Lewis acids and bases

From Wikipedia, the free encyclopedia

Diagram of some Lewis bases and acids

A Lewis acid is a chemical species that contains an empty orbital which is capable of accepting an electron pair from a Lewis base to form a Lewis adduct. A Lewis base, then, is any species that has a filled orbital containing an electron pair which is not involved in bonding but may form a dative bond with a Lewis acid to form a Lewis adduct. For example, NH₃ is a Lewis base, because it can donate its lone pair of electrons. Trimethylborane (Me₃B) is a Lewis acid as it is capable of accepting a lone pair. In a Lewis adduct, the Lewis acid and base share an electron pair furnished by the Lewis base, forming a dative bond.^[1] In the context of a specific chemical reaction between NH₃ and Me₃B, the lone pair from NH₃ will form a dative bond with the empty orbital of Me₃B to form an adduct NH₃•BMe₃. The terminology refers to the contributions of Gilbert N. Lewis.^[2]

Depicting adducts

In many cases, the interaction between the Lewis base and Lewis acid in a complex is indicated by an arrow indicating the Lewis base donating electrons toward the Lewis acid using the notation of a dative bond—for example, Me₃B←NH₃. Some sources indicate the Lewis base with a pair of dots (the explicit electrons being donated), which allows consistent representation of the transition from the base itself to the complex with the acid:

Me₃B +  :NH₃ → Me₃B:NH₃

A center dot may also be used to represent a Lewis adduct, such as Me₃B•NH₃. Another example is boron trifluoride diethyl etherate, BF₃•Et₂O. The center dot is also used to represent hydrate coordination in various crystals, as in MgSO₄·7H₂O for hydrated magnesium sulfate. In general, however, the donor–acceptor bond is viewed as simply somewhere along a continuum between idealized covalent bonding and ionic bonding.^[3]

Examples

Major structural changes accompany binding of the Lewis base to the coordinatively unsaturated, planar Lewis acid BF₃.

Classically, the term "Lewis acid" is restricted to trigonal planar species with an empty p orbital, such as BR₃ where R can be an organic substituent or a halide.^{[citation needed]} For the purposes of discussion, even complex compounds such as Et₃Al₂Cl₃ and AlCl₃ are treated as trigonal planar Lewis acids. Metal ions such as Na⁺, Mg²⁺, and Ce³⁺, which are invariably complexed with additional ligands, are often sources of coordinatively unsaturated derivatives that form Lewis adducts upon reaction with a Lewis base. Other reactions might simply be referred to as "acid-catalyzed" reactions. Some compounds, such as H₂O, are both Lewis acids and Lewis bases, because they can either accept a pair of electrons or donate a pair of electrons, depending upon the reaction.

Lewis acids are diverse. Simplest are those that react directly with the Lewis base. But more common are those that undergo a reaction prior to forming the adduct.

• Examples of Lewis acids based on the general definition of electron pair acceptor include:
  • the proton (H⁺) and acidic compounds onium ions, such as NH₄⁺ and H₃O⁺
  • high oxidation state transition metal cations, e.g., Fe³⁺;
  • other metal cations, such as Li⁺ and Mg²⁺, often as their aquo or ether complexes,
  • trigonal planar species, such as BF₃ and carbocations H₃C⁺
  • pentahalides of phosphorus, arsenic, and antimony
  • electron poor π-systems, such as enones and tetracyanoethylenes.

Again, the description of a Lewis acid is often used loosely. For example, in solution, bare protons do not exist.

Simple Lewis acids

Some of the most studied examples of such Lewis acids are the boron trihalides and organoboranes, but other compounds exhibit this behavior:

BF₃ + F⁻ → BF₄⁻

In this adduct, all four fluoride centres (or more accurately, ligands) are equivalent.

BF₃ + OMe₂ → BF₃OMe₂

Both BF₄⁻ and BF₃OMe₂ are Lewis base adducts of boron trifluoride.

In many cases, the adducts violate the octet rule, such as the triiodide anion:

I₂ + I⁻ → I₃⁻

The variability of the colors of iodine solutions reflects the variable abilities of the solvent to form adducts with the Lewis acid I₂.

In some cases, the Lewis acids are capable of binding two Lewis bases, a famous example being the formation of hexafluorosilicate:

SiF₄ + 2 F⁻ → SiF₆²⁻

Complex Lewis acids

Most compounds considered to be Lewis acids require an activation step prior to formation of the adduct with the Lewis base. Well known cases are the aluminium trihalides, which are widely viewed as Lewis acids. Aluminium trihalides, unlike the boron trihalides, do not exist in the form AlX₃, but as aggregates and polymers that must be degraded by the Lewis base.^[4] A simpler case is the formation of adducts of borane. Monomeric BH₃ does not exist appreciably, so the adducts of borane are generated by degradation of diborane:

B₂H₆ + 2 H⁻ → 2 BH₄⁻

In this case, an intermediate B₂H₇⁻ can be isolated.

Many metal complexes serve as Lewis acids, but usually only after dissociating a more weakly bound Lewis base, often water.

[Mg(H₂O)₆]²⁺ + 6 NH₃ → [Mg(NH₃)₆]²⁺ + 6 H₂O

H⁺ as Lewis acid

The proton (H⁺) ^[5] is one of the strongest but is also one of the most complicated Lewis acids. It is convention to ignore the fact that a proton is heavily solvated (bound to solvent). With this simplification in mind, acid-base reactions can be viewed as the formation of adducts:

• H⁺ + NH₃ → NH₄⁺
• H⁺ + OH⁻ → H₂O

Applications of Lewis acids

A typical example of a Lewis acid in action is in the Friedel–Crafts alkylation reaction.^[3] The key step is the acceptance by AlCl₃ of a chloride ion lone-pair, forming AlCl₄⁻ and creating the strongly acidic, that is, electrophilic, carbonium ion.

RCl +AlCl₃ → R⁺ + AlCl₄⁻

Lewis bases

A Lewis base is an atomic or molecular species where the highest occupied molecular orbital (HOMO) is highly localized. Typical Lewis bases are conventional amines such as ammonia and alkyl amines. Other common Lewis bases include pyridine and its derivatives. Some of the main classes of Lewis bases are

• amines of the formula NH_3−xR_x where R = alkyl or aryl. Related to these are pyridine and its derivatives.
• phosphines of the formula PR_3−xA_x, where R = alkyl, A = aryl.
• compounds of O, S, Se and Te in oxidation state 2, including water, ethers, ketones

The most common Lewis bases are anions. The strength of Lewis basicity correlates with the pK_a of the parent acid: acids with high pK_a's give good Lewis bases. As usual, a weaker acid has a stronger conjugate base.

• Examples of Lewis bases based on the general definition of electron pair donor include:
  • simple anions, such as H⁻ and F⁻.
  • other lone-pair-containing species, such as H₂O, NH₃, HO⁻, and CH₃⁻
  • complex anions, such as sulfate
  • electron rich π-system Lewis bases, such as ethyne, ethene, and benzene

The strength of Lewis bases have been evaluated for various Lewis acids, such as I₂, SbCl₅, and BF₃.<sup>[6]

Heats of binding of various bases to BF₃</sup>

Lewis base	Donor atom	Enthalpy of complexation (kJ/mol)
Et3N	N	135
quinuclidine	N	150
pyridine	N	128
Acetonitrile	N	60
Et2O	O	78.8
THF	O	90.4
acetone	O	76.0
EtOAc	O	75.5
DMA	O	112
DMSO	O	105
Tetrahydrothiophene	S	51.6
Trimethylphosphine	P	97.3

Applications of Lewis bases

Nearly all electron pair donors that form compounds by binding transition elements can be viewed as a collections of the Lewis bases – or ligands. Thus a large application of Lewis bases is to modify the activity and selectivity of metal catalysts. Chiral Lewis bases thus confer chirality on a catalyst, enabling asymmetric catalysis, which is useful for the production of pharmaceuticals.

Many Lewis bases are "multidentate," that is they can form several bonds to the Lewis acid. These multidentate Lewis bases are called chelating agents.

Hard and soft classification

Lewis acids and bases are commonly classified according to their hardness or softness. In this context hard implies small and nonpolarizable and soft indicates larger atoms that are more polarizable.

• typical hard acids: H⁺, alkali/alkaline earth metal cations, boranes, Zn²⁺
• typical soft acids: Ag⁺, Mo(0), Ni(0), Pt²⁺
• typical hard bases: ammonia and amines, water, carboxylates, fluoride and chloride
• typical soft bases: organophosphines, thioethers, carbon monoxide, iodide

For example, an amine will displace phosphine from the adduct with the acid BF₃. In the same way, bases could be classified. For example, bases donating a lone pair from an oxygen atom are harder than bases donating through a nitrogen atom. Although the classification was never quantified it proved to be very useful in predicting the strength of adduct formation, using the key concepts that hard acid — hard base and soft acid — soft base interactions are stronger than hard acid — soft base or soft acid — hard base interactions. Later investigation of the thermodynamics of the interaction suggested that hard—hard interactions are enthalpy favored, whereas soft—soft are entropy favored.

ECW model

The ECW model is quantitative model that describes and predicts the strength of Lewis acid base interactions, −ΔH . The model assigned E and C parameters to many Lewis acids and bases. Each acid is characterized by an E_A and a C_A. Each base is likewise characterized by its own E_B and C_B. The E and C parameters refer, respectively, to the electrostatic and covalent contributions to the strength of the bonds that the acid and base will form. The equation is

−ΔH = E_AE_B + C_AC_B + W

The W term represents a constant energy contribution for acid–base reaction such as the cleavage of a dimeric acid or base. The equation predicts reversal of acids and base strengths. The graphical presentations of the equation show that there is no single order of Lewis base strengths or Lewis acid strengths.^[7]

There is no single order of Lewis base strengths

Cramer–Bopp plots show graphically using the E and C parameters of the ECW model that there is no one single order of Lewis base strengths (or acid strengths).^[8] Single property or variable scales are limited to a small range of acids or bases.

History

MO diagram depicting the formation of a dative covalent bond between two atoms.

The concept originated with Gilbert N. Lewis who studied chemical bonding. In 1923, Lewis wrote An acid substance is one which can employ an electron lone pair from another molecule in completing the stable group of one of its own atoms.^[9][10] The Brønsted–Lowry acid–base theory was published in the same year. The two theories are distinct but complementary. A Lewis base is also a Brønsted–Lowry base, but a Lewis acid doesn't need to be a Brønsted–Lowry acid. The classification into hard and soft acids and bases (HSAB theory) followed in 1963. The strength of Lewis acid-base interactions, as measured by the standard enthalpy of formation of an adduct can be predicted by the Drago–Wayland two-parameter equation.

Reformulation of Lewis theory

Lewis had suggested in 1916 that two atoms are held together in a chemical bond by sharing a pair of electrons. When each atom contributed one electron to the bond it was called a covalent bond. When both electrons come from one of the atoms it was called a dative covalent bond or coordinate bond. The distinction is not very clear-cut. For example, in the formation of an ammonium ion from ammonia and hydrogen the ammonia molecule donates a pair of electrons to the proton;^[5] the identity of the electrons is lost in the ammonium ion that is formed. Nevertheless, Lewis suggested that an electron-pair donor be classified as a base and an electron-pair acceptor be classified as acid.

A more modern definition of a Lewis acid is an atomic or molecular species with a localized empty atomic or molecular orbital of low energy. This lowest energy molecular orbital (LUMO) can accommodate a pair of electrons.

Comparison with Brønsted–Lowry theory

A Lewis base is often a Brønsted–Lowry base as it can donate a pair of electrons to H⁺;^[5] the proton is a Lewis acid as it can accept a pair of electrons. The conjugate base of a Brønsted–Lowry acid is also a Lewis base as loss of H⁺ from the acid leaves those electrons which were used for the A—H bond as a lone pair on the conjugate base. However, a Lewis base can be very difficult to protonate, yet still react with a Lewis acid. For example, carbon monoxide is a very weak Brønsted–Lowry base but it forms a strong adduct with BF₃.

In another comparison of Lewis and Brønsted–Lowry acidity by Brown and Kanner,^[11] 2,6-di-t-butylpyridine reacts to form the hydrochloride salt with HCl but does not react with BF₃. This example demonstrates that steric factors, in addition to electron configuration factors, play a role in determining the strength of the interaction between the bulky di-t-butylpyridine and tiny proton.

A Brønsted–Lowry acid is a proton donor, not an electron-pair acceptor.

Radical (chemistry)

From Wikipedia, the free encyclopedia

The hydroxyl radical, Lewis structure shown, contains one unpaired electron

In chemistry, a radical (more precisely, a free radical) is an atom, molecule, or ion that has an unpaired valence electron.^[1][2] With some exceptions, these unpaired electrons make free radicals highly chemically reactive. Many free radicals spontaneously dimerize. Most organic radicals have short lifetimes.

A notable example of a free radical is the hydroxyl radical (HO•), a molecule that has one unpaired electron on the oxygen atom. Two other examples are triplet oxygen and triplet carbene (:CH
₂) which have two unpaired electrons.

Free radicals may be generated in a number of ways, but typical methods involve redox reactions. Ionizing radiation, heat, electrical discharges, electrolysis, are known to produce radicals. Radicals are intermediates in many chemical reactions, more so than is apparent from the balanced equations.

Free radicals are important in combustion, atmospheric chemistry, polymerization, plasma chemistry, biochemistry, and many other chemical processes. A large fraction of natural products are generated by radical-generating enzymes. In living organisms, the free radicals superoxide and nitric oxide and their reaction products regulate many processes, such as control of vascular tone and thus blood pressure. They also play a key role in the intermediary metabolism of various biological compounds. Such radicals can even be messengers in a process dubbed redox signaling. A radical may be trapped within a solvent cage or be otherwise bound.

Depiction in chemical reactions

In chemical equations, free radicals are frequently denoted by a dot placed immediately to the right of the atomic symbol or molecular formula as follows:

${\displaystyle \mathrm {Cl} _{2}\;{\xrightarrow {UV}}\;{\mathrm {Cl} \cdot }+{\mathrm {Cl} \cdot }}$

Chlorine gas can be broken down by ultraviolet light to form atomic chlorine radicals.

Radical reaction mechanisms use single-headed arrows to depict the movement of single electrons:

The homolytic cleavage of the breaking bond is drawn with a 'fish-hook' arrow to distinguish from the usual movement of two electrons depicted by a standard curly arrow. The second electron of the breaking bond also moves to pair up with the attacking radical electron; this is not explicitly indicated in this case.

Free radicals also take part in radical addition and radical substitution as reactive intermediates. Chain reactions involving free radicals can usually be divided into three distinct processes. These are initiation, propagation, and termination.

• Initiation reactions are those that result in a net increase in the number of free radicals. They may involve the formation of free radicals from stable species as in Reaction 1 above or they may involve reactions of free radicals with stable species to form more free radicals.
• Propagation reactions are those reactions involving free radicals in which the total number of free radicals remains the same.
• Termination reactions are those reactions resulting in a net decrease in the number of free radicals. Typically two free radicals combine to form a more stable species, for example: 2Cl·→ Cl₂

Formation

The formation of radicals may involve the breaking of covalent bonds by homolysis, a process that requires significant amounts of energy. Such energies are known as homolytic bond dissociation energies, usually abbreviated as "ΔH °". Splitting H₂ into 2H•, for example, requires a ΔH ° of +435 kJ·mol^-1, while splitting Cl₂ into 2Cl• requires a ΔH ° of +243 kJ·mol^-1.

The energy needed to break a specific bond (generally covalent) between two atoms known as bond energy is a result of all the relative attractions and repulsions between the atoms of the molecule, however the most relevant are the bond's atoms and the immediate neighbors. As an approximation the most important parameters that influence the bonding between two atoms in a molecule are the mutual energy match and overlap of covalent orbitals and the repulsion between nonbonding orbitals. Likewise, radicals requiring more energy to form are less stable than those requiring less energy. An additional barrier can be the selection rule. Propagation, however, is very exothermic.

Radical formation through homolytic bond cleavage most often happens between two atoms of similar electronegativity; in organic chemistry, this is often between the O–O bond in peroxide species or between O–N bonds. Radicals may also be formed by single-electron oxidation or reduction of an atom or molecule: an example is the production of superoxide by the electron transport chain. Early studies in organometallic chemistry – especially F. A. Paneth and K. Hahnfeld's studies of tetra-alkyl lead species during the 1930s – supported the heterolytic fission of bonds and a radical-based mechanism. Although radical ions do exist, most species are electrically neutral.

Persistence and stability

The radical derived from α-tocopherol

Although radicals are generally short-lived due to their reactivity, there are long-lived radicals. These are categorized as follows:

Stable radicals

The prime example of a stable radical is molecular dioxygen (O₂). Another common example is nitric oxide (NO). Organic radicals can be long lived if they occur in a conjugated π system, such as the radical derived from α-tocopherol (vitamin E). There are also hundreds of examples of thiazyl radicals, which show low reactivity and remarkable thermodynamic stability with only a very limited extent of π resonance stabilization.^[3][4]

Persistent radicals

Persistent radical compounds are those whose longevity is due to steric crowding around the radical center, which makes it physically difficult for the radical to react with another molecule.^[5] Examples of these include Gomberg's triphenylmethyl radical, Fremy's salt (Potassium nitrosodisulfonate, (KSO₃)₂NO·), aminoxyls, (general formula R₂NO·) such as TEMPO, TEMPOL, nitronyl nitroxides, and azephenylenyls and radicals derived from PTM (perchlorophenylmethyl radical) and TTM (tris(2,4,6-trichlorophenyl)methyl radical). Persistent radicals are generated in great quantity during combustion, and "may be responsible for the oxidative stress resulting in cardiopulmonary disease and probably cancer that has been attributed to exposure to airborne fine particles."^[6]

Diradicals

Diradicals are molecules containing two radical centers. Multiple radical centers can exist in a molecule. Atmospheric oxygen naturally exists as a diradical in its ground state as triplet oxygen. The low reactivity of atmospheric oxygen is due to its diradical state. Non-radical states of dioxygen are actually less stable than the diradical. The relative stability of the oxygen diradical is primarily due to the spin-forbidden nature of the triplet-singlet transition required for it to grab electrons, i.e., "oxidize". The diradical state of oxygen also results in its paramagnetic character, which is demonstrated by its attraction to an external magnet.^[7] Diradicals can also occur in metal-oxo complexes, lending themselves for studies of spin forbidden reactions in transition metal chemistry.^[8]

Reactivity

Radical alkyl intermediates are stabilized by similar physical processes to carbocations: as a general rule, the more substituted the radical center is, the more stable it is. This directs their reactions. Thus, formation of a tertiary radical (R₃C·) is favored over secondary (R₂HC·), which is favored over primary (RH₂C·). Likewise, radicals next to functional groups such as carbonyl, nitrile, and ether are more stable than tertiary alkyl radicals.

Radicals attack double bonds. However, unlike similar ions, such radical reactions are not as much directed by electrostatic interactions. For example, the reactivity of nucleophilic ions with α,β-unsaturated compounds (C=C–C=O) is directed by the electron-withdrawing effect of the oxygen, resulting in a partial positive charge on the carbonyl carbon. There are two reactions that are observed in the ionic case: the carbonyl is attacked in a direct addition to carbonyl, or the vinyl is attacked in conjugate addition, and in either case, the charge on the nucleophile is taken by the oxygen. Radicals add rapidly to the double bond, and the resulting α-radical carbonyl is relatively stable; it can couple with another molecule or be oxidized. Nonetheless, the electrophilic/neutrophilic character of radicals has been shown in a variety of instances. One example is the alternating tendency of the copolymerization of maleic anhydride (electrophilic) and styrene (slightly nucleophilic).

In intramolecular reactions, precise control can be achieved despite the extreme reactivity of radicals. In general, radicals attack the closest reactive site the most readily. Therefore, when there is a choice, a preference for five-membered rings is observed: four-membered rings are too strained, and collisions with carbons six or more atoms away in the chain are infrequent.

Triplet carbenes and nitrenes, which are diradicals, have distinctive chemistry.

Combustion

Spectrum of the blue flame from a butane torch showing excited molecular radical band emission and Swan bands

A familiar free-radical reaction is combustion. The oxygen molecule is a stable diradical, best represented by ·O-O·. Because spins of the electrons are parallel, this molecule is stable. While the ground state of oxygen is this unreactive spin-unpaired (triplet) diradical, an extremely reactive spin-paired (singlet) state is available. For combustion to occur, the energy barrier between these must be overcome. This barrier can be overcome by heat, requiring high temperatures. The triplet-singlet transition is also "forbidden". This presents an additional barrier to the reaction. It also means molecular oxygen is relatively unreactive at room temperature except in the presence of a catalytic heavy atom such as iron or copper.

Combustion consists of various radical chain reactions that the singlet radical can initiate. The flammability of a given material strongly depends on the concentration of free radicals that must be obtained before initiation and propagation reactions dominate leading to combustion of the material. Once the combustible material has been consumed, termination reactions again dominate and the flame dies out. As indicated, promotion of propagation or termination reactions alters flammability. For example, because lead itself deactivates free radicals in the gasoline-air mixture, tetraethyl lead was once commonly added to gasoline. This prevents the combustion from initiating in an uncontrolled manner or in unburnt residues (engine knocking) or premature ignition (preignition).

When a hydrocarbon is burned, a large number of different oxygen radicals are involved. Initially, hydroperoxyl radical (HOO·) are formed. These then react further to give organic hydroperoxides that break up into hydroxyl radicals (HO·).

Polymerization

In addition to combustion, many polymerization reactions involve free radicals. As a result, many plastics, enamels, and other polymers are formed through radical polymerization. For instance, drying oils and alkyd paints harden due to radical crosslinking by oxygen from the atmosphere.

Recent advances in radical polymerization methods, known as living radical polymerization, include:

• Reversible addition-fragmentation chain transfer (RAFT)
• Atom transfer radical polymerization (ATRP)
• Nitroxide mediated polymerization (NMP)

These methods produce polymers with a much narrower distribution of molecular weights.

Atmospheric radicals

The most common radical in the lower atmosphere is molecular dioxygen. Photodissociation of source molecules produces other free radicals. In the lower atmosphere, important free radical are produced by the photodissociation of nitrogen dioxide to an oxygen atom and nitric oxide (see eq. 1. 1 below), which plays a key role in smog formation—and the photodissociation of ozone to give the excited oxygen atom O(1D) (see eq. 1. 2 below). The net and return reactions are also shown (eq. 1. 3 and eq. 1. 4, respectively).

${\displaystyle {\ce {{NO2}->[h \nu] {NO}+ {O}}}}$

(eq. 1. 1)

${\displaystyle {\ce {{O}+ {O2}-> {O3}}}}$

(eq. 1. 2)

${\displaystyle {\ce {{NO2}+ {O2}->[h \nu] {NO}+ {O3}}}}$

(eq. 1. 3)

${\displaystyle {\ce {{NO}+ {O3}-> {NO2}+ {O2}}}}$

(eq. 1. 4)

In the upper atmosphere, the photodissociation of normally unreactive chlorofluorocarbons (CFCs) by solar ultraviolet radiation is an important source of radicals (see eq. 1 below). These reactions give the chlorine radical, Cl•, which catalyzes the conversion of ozone to O₂, i.e., Ozone depletion (eq. 2. 2–eq. 2. 4 below).

${\displaystyle {\ce {{CFCS}->[h \nu] {Cl.}}}}$

(eq. 2. 1)

${\displaystyle {\ce {{Cl.}+ {O3}-> {ClO.}+ {O2}}}}$

(eq. 2. 2)

${\displaystyle {\ce {{O3}->[h \nu] {O}+ {O2}}}}$

(eq. 2. 3)

${\displaystyle {\ce {{O}+ {ClO.}-> {Cl.}+ {O2}}}}$

(eq. 2. 4)

${\displaystyle {\ce {{2O3}->[h \nu] 3O2}}}$

(eq. 2. 5)

Such reactions cause the depletion of the ozone layer, especially since the chlorine radical is free to engage in another reaction chain; consequently, the use of chlorofluorocarbons as refrigerants has been restricted.

In biology

Free radicals play important roles in biology. Many of these are necessary for life, such as the intracellular killing of bacteria by phagocytic cells such as granulocytes and macrophages. Free radicals are involved in cell signalling processes,^[9] known as redox signaling. For example, free radical attack of linoleic acid produces a series of 13-Hydroxyoctadecadienoic acids and 9-Hydroxyoctadecadienoic acids, which may act to regulate localized tissue inflammatory and/or healing responses, pain perception, and the proliferation of malignant cells. Free radical attacks on arachidonic acid and docosahexaenoic acid produce a similar but broader array of signaling products.^[10]
 

Structure of the deoxyadenosyl radical, a common biosynthetic intermediate.^[11]

Free radicals may also be involved in Parkinson's disease, senile and drug-induced deafness, schizophrenia, and Alzheimer's.^[12] The classic free-radical syndrome, the iron-storage disease hemochromatosis, is typically associated with a constellation of free-radical-related symptoms including movement disorder, psychosis, skin pigmentary melanin abnormalities, deafness, arthritis, and diabetes mellitus. The free-radical theory of aging proposes that free radicals underlie the aging process itself. Similarly, the process of mitohormesis suggests that repeated exposure to free radicals may extend life span.

Because free radicals are necessary for life, the body has a number of mechanisms to minimize free-radical-induced damage and to repair damage that occurs, such as the enzymes superoxide dismutase, catalase, glutathione peroxidase and glutathione reductase. In addition, antioxidants play a key role in these defense mechanisms. These are often the three vitamins, vitamin A, vitamin C and vitamin E and polyphenol antioxidants. Furthermore, there is good evidence indicating that bilirubin and uric acid can act as antioxidants to help neutralize certain free radicals. Bilirubin comes from the breakdown of red blood cells' contents, while uric acid is a breakdown product of purines. Too much bilirubin, though, can lead to jaundice, which could eventually damage the central nervous system, while too much uric acid causes gout.^[13]

Reactive oxygen species

Reactive oxygen species or ROS are species such as superoxide, hydrogen peroxide, and hydroxyl radical, commonly associated with cell damage. ROS form as a natural by-product of the normal metabolism of oxygen and have important roles in cell signaling. Two important oxygen-centered free radicals are superoxide and hydroxyl radical. They derive from molecular oxygen under reducing conditions. However, because of their reactivity, these same free radicals can participate in unwanted side reactions resulting in cell damage. Excessive amounts of these free radicals can lead to cell injury and death, which may contribute to many diseases such as cancer, stroke, myocardial infarction, diabetes and major disorders.^[14] Many forms of cancer are thought to be the result of reactions between free radicals and DNA, potentially resulting in mutations that can adversely affect the cell cycle and potentially lead to malignancy.^[15] Some of the symptoms of aging such as atherosclerosis are also attributed to free-radical induced oxidation of cholesterol to 7-ketocholesterol.^[16] In addition free radicals contribute to alcohol-induced liver damage, perhaps more than alcohol itself. Free radicals produced by cigarette smoke are implicated in inactivation of alpha 1-antitrypsin in the lung. This process promotes the development of emphysema.

Oxybenzone has been found to form free radicals in sunlight, and therefore may be associated with cell damage as well. This only occurred when it was combined with other ingredients commonly found in sunscreens, like titanium oxide and octyl methoxycinnamate.^[17]

ROS attack the polyunsaturated fatty acid, linoleic acid, to form a series of 13-Hydroxyoctadecadienoic acid and 9-Hydroxyoctadecadienoic acid products that serve as signaling molecules that may trigger responses that counter the tissue injury which caused their formation. ROS attacks other polyunsaturated fatty acids, e.g. arachidonic acid and docosahexaenoic acid, to produce a similar series of signaling products.^[18]

History and nomenclature

Moses Gomberg (1866–1947), the founder of radical chemistry

Until late in the 20th century the word "radical" was used in chemistry to indicate any connected group of atoms, such as a methyl group or a carboxyl, whether it was part of a larger molecule or a molecule on its own. The qualifier "free" was then needed to specify the unbound case. Following recent nomenclature revisions, a part of a larger molecule is now called a functional group or substituent, and "radical" now implies "free". However, the old nomenclature may still appear in some books.

The term radical was already in use when the now obsolete radical theory was developed. Louis-Bernard Guyton de Morveau introduced the phrase "radical" in 1785 and the phrase was employed by Antoine Lavoisier in 1789 in his Traité Élémentaire de Chimie. A radical was then identified as the root base of certain acids (the Latin word "radix" meaning "root"). Historically, the term radical in radical theory was also used for bound parts of the molecule, especially when they remain unchanged in reactions. These are now called functional groups. For example, methyl alcohol was described as consisting of a methyl "radical" and a hydroxyl "radical". Neither are radicals in the modern chemical sense, as they are permanently bound to each other, and have no unpaired, reactive electrons; however, they can be observed as radicals in mass spectrometry when broken apart by irradiation with energetic electrons.

In a modern context the first organic (carbon–containing) free radical identified was triphenylmethyl radical, (C₆H₅)₃C•. This species was discovered by Moses Gomberg in 1900. In 1933 Morris Kharash and Frank Mayo proposed that free radicals were responsible for anti-Markovnikov addition of hydrogen bromide to allyl bromide.^[19][20]

In most fields of chemistry, the historical definition of radicals contends that the molecules have nonzero electron spin. However, in fields including spectroscopy, chemical reaction, and astrochemistry, the definition is slightly different. Gerhard Herzberg, who won the Nobel prize for his research into the electron structure and geometry of radicals, suggested a looser definition of free radicals: "any transient (chemically unstable) species (atom, molecule, or ion)".^[21] The main point of his suggestion is that there are many chemically unstable molecules that have zero spin, such as C₂, C₃, CH₂ and so on. This definition is more convenient for discussions of transient chemical processes and astrochemistry; therefore researchers in these fields prefer to use this loose definition.^[22]

Diagnostics

Radicals typically exhibit paramagnetism, but the bulk magnetic properties of a ion or molecule are often not conveniently measured. Electron spin resonance is instead the definitive and most widely used technique for characterizing free radicals. The nature of the atom bearing the unpaired electron and its neighboring atoms can often be deduced by the EPR spectrum.^[23]

The presence of free radicals can also be detected or inferred by chemical reagents that trap (i.e. combine with) radicals. Often these traps are themselves radicals, such as TEMPO.