Le Chatelier's principle

Le Chatelier's principle (UK: /lə ʃæˈtɛljeɪ/, US: /ˈʃɑːtəljeɪ/), also called Chatelier's principle or "The Equilibrium Law", can be used to predict the effect of a change in conditions on some chemical equilibria. The principle is named after Henry Louis Le Chatelier and sometimes Karl Ferdinand Braun who discovered it independently. It can be stated as:

When any system at equilibrium for a long period of time is subjected to change in concentration, temperature, volume, or pressure, (1) the system changes to a new equilibrium and (2) this change partly counteracts the applied change.

It is common to treat the principle as a more general observation,[1] such as

When a settled system is disturbed, it will adjust to diminish the change that has been made to it,

or, "roughly stated",[1]

Any change in status quo prompts an opposing reaction in the responding system.

The principle has a variety of names, depending upon the discipline using it (see homeostasis, a term commonly used in biology).

In chemistry, the principle is used to manipulate the outcomes of reversible reactions, often to increase the yield of reactions. In pharmacology, the binding of ligands to the receptor may shift the equilibrium according to Le Chatelier's principle, thereby explaining the diverse phenomena of receptor activation and desensitization.[2] In economics, the principle has been generalized to help explain the price equilibrium of efficient economic systems.

Phenomena in apparent contradiction to Le Chatelier's principle can arise in systems of simultaneous equilibrium: see the article on the theory of response reactions.

Status as a physical law

Le Chatelier's principle describes the qualitative behavior of systems where there is an externally induced, instantaneous change in one parameter of a system; it states that a behavioural shift occurs in the system so as to oppose (partly cancel) the parameter change. The duration of adjustment depends on the strength of the negative feedback to the initial shock. Where a shock initially induces positive feedback (such as thermal runaway), the new equilibrium can be far from the old one, and can take a long time to reach. In some dynamic systems, the end-state cannot be determined from the shock. The principle is typically used to describe closed negative-feedback systems, but applies, in general, to thermodynamically closed and isolated systems in nature, since the second law of thermodynamics ensures that the disequilibrium caused by an instantaneous shock must have a finite half-life.[3] The principle has analogs throughout the entire physical world.

The principle while well rooted in chemical equilibrium and extended into economic theory, can also be used in describing mechanical systems in that the system put under stress will respond in a way such as to reduce or minimize that stress. Moreover, the response will generally be via the mechanism that most easily relieves that stress. Shear pins and other such sacrificial devices are design elements that protect systems against stress applied in undesired manners to relieve it so as to prevent more extensive damage to the entire system, a practical engineering application of Le Chatelier's principle.


Effect of change in concentration

Changing the concentration of a chemical will shift the equilibrium to the side that would reduce that change in concentration. The chemical system will attempt to partly oppose the change affected to the original state of equilibrium. In turn, the rate of reaction, extent, and yield of products will be altered corresponding to the impact on the system.

This can be illustrated by the equilibrium of carbon monoxide and hydrogen gas, reacting to form methanol.

CO + 2 H2 ⇌ CH3OH

Suppose we were to increase the concentration of CO in the system. Using Le Chatelier's principle, we can predict that the amount of methanol will increase, decreasing the total change in CO. If we are to add a species to the overall reaction, the reaction will favor the side opposing the addition of the species. Likewise, the subtraction of a species would cause the reaction to "fill the gap" and favor the side where the species was reduced. This observation is supported by the collision theory. As the concentration of CO is increased, the frequency of successful collisions of that reactant would increase also, allowing for an increase in forward reaction, and generation of the product. Even if the desired product is not thermodynamically favored, the end-product can be obtained if it is continuously removed from the solution.

The effect of a change in concentration is often exploited synthetically for condensation reactions (i.e., reactions that extrude water) that are equilibrium processes (e.g., formation of an ester from carboxylic acid and alcohol or an imine from an amine and aldehyde). This can be achieved by physically sequestering water, by adding desiccants like anhydrous magnesium sulfate or molecular sieves, or by continuous removal of water by distillation, often facilitated by a Dean-Stark apparatus.

Effect of change in temperature

The reversible reaction N2O4(g) ⇌ 2NO2(g) is endothermic, so the equilibrium position can be shifted by changing the temperature.
When heat is added and the temperature increases, the reaction shifts to the right and the flask turns reddish brown due to an increase in NO2. This demonstrates Le Chatelier's Principle: the equilibrium shifts in the direction that consumes energy.
When heat is removed and the temperature decreases, the reaction shifts to the left and flask turns colorless due to an increase in N2O4: again, according to the Principle.

The effect of changing the temperature in the equilibrium can be made clear by 1) incorporating heat as either a reactant or a product, and 2) assuming that an increase in temperature increases the heat content of a system. When the reaction is exothermicH is negative, puts energy out), heat is included as a product, and, when the reaction is endothermicH is positive, takes energy in), heat is included as a reactant. Hence, whether increasing or decreasing the temperature would favor the forward or the reverse reaction can be determined by applying the same principle as with concentration changes.

Take, for example, the reversible reaction of nitrogen gas with hydrogen gas to form ammonia:

N2(g) + 3 H2(g) ⇌ 2 NH3(g)    ΔH = -92 kJ mol−1

Because this reaction is exothermic, it produces heat:

N2(g) + 3 H2(g) ⇌ 2 NH3(g) + heat

If the temperature were increased, the heat content of the system would increase, so the system would consume some of that heat by shifting the equilibrium to the left, thereby producing less ammonia. More ammonia would be produced if the reaction were run at a lower temperature, but a lower temperature also lowers the rate of the process, so, in practice (the Haber process) the temperature is set at a compromise value that allows ammonia to be made at a reasonable rate with an equilibrium concentration that is not too unfavorable.

In exothermic reactions, increase in temperature decreases the equilibrium constant, K, whereas, in endothermic reactions, increase in temperature increases the K value.

Le Chatelier's principle applied to changes in concentration or pressure can be understood by having K have a constant value. The effect of temperature on equilibria, however, involves a change in the equilibrium constant. The dependence of K on temperature is determined by the sign of ΔH. The theoretical basis of this dependence is given by the Van 't Hoff equation.

Effect of change in pressure

The equilibrium concentrations of the products and reactants do not directly depend on the total pressure of the system. They may depend on the partial pressures of the products and reactants, but if the number of moles of gaseous reactants is equal to the number of moles of gaseous products, pressure has no effect on equilibrium.

Changing total pressure by adding an inert gas at constant volume does not affect the equilibrium concentrations (see §Effect of adding an inert gas below).

Changing total pressure by changing the volume of the system changes the partial pressures of the products and reactants and can affect the equilibrium concentrations (see §Effect of change in volume below).

Effect of change in volume

Changing the volume of the system changes the partial pressures of the products and reactants and can affect the equilibrium concentrations. With a pressure increase due to a decrease in volume, the side of the equilibrium with fewer moles is more favorable[4] and with a pressure decrease due to an increase in volume, the side with more moles is more favorable. There is no effect on a reaction where the number of moles of gas is the same on each side of the chemical equation.

Considering the reaction of nitrogen gas with hydrogen gas to form ammonia:

    ΔH = -92kJ mol−1

Note the number of moles of gas on the left-hand side and the number of moles of gas on the right-hand side. When the volume of the system is changed, the partial pressures of the gases change. If we were to decrease pressure by increasing volume, the equilibrium of the above reaction will shift to the left, because the reactant side has a greater number of moles than does the product side. The system tries to counteract the decrease in partial pressure of gas molecules by shifting to the side that exerts greater pressure. Similarly, if we were to increase pressure by decreasing volume, the equilibrium shifts to the right, counteracting the pressure increase by shifting to the side with fewer moles of gas that exert less pressure. If the volume is increased because there are more moles of gas on the reactant side, this change is more significant in the denominator of the equilibrium constant expression, causing a shift in equilibrium.

Effect of adding an inert gas

An inert gas (or noble gas), such as helium, is one that does not react with other elements or compounds. Adding an inert gas into a gas-phase equilibrium at constant volume does not result in a shift.[4] This is because the addition of a non-reactive gas does not change the equilibrium equation, as the inert gas appears on both sides of the chemical reaction equation. For example, if A and B react to form C and D, but X does not participate in the reaction: . While it is true that the total pressure of the system increases, the total pressure does not have any effect on the equilibrium constant; rather, it is a change in partial pressures that will cause a shift in the equilibrium. If, however, the volume is allowed to increase in the process, the partial pressures of all gases would be decreased resulting in a shift towards the side with the greater number of moles of gas. The shift will never occur on the side with fewer moles of gas. It is also known as Le Chatelier's postulate.

Effect of a catalyst

A catalyst increases the rate of a reaction without being consumed in the reaction. The use of a catalyst does not affect the position and composition of the equilibrium of a reaction, because both the forward and backward reactions are sped up by the same factor.

For example, consider the Haber process for the synthesis of ammonia (NH3):

N2 + 3 H2 ⇌ 2 NH3

In the above reaction, iron (Fe) and molybdenum (Mo) will function as catalysts if present. They will accelerate any reactions, but they do not affect the state of the equilibrium.

General statements related to Le Chatelier's principle

For thermodynamics, when complicated equilibria are considered, it has turned out to be difficult or unfeasible to make valid and very general statements that echo Le Chatelier's principle.[5][6] Prigogine & Defay demonstrate that a thermodynamic system may or may not exhibit moderation, depending upon exactly what conditions are imposed after the perturbation.[7] Le Chatelier's principle refers to thermodynamic equilibria, which are stable, and mostly does not apply to metastable and unstable equilibria.

Applications in economics

In economics, a similar concept also named after Le Chatelier was introduced by U.S. economist Paul Samuelson in 1947. There the generalized Le Chatelier principle is for a maximum condition of economic equilibrium: Where all unknowns of a function are independently variable, auxiliary constraints—"just-binding" in leaving initial equilibrium unchanged—reduce the response to a parameter change. Thus, factor-demand and commodity-supply elasticities are hypothesized to be lower in the short run than in the long run because of the fixed-cost constraint in the short run.[8]

Since the change of the value of an objective function in a neighbourhood of the maximum position is described by the envelope theorem, Le Chatelier's principle can be shown to be a corollary thereof.[9]

See also


  1. ^ a b Gall, John (2002). The Systems Bible (3rd ed.). General Systemantics Press. The System always kicks back
  2. ^ "The Biophysical Basis for the Graphical Representations". Retrieved 2009-05-04.
  3. ^ Kay, J. J. (February 2000) [1999]. "Application of the Second Law of Thermodynamics and Le Chatelier's Principle to the Developing Ecosystem". In Muller, F. (ed.). Handbook of Ecosystem Theories and Management. Environmental & Ecological (Math) Modeling. CRC Press. ISBN 978-1-56670-253-9. As systems are moved away from equilibrium, they will utilize all available avenues to counter the applied gradients... Le Chatelier's principle is an example of this equilibrium seeking principle.
    For full details, see: "Ecosystems as Self-organizing Holarchic Open Systems: Narratives and the Second Law of Thermodynamics". 2000: 5. CiteSeerX Cite journal requires |journal= (help)
  4. ^ a b Atkins1993, p. 114
  5. ^ Münster, A. (1970), Classical Thermodynamics, translated by E.S. Halberstadt, Wiley–Interscience, London, ISBN 0-471-62430-6, pp. 173–174.
  6. ^ Prigogine, I., Defay, R. (1950/1954). Chemical Thermodynamics, Longmans, Green & Co, London, pp. 268–269.
  7. ^ Prigogine, I., Defay, R. (1950/1954). Chemical Thermodynamics, Longmans, Green & Co, London, p. 265.
  8. ^ Samuelson, Paul A (1983). Foundations of Economic Analysis. Harvard University Press. ISBN 0-674-31301-1.
  9. ^ Silberberg, Eugene (1971). "The Le Chatelier Principle as a Corollary to a Generalized Envelope Theorem". Journal of Economic Theory. 3 (2): 146–155. doi:10.1016/0022-0531(71)90012-3.


  • Atkins, P.W. (1993). The Elements of Physical Chemistry (3rd ed.). Oxford University Press.
  • Le Chatelier, H. and Boudouard O. (1898), "Limits of Flammability of Gaseous Mixtures", Bulletin de la Société Chimique de France (Paris), v. 19, pp. 483–488.
  • Hatta, Tatsuo (1987), "Le Chatelier principle," The New Palgrave: A Dictionary of Economics, v. 3, pp. 155–57.
  • Samuelson, Paul A. (1947, Enlarged ed. 1983). Foundations of Economic Analysis, Harvard University Press. ISBN 0-674-31301-1
  • D.J. Evans, D.J. Searles and E. Mittag (2001), "Fluctuation theorem for Hamiltonian systems—Le Chatelier's principle", Physical Review E, 63, 051105(4).

External links

1884 in science

The year 1884 in science and technology involved some significant events, listed below.


Antiperistasis, in philosophy, is a general term for various processes, real or contrived, in which one quality heightens the force of another, opposing, quality.

Chemical Monitoring and Management

The Chemical Monitoring and Management Module is part of the New South Wales, Higher School Certificate (HSC) Chemistry course studied by Secondary Students in their final year of schooling (Year 12). Students study four modules, 3 compulsory, and 1 of the 5 elective modules.The 3 compulsory modules are:

Identification and Production of Materials

The Acidic Environment

Chemical Monitoring and ManagementThe five option modules, of which one may be studied are:

Industrial Chemistry

Shipwrecks and Salvage

Forensic Chemistry

The Biochemistry of Movement

The Chemistry of ArtThe module "Chemical Monitoring and Management" is designed to teach students studying Chemistry:

The Role of Chemists in Monitoring and Management of Chemical Reactions

Various Methods of Chemical Analysis

The Production of Ammonia (The Haber/Bosch Process)

Chemical Equilibrium

Le Chatelier's Principle

The role of Catalysts

Identification of chemicals using chemical tests and Spectroscopy

The Chemical Monitroing and Management of the atmosphere and waterwaysThe syllabus was created by the New South Wales Board of Studies.

Climate change feedback

Climate change feedback is important in the understanding of global warming because feedback processes may amplify or diminish the effect of each climate forcing, and so play an important part in determining the climate sensitivity and future climate state. Feedback in general is the process in which changing one quantity changes a second quantity, and the change in the second quantity in turn changes the first. Positive feedback amplifies the change in the first quantity while negative feedback reduces it.The term "forcing" means a change which may "push" the climate system in the direction of warming or cooling. An example of a climate forcing is increased atmospheric concentrations of greenhouse gases. By definition, forcings are external to the climate system while feedbacks are internal; in essence, feedbacks represent the internal processes of the system. Some feedbacks may act in relative isolation to the rest of the climate system; others may be tightly coupled; hence it may be difficult to tell just how much a particular process contributes.Forcings and feedbacks together determine how much and how fast the climate changes. The main positive feedback in global warming is the tendency of warming to increase the amount of water vapor in the atmosphere, which in turn leads to further warming. The main negative feedback comes from the Stefan–Boltzmann law, the amount of heat radiated from the Earth into space changes with the fourth power of the temperature of Earth's surface and atmosphere. Observations and modelling studies indicate that there is a net positive feedback to warming. Large positive feedbacks can lead to effects that are abrupt or irreversible, depending upon the rate and magnitude of the climate change.

Common-ion effect

The common-ion effect states that in a chemical solution in which several species reversibly associate with each other by an equilibrium process, increasing the concentration of any one of its dissociated components by adding another chemical that also contains it will cause an increased amount of association. This result is a consequence of Le Chatelier's principle for the equilibrium reaction of the association/dissociation. The effect is commonly seen as an effect on the solubility of salts and other weak electrolytes. Adding an additional amount of one of the ions of the salt generally leads to increased precipitation of the salt, which reduces the concentration of both ions of the salt until the solubility equilibrium is reached. The effect is based on the fact that both the original salt and the other added chemical have one ion in common with each other.


Diimide, also called diazene or diimine, is a compound having the formula (NH)2. It exists as two geometric isomers, E (trans) and Z (cis). The term diazene is more common for organic derivatives of diimide. Thus, azobenzene is an example of an organic diazene.

Free base

Free base (freebase, free-base) is the conjugate base (deprotonated) form of an amine, as opposed to its conjugate acid (protonated) form. The amine is often an alkaloid, such as nicotine, cocaine, morphine, and ephedrine, or derivatives thereof.

Freebasing is a more efficient method of self-administering alkaloids via the smoking route.

Haldane effect

The Haldane effect is a property of haemoglobin first described by John Scott Haldane. Oxygenation of blood in the lungs displaces carbon dioxide from hemoglobin which increases the removal of carbon dioxide. This property is the Haldane effect. Consequently, oxygenated blood has a reduced affinity for carbon dioxide. Thus, the Haldane effect describes the ability of hemoglobin to carry increased amounts of carbon dioxide (CO2) in the deoxygenated state as opposed to the oxygenated state. A high concentration of CO2 facilitates dissociation of oxyhaemoglobin.


A hemiacetal or a hemiketal is a compound that results from the addition of an alcohol to an aldehyde or a ketone, respectively. The Greek word hèmi, meaning half(semi), refers to the fact that a single alcohol has been added to the carbonyl group, in contrast to acetals or ketals, which are formed when a second alkoxy group has been added to the structure.

Henry Louis Le Chatelier

Henry Louis Le Chatelier (French pronunciation: ​[ɑ̃ʁi lwi lə ʃɑtlje]; 8 October 1850 – 17 September 1936) was a French chemist of the late 19th and early 20th centuries. He devised Le Chatelier's principle, used by chemists to predict the effect a changing condition has on a system in chemical equilibrium.

Holon (philosophy)

A holon (Greek: ὅλον, holon neuter form of ὅλος, holos "whole") is something that is simultaneously a whole and a part. The word was used by Arthur Koestler in his book The Ghost in the Machine (1967, p. 48) and the phrase to hólon is a Greek translation from the Latin word universum, in the sense of totality, a whole.Koestler was influenced by two observations in proposing the notion of the holon. The first observation was influenced by Herbert A. Simon's parable of the two watchmakers—in which Simon concludes that complex systems evolve from simple systems much more rapidly when there are stable intermediate forms present in the evolutionary process than if they are not present.The second observation was made by Koestler himself in his analysis of hierarchies and stable intermediate forms in non-living matter (atomic and molecular structure), living organisms, and social organizations. He concluded that, although it is easy to identify sub-wholes or parts, wholes and parts in an absolute sense do not exist anywhere. Koestler proposed the word holon to describe the hybrid nature of sub-wholes and parts within in vivo systems. From this perspective, holons exist simultaneously as self-contained wholes in relation to their sub-ordinate parts, and as dependent parts when considered from the inverse direction.

Koestler also says that holons are self-reliant units that possess a degree of independence and can handle contingencies without asking higher authorities for instructions. I.e. they have a degree of autonomy. These holons are also simultaneously subject to control from one or more of these higher authorities. The first property ensures that holons are stable forms that are able to withstand disturbances, while the latter property signifies that they are intermediate forms, providing a context for the proper functionality for the larger whole.

Finally, Koestler defines a holarchy as a hierarchy of self-regulating holons that function first as autonomous wholes in supra-ordination to their parts, secondly as dependent parts in sub-ordination to controls on higher levels, and thirdly in coordination with their local environment.

Le Chatelier

Le Chatelier can refer to:

Alfred Le Chatelier (1855–1929), French soldier, explorer and professor

Bénédicte Le Chatelier (born 1976), French television journalist

Henry Louis Le Chatelier, 19th-century chemist

Le Châtelier's principle, named after Henry Louis

Louis Le Chatelier, 19th-century chemist and industrialist, father of Henri Louis

Le Châtelier, a commune in the Marne département, France

Pericyclic reaction

In organic chemistry, a pericyclic reaction is a type of organic reaction wherein the transition state of the molecule has a cyclic geometry, the reaction progresses in a concerted fashion, and the bond orbitals involved in the reaction overlap in a continuous cycle at the transition state. Pericyclic reactions stand in contrast to linear reactions, encompassing most organic transformations and proceeding through an acyclic transition state, on the one hand and coarctate reactions, which proceed through a doubly cyclic, concerted transition state on the other hand. Pericyclic reactions are usually rearrangement or addition reactions. The major classes of pericyclic reactions are given in the table below (the three most important classes are shown in bold). Ene reactions and cheletropic reactions are often classed as group transfer reactions and cycloadditions/reversions, respectively, while dyotropic reactions and group transfer reactions (if ene reactions are excluded) are rarely encountered.

In general, these are considered to be equilibrium processes, although it is possible to push the reaction in one direction by designing a reaction by which the product is at a significantly lower energy level; this is due to a unimolecular interpretation of Le Chatelier's principle. There is thus a set of "retro" pericyclic reactions.


Phenolphthalein is a chemical compound with the formula C20H14O4 and is often written as "HIn" or "phph" in shorthand notation. Phenolphthalein is often used as an indicator in acid–base titrations. For this application, it turns colorless in acidic solutions and pink in basic solutions. It belongs to the class of dyes known as phthalein dyes.

Phenolphthalein is slightly soluble in water and usually is dissolved in alcohols for use in experiments. It is a weak acid, which can lose H+ ions in solution. The phenolphthalein molecule is colorless, and the phenolphthalein ion is pink. When a base is added to the phenolphthalein, the equilibrium shifts, leading to more ionization as H+ ions are removed. This is predicted by Le Chatelier's principle.

Phenylalanine racemase (ATP-hydrolysing)

The enzyme phenylalanine racemase (EC, phenylalanine racemase, phenylalanine racemase (adenosine triphosphate-hydrolysing), gramicidin S synthetase I) is the enzyme that acts on amino acids and derivatives. It activates both the L & D stereo isomers of phenylalanine to form L-phenylalanyl adenylate and D-phenylalanyl adenylate, which are bound to the enzyme. These bound compounds are then transferred to the thiol group of the enzyme followed by conversion of its configuration, the D-isomer being the more favorable configuration of the two, with a 7 to 3 ratio between the two isomers. The racemisation reaction of phenylalanine is coupled with the highly favorable hydrolysis of adenosine triphosphate (ATP) to adenosine monophosphate (AMP) and pyrophosphate (PP), thermodynamically allowing it to proceed. This reaction is then drawn forward by further hydrolyzing PP to inorganic phosphate (Pi), via Le Chatelier's principle.

Reversible reaction

A reversible reaction is a reaction where the reactants form products, which react together to give the reactants back.

A and B can react to form C and D or, in the reverse reaction, C and D can react to form A and B. This is distinct from reversible process in thermodynamics.

Weak acids and bases undertake reversible reactions. For example, carbonic acid:

H2CO3 (l) + H2O(l) ⇌ HCO3(aq) + H3O+(aq).

The concentrations of reactants and products in an equilibrium mixture are determined by the analytical concentrations of the reagents (A and B or C and D) and the equilibrium constant, K. The magnitude of the equilibrium constant depends on the Gibbs free energy change for the reaction. So, when the free energy change is large (more than about 30 kJ mol−1), then the equilibrium constant is large (log K > 3) and the concentrations of the reactants at equilibrium are very small. Such a reaction is sometimes considered to be an irreversible reaction, although in reality small amounts of the reactants are still expected to be present in the reacting system. A truly irreversible chemical reaction is usually achieved when one of the products exits the reacting system, for example, as does carbon dioxide (volatile) in the reaction

CaCO3 + 2HCl → CaCl2 + H2O + CO2
Substrate (chemistry)

In chemistry, a substrate is typically the chemical species being observed in a chemical reaction, which reacts with a reagent to generate a product. In synthetic and organic chemistry, the substrate is the chemical of interest that is being modified. In biochemistry, an enzyme substrate is the material upon which an enzyme acts. When referring to Le Chatelier's principle, the substrate is the reagent whose concentration is changed. The term substrate is highly context-dependent. It essentially refers to the part of the molecule that is precursor to a product.

Triphenylmethyl radical

The triphenylmethyl radical (often shorted to trityl radical) is a persistent radical and the first radical ever described in organic chemistry. It can be prepared by homolysis of triphenylmethyl chloride 1 by a metal like silver or zinc in benzene or diethyl ether. The radical 2 forms a chemical equilibrium with the quinoid-type dimer 3 (3-triphenylmethyl-6-diphenylmethylidene-1,4-cyclohexadiene; Gomberg's dimer). In benzene the concentration of the radical is 2%.

Solutions containing the radical are yellow; when the temperature of the solution is raised, the yellow color becomes more intense as the equilibrium is shifted in favor of the radical (in accordance with Le Châtelier's principle).

When exposed to air, the radical rapidly oxidizes to the peroxide, and the color of the solution changes from yellow to colorless. Likewise, the radical reacts with iodine to triphenylmethyl iodide.

The radical was discovered by Moses Gomberg in 1900 at the University of Michigan. He tried to prepare hexaphenylethane from triphenylmethyl chloride and zinc in benzene in a Wurtz reaction and found that the product, based on its behaviour towards iodine and oxygen, was far more reactive than anticipated. The discovered structure was used in the development of ESR spectroscopy and confirmed by it.The correct quinoid structure for the dimer was suggested as early as 1904 but this structure was soon after abandoned by the scientific community in favor of hexaphenylethane (4). It subsequently took until 1968 for its rediscovery when researchers at the Vrije Universiteit Amsterdam published proton NMR data.While the trityl radical forms a quinoid dimer, derivatives thereof with the appropriate substitution pattern do form dimers with a hexaphenylethane structure. X-ray studies give a bond length of 1.67 Å for hexakis(3,5-di-t-butylphenyl)ethane. Theoretical calculations on a very high level of theory indicate that van der Waals attraction between the tert-butyl groups create a potential minimum that is absent in the unsubstituted molecule. Other derivatives have been reported as the quinoid dimer


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