So while either may be quite destructive,
strong acids are called corrosive, and strong bases are
referred to as caustic. Superbases are a class of especially
basic compounds and non-nucleophilic bases are a special class of
strong bases with poor nucleophilicity. Bases may also be weak
bases
such as ammonia, which is used for cleaning. Arrhenius bases are
water-soluble and these solutions always have a pH greater than 7 at standard
conditions. An alkali is a special example of a
base, where in an aqueous environment, hydroxide ions are donated. There are
other more generalized and advanced definitions of acids and bases.
The notion of a base as a concept in chemistry was first
introduced by the French chemist Guillaume François Rouelle in 1754. He noted that
acids, which in those days were mostly volatile liquids (like acetic acid), turned into solid salts
only when combined with specific substances. Rouelle considered that such a
substance serves as a base for the salt, giving the salt a
"concrete or solid form”.[3]
Properties
Some general properties of bases include
- Slimy
or soapy feel on fingers, due to saponification of the lipids in human skin.
- Concentrated
or strong bases are caustic on organic matter and react violently with acidic
substances.
- Aqueous
solutions or molten bases dissociate in ions and conduct electricity.
- The
pH level of a basic solution is higher than 7.
- Bases
are bitter in taste.[4]
Bases and pH
The pH of an aqueous sample (water) is a measure of its acidity. In pure water, about one
in ten million molecules dissociate into hydronium ions and hydroxide ions
according to the following equation:
2H2O(l) ⇌ H3O+(aq)
+ OH−(aq)
The concentration, measured in molarity (M or moles per litre), of the ions is
indicated as [H3O+] and [OH−]; their product
is the dissociation constant which has the value of 10−14
M2. The pH is defined as −log [H3O+];
thus, pure water has a pH of 7. (These numbers are correct at 23 °C and
are slightly different at other temperatures.)
A base accepts protons from hydronium ions, or donates
hydroxide ions to the solution. Both actions will lower the concentration of
hydronium ions, and thus raise the pH. By contrast, an acid donates protons to
water or accepts OH−, thus increasing the concentration of hydronium
and lowering the pH.
For example, if 0.1 mol (4 g) of sodium
hydroxide (NaOH) are dissolved in water to make 1 litre of solution, the
concentration of hydroxide ions becomes [OH−] = 0.1 mol/L.
As the ionic product remains a constant value, [H+] = 1×10−14/[OH−]
= 1×10−13 mol/L, and pH = −log 10−13 = 13.
The base dissociation constant, Kb, is
a measure of basicity. It is related to the acid dissociation constant, Ka,
by the simple relationship pKa + pKb = 14,
where pKb and pKa are the negative
logarithms of Kb and Ka, respectively.
Alkalinity is a measure of the
ability of a solution to neutralize acids to the equivalence points of
carbonates or bicarbonates.
Neutralization of acids
When dissolved in water, the strong base sodium hydroxide
ionizes into hydroxide and sodium ions:
NaOH → Na+ + OH−
HCl + H2O → H3O+
+ Cl−
When the two solutions are mixed, the H3O+
and OH− ions combine to form water molecules:
H3O+
+ OH− → 2 H2O
If equal quantities of NaOH and HCl are dissolved, the
base and the acid neutralize exactly, leaving only NaCl, effectively table
salt,
in solution.
Alkalinity of
non-hydroxides
Bases are generally compounds that can neutralize an
amount of acids. Both sodium
carbonate
and ammonia are bases, although neither of these
substances contains OH− groups. Both compounds accept H+
when dissolved in protic
solvents
such as water:
Na2CO3
+ H2O → 2 Na+ + HCO3- + OH-
NH3 + H2O
→ NH4+ + OH-
From this, a pH, or acidity, can be calculated for aqueous
solutions of bases. Bases also directly act as electron-pair donors themselves:
CO32-
+ H+ → HCO3-
NH3 + H+
→ NH4+
Carbon can act as a base as well
as nitrogen and oxygen. This occurs typically in
compounds such as butyl
lithium, alkoxides, and metal amides such as sodium amide. Bases of carbon, nitrogen
and oxygen without resonance stabilization are usually
very strong, or superbases, which cannot exist in a
water solution due to the acidity of water. Resonance stabilization, however,
enables weaker bases such as carboxylates; for example, sodium acetate is a weak
base.
Strong bases
A strong base is a basic chemical compound that is able
to deprotonate very weak acids in an acid-base reaction. Common examples of
strong bases are the hydroxides of alkali metals and alkaline earth metals like
NaOH and Ca(OH)2. Very strong bases are even able to deprotonate
very weakly acidic C–H groups in the absence of water. Here is a list of
several strong bases:
- Potassium hydroxide (KOH)
- Barium hydroxide (Ba(OH)2)
- Caesium hydroxide (CsOH)
- Sodium hydroxide (NaOH)
- Strontium hydroxide (Sr(OH)2)
- Calcium hydroxide (Ca(OH)2)
- Magnesium hydroxide (Mg(OH)2)
- Lithium hydroxide (LiOH)
- Rubidium hydroxide (RbOH)
The cations of these strong bases appear in the first and
second groups of the periodic table (alkali and earth alkali metals).
Acids with a pKa of more than about 13
are considered very weak, and their conjugate
bases
are strong bases.
Group 1 salts of carbanions, amides, and hydrides tend to
be even stronger bases due to the extreme weakness of their conjugate acids,
which are stable hydrocarbons, amines, and dihydrogen. Usually these bases are
created by adding pure alkali metals such as sodium into the conjugate acid.
They are called superbases and it is not possible to
keep them in water solution, due to the fact they are stronger bases than the
hydroxide ion and as such they will deprotonate the conjugate acid water. For
example, the ethoxide ion (conjugate base of ethanol) in the presence of water
will undergo this reaction.
CH3CH2O−
+ H2O → CH3CH2OH + OH−
Here are some superbases:
- Butyl
lithium (n-BuLi)
- Lithium diisopropylamide (LDA) (C6H14LiN)
- Lithium
diethylamide (LDEA)
- Sodium
amide (NaNH2)
- Sodium
hydride (NaH)
- Lithium bis(trimethylsilyl)amide (((CH3)3Si)2NLi)
Bases as catalysts
Basic substances can be used as insoluble heterogeneous catalysts for chemical reactions. Some examples are metal
oxides such as magnesium
oxide, calcium oxide, and barium oxide as well as potassium fluoride on alumina and some zeolites. Many transition metals make good catalysts, many
of which form basic substances. Basic catalysts have been used for hydrogenations, the migration of double bonds, in the Meerwein-Ponndorf-Verley reduction, the Michael reaction, and many other reactions.
SALT
In chemistry, salts are ionic compounds that
result from the neutralization reaction
of an acid and a base. They are
composed of cations (positively charged ions) and anions (negative ions) so that the product
is electrically neutral (without a net
charge). These component ions can be inorganic such as
chloride (Cl−), as well as organic such as acetate (CH3COO−)
and monatomic ions such as
fluoride (F−), as well as polyatomic ions such as sulfate (SO42−).
There are several varieties of salts. Salts
that hydrolyze to produce hydroxide ions when dissolved
in water are basic salts and salts
that hydrolyze to produce hydronium ions in water are acid salts. Neutral salts
are those that are neither acid nor basic salts. Zwitterions contain an anionic
center and a cationic center in the same molecule but are not
considered to be salts. Examples include amino acids, many metabolites, peptides, and proteins.
Molten salts and solutions containing
dissolved salts (e.g., sodium chloride in water) are called electrolytes, as they
are able to conduct
electricity. As observed in the cytoplasm of cells, in blood, urine, plant saps and mineral waters, mixtures
of many different ions in solution usually do not form defined salts after
evaporation of the water. Therefore, their salt content is given for the
respective ions.
Color
Potassium
dichromate, a bright orange salt used as a pigment
Manganese
dioxide, an opaque black salt
Salts can appear to be clear and transparent (sodium
chloride), opaque, and even metallic and
lustrous (iron disulfide). In many cases the apparent opacity or transparency are only related to the
difference in size of the individual monocrystals. Since light reflects from
the grain boundaries (boundaries between
crystallites), larger crystals tend to be transparent,
while polycrystalline aggregates look like white
powders.
- yellow
(sodium
chromate)
- orange
(potassium dichromate)
- red
(potassium ferricyanide)
- mauve
(cobalt chloride hexahydrate)
- blue
(copper sulfate pentahydrate, ferric
hexacyanoferrate)
- purple
(potassium permanganate)
- green
(nickel chloride hexahydrate)
- white
(sodium
chloride)
- colorless
(magnesium sulfate heptahydrate)
- black
(manganese dioxide)
Taste
Different salts can elicit all five basic tastes, e.g., salty (sodium chloride), sweet (lead diacetate, which will cause lead poisoning if ingested), sour (potassium bitartrate), bitter (magnesium
sulfate),
and umami or savory (monosodium glutamate).
Odour
Salts of strong acids and strong bases ("strong salts") are non-volatile and odourless, whereas
salts of either weak acids or weak bases ("weak
salts")
may smell after the conjugate
acid
(e.g., acetates like acetic acid (vinegar) and cyanides like
hydrogen cyanide (almonds)) or the conjugate base
(e.g., ammonium salts like ammonia) of the component ions.
That slow, partial decomposition is usually accelerated by the presence of
water, since hydrolysis is the other half of the reversible reaction equation of formation of weak
salts.
Degenerationism
The name of a salt starts with the name of the cation
(e.g., sodium or ammonium) followed by the name of the anion
(e.g., chloride or acetate). Salts are often referred to only by
the name of the cation (e.g., sodium salt or ammonium salt) or by
the name of the anion (e.g., chloride or acetate).
Common salt-forming cations include:
Iron(II)
oxide (FeO)
Iron(III)
oxide (Fe2O3)
- Calcium Ca2+
- Iron
Fe2+ and Fe3+
- Magnesium Mg2+
- Potassium K+
- Pyridinium C5H5NH+
- Quaternary ammonium NR4+
- Sodium Na+
Common salt-forming anions (parent acids in parentheses
where available) include:
- Acetate CH3COO− (acetic
acid)
- Carbonate CO32− (carbonic
acid)
- Chloride Cl− (hydrochloric acid)
- Citrate HOC(COO−)(CH2COO−)2
(citric
acid)
- Cyanide C≡N− (N/A)
- Hydroxide OH− (N/A)
- Nitrate NO3− (nitric acid)
- Nitrite NO2− (nitrous
acid)
- Oxide
O2− (N/A)
- Phosphate PO43− (phosphoric
acid)
- Sulfate SO42− (sulfuric
acid)
Formation
Solid
lead(II) sulfate (PbSO4)
- A
base and an acid, e.g., NH3 + HCl → NH4Cl
- A
metal and an acid, e.g., Mg + H2SO4 → MgSO4 + H2
- A
metal and a non-metal, e.g., Ca + Cl2 → CaCl2
- A
base and an acid
anhydride, e.g., 2 NaOH + Cl2O → 2 NaClO + H2O
- An
acid and a basic
anhydride, e.g., 2 HNO3 + Na2O → 2 NaNO3 + H2O
Pb(NO3)2(aq)
+ Na2SO4(aq) → PbSO4(s) + 2 NaNO3(aq)
See also
- Acid salt also known as Hydrogen salt
- Alkali
salts also known as Basic
salt
- Edible salt
- Electrolyte
- Halide
- Ionic bonds
- Kosher salt
- Natron
- Old
Salt Route
- Road salt
- Salting the earth (the deliberate massive use of salt to render a
soil unsuitable for cultivation and thus discourage habitation)
- Sea salt
- Sodium
- Table salt
- Zwitterion
- Salinity
- Hypertension
- Bresle
method (The method used to test
for salt presence during coating applications.
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