Organic chemistry is a subdiscipline of chemistry that studies the structure and reactions of organic compounds, which contain carbon in covalent bonding. Study of structure determines their chemical formula. Study of properties includes physical and chemical properties, evaluation of chemical reactivity to understand their behavior; the study of organic reactions includes the chemical synthesis of natural products and polymers, study of individual organic molecules in the laboratory and via theoretical study. The range of chemicals studied in organic chemistry includes hydrocarbons as well as compounds based on carbon, but containing other elements oxygen, sulfur and the halogens. Organometallic chemistry is the study of compounds containing carbon–metal bonds. In addition, contemporary research focuses on organic chemistry involving other organometallics including the lanthanides, but the transition metals zinc, palladium, cobalt and chromium. Organic compounds constitute the majority of known chemicals.
The bonding patterns of carbon, with its valence of four—formal single and triple bonds, plus structures with delocalized electrons—make the array of organic compounds structurally diverse, their range of applications enormous. They form the basis of, or are constituents of, many commercial products including pharmaceuticals; the study of organic chemistry overlaps organometallic chemistry and biochemistry, but with medicinal chemistry, polymer chemistry, materials science. Before the nineteenth century, chemists believed that compounds obtained from living organisms were endowed with a vital force that distinguished them from inorganic compounds. According to the concept of vitalism, organic matter was endowed with a "vital force". During the first half of the nineteenth century, some of the first systematic studies of organic compounds were reported. Around 1816 Michel Chevreul started a study of soaps made from various alkalis, he separated the different acids. Since these were all individual compounds, he demonstrated that it was possible to make a chemical change in various fats, producing new compounds, without "vital force".
In 1828 Friedrich Wöhler produced the organic chemical urea, a constituent of urine, from inorganic starting materials, in what is now called the Wöhler synthesis. Although Wöhler himself was cautious about claiming he had disproved vitalism, this was the first time a substance thought to be organic was synthesized in the laboratory without biological starting materials; the event is now accepted as indeed disproving the doctrine of vitalism. In 1856 William Henry Perkin, while trying to manufacture quinine accidentally produced the organic dye now known as Perkin's mauve, his discovery, made known through its financial success increased interest in organic chemistry. A crucial breakthrough for organic chemistry was the concept of chemical structure, developed independently in 1858 by both Friedrich August Kekulé and Archibald Scott Couper. Both researchers suggested that tetravalent carbon atoms could link to each other to form a carbon lattice, that the detailed patterns of atomic bonding could be discerned by skillful interpretations of appropriate chemical reactions.
The era of the pharmaceutical industry began in the last decade of the 19th century when the manufacturing of acetylsalicylic acid—more referred to as aspirin—in Germany was started by Bayer. By 1910 Paul Ehrlich and his laboratory group began developing arsenic-based arsphenamine, as the first effective medicinal treatment of syphilis, thereby initiated the medical practice of chemotherapy. Ehrlich popularized the concepts of "magic bullet" drugs and of systematically improving drug therapies, his laboratory made decisive contributions to developing antiserum for diphtheria and standardizing therapeutic serums. Early examples of organic reactions and applications were found because of a combination of luck and preparation for unexpected observations; the latter half of the 19th century however witnessed systematic studies of organic compounds. The development of synthetic indigo is illustrative; the production of indigo from plant sources dropped from 19,000 tons in 1897 to 1,000 tons by 1914 thanks to the synthetic methods developed by Adolf von Baeyer.
In 2002, 17,000 tons of synthetic indigo were produced from petrochemicals. In the early part of the 20th century and enzymes were shown to be large organic molecules, petroleum was shown to be of biological origin; the multiple-step synthesis of complex organic compounds is called total synthesis. Total synthesis of complex natural compounds increased in complexity to terpineol. For example, cholesterol-related compounds have opened ways to synthesize complex human hormones and their modified derivatives. Since the start of the 20th century, complexity of total syntheses has been increased to include molecules of high complexity such as lysergic acid and vitamin B12; the discovery of petroleum and the development of the petrochemical industry spurred the development of organic chemistry. Converting individual petroleum compounds into different types of compounds by various chemical processes led to organic reactions enabling a broad range of
In chemistry, a coordination complex consists of a central atom or ion, metallic and is called the coordination centre, a surrounding array of bound molecules or ions, that are in turn known as ligands or complexing agents. Many metal-containing compounds those of transition metals, are coordination complexes. A coordination complex whose centre is a metal atom is called a metal complex. Coordination complexes are so pervasive that their structures and reactions are described in many ways, sometimes confusingly; the atom within a ligand, bonded to the central metal atom or ion is called the donor atom. In a typical complex, a metal ion is bonded to several donor atoms, which can be the same or different. A polydentate ligand is a molecule or ion that bonds to the central atom through several of the ligand's atoms; these complexes are called chelate complexes. The central atom or ion, together with all ligands, comprise the coordination sphere; the central atoms or ion and the donor atoms comprise the first coordination sphere.
Coordination refers to the "coordinate covalent bonds" between the central atom. A complex implied a reversible association of molecules, atoms, or ions through such weak chemical bonds; as applied to coordination chemistry, this meaning has evolved. Some metal complexes are formed irreversibly and many are bound together by bonds that are quite strong; the number of donor atoms attached to the central atom or ion is called the coordination number. The most common coordination numbers are 2, 4, 6. A hydrated ion is one kind of a complex ion, a species formed between a central metal ion and one or more surrounding ligands, molecules or ions that contain at least one lone pair of electrons. If all the ligands are monodentate the number of donor atoms equals the number of ligands. For example, the cobalt hexahydrate ion or the hexaaquacobalt ion 2+ is a hydrated-complex ion that consists of six water molecules attached to a metal ion Co; the oxidation state and the coordination number reflect the number of bonds formed between the metal ion and the ligands in the complex ion.
However, the coordination number of Pt2+2 is 4 since it has two bidentate ligands, which contain four donor atoms in total. Any donor atom will give a pair of electrons. There are some donor groups which can offer more than one pair of electrons; such are called polydentate. In some cases an atom or a group offers a pair of electrons to two similar or different central metal atoms or acceptors—by division of the electron pair—into a three-center two-electron bond; these are called bridging ligands. Coordination complexes have been known since the beginning of modern chemistry. Early well-known coordination complexes include dyes such as Prussian blue, their properties were first well understood in the late 1800s, following the 1869 work of Christian Wilhelm Blomstrand. Blomstrand developed; the theory claimed that the reason coordination complexes form is because in solution, ions would be bound via ammonia chains. He compared this effect to the way. Following this theory, Danish scientist Sophus Mads Jørgensen made improvements to it.
In his version of the theory, Jørgensen claimed that when a molecule dissociates in a solution there were two possible outcomes: the ions would bind via the ammonia chains Blomstrand had described or the ions would bind directly to the metal. It was not until 1893 that the most accepted version of the theory today was published by Alfred Werner. Werner’s work included two important changes to the Blomstrand theory; the first was that Werner described the two different ion possibilities in terms of location in the coordination sphere. He claimed that if the ions were to form a chain this would occur outside of the coordination sphere while the ions that bound directly to the metal would do so within the coordination sphere. In one of Werner’s most important discoveries however he disproved the majority of the chain theory. Werner was able to discover the spatial arrangements of the ligands that were involved in the formation of the complex hexacoordinate cobalt, his theory allows one to understand the difference between a coordinated ligand and a charge balancing ion in a compound, for example the chloride ion in the cobaltammine chlorides and to explain many of the inexplicable isomers.
In 1914, Werner first resolved the coordination complex, called hexol, into optical isomers, overthrowing the theory that only carbon compounds could possess chirality. The ions or molecules surrounding the central atom are called ligands. Ligands are bound to the central atom by a coordinate covalent bond, are said to be coordinated to the atom. There are organic ligands such as alkenes whose pi bonds can coordinate to empty metal orbitals. An example is ethene in the complex known as Zeise's salt, K+−. In coordination chemistry, a structure is first described by its coordination number, the number of ligands attached to the metal. One can count the ligands attached, but sometimes the counting can become ambiguous. Coordination numbers are between two and nine, but large numbers of ligands are not uncommon for the lanthanides and actinides; the number of bonds
Hexaamminecobalt chloride is the chemical compound with the formula Cl3. It is the chloride salt of the coordination complex 3+, considered an archetypal "Werner complex", named after the pioneer of coordination chemistry, Alfred Werner; the cation itself is a metal ammine complex with six ammonia ligands attached to the cobalt ion. Salts of 3+ were described as the luteo complex of cobalt; this name has been discarded as modern chemistry considers color less important than molecular structure. Other similar complexes had color names, such as purpureo for a cobalt pentammine complex, praseo and violeo for two isomeric tetrammine complexes. 3+ is diamagnetic, with a low-spin 3d6 octahedral Co center. The cation obeys the 18-electron rule and is considered to be a classic example of an exchange inert metal complex; as a manifestation of its inertness, Cl3 can be recrystallized unchanged from concentrated hydrochloric acid: the NH3 is so bound to the Co centers that it does not dissociate to allow its protonation.
In contrast, labile metal ammine complexes, such as Cl2, react with acids, reflecting the lability of the Ni–NH3 bonds. Upon heating, hexamminecobalt begins to lose some of its ammine ligands producing a stronger oxidant; the chloride ions in Cl3 can be exchanged with a variety of other anions such as nitrate, iodide, sulfamate to afford the corresponding X3 derivative. Such salts display varying degrees of water solubility; the chloride ion can be exchanged with more complex anions such as the hexathiocyanatochromate, yielding a pink compound with formula, or the ferricyanide ion. Cl3 is prepared by treating cobalt chloride with ammonia and ammonium chloride followed by oxidation. Oxidants include hydrogen oxygen in the presence of charcoal catalyst; this salt appears to have been first reported by Fremy. The acetate salt can be prepared by aerobic oxidation of cobalt acetate, ammonium acetate, ammonia in methanol; the acetate salt is water-soluble to the level of 1.9 M, versus 0.26 M for the trichloride.
3+ is a component of some structural biology methods, to help solve their structures by X-ray crystallography or by nuclear magnetic resonance. In the biological system, the counterions would more be Mg2+, but the heavy atoms of cobalt provide anomalous scattering to solve the phase problem and produce an electron-density map of the structure.3+ is an unusual example of a water-soluble trivalent metal complex and is of utility for charge-shielding applications such as the stabilization of negatively charged complexes e.g. of interactions with and between nucleic acids
Electric dipole moment
The electric dipole moment is a measure of the separation of positive and negative electrical charges within a system, that is, a measure of the system's overall polarity. The SI units for electric dipole moment are coulomb-meter. Theoretically, an electric dipole is defined by the first-order term of the multipole expansion; this is unrealistic. However, because the charge separation is small compared to everyday lengths, the error introduced by treating real dipoles like they are theoretically perfect is negligible; the dipole's direction points from the negative charge towards the positive charge. In physics the dimensions of a massive object can be ignored and can be treated as a pointlike object, i.e. a point particle. Point particles with electric charge are referred to as point charges. Two point charges, one with charge +q and the other one with charge −q separated by a distance d, constitute an electric dipole. For this case, the electric dipole moment has a magnitude p = q d and is directed from the negative charge to the positive one.
Some authors may split d in half and use s = d/2 since this quantity is the distance between either charge and the center of the dipole, leading to a factor of two in the definition. A stronger mathematical definition is to use vector algebra, since a quantity with magnitude and direction, like the dipole moment of two point charges, can be expressed in vector form p = q d where d is the displacement vector pointing from the negative charge to the positive charge; the electric dipole moment vector p points from the negative charge to the positive charge. An idealization of this two-charge system is the electrical point dipole consisting of two charges only infinitesimally separated, but with a finite p; this quantity is used in the definition of polarization density. An object with an electric dipole moment is subject to a torque τ when placed in an external electric field; the torque tends to align the dipole with the field. A dipole aligned parallel to an electric field has lower potential energy than a dipole making some angle with it.
For a spatially uniform electric field E, the torque is given by: τ = p × E, where p is the dipole moment, the symbol "×" refers to the vector cross product. The field vector and the dipole vector define a plane, the torque is directed normal to that plane with the direction given by the right-hand rule. A dipole oriented co- or anti-parallel to the direction in which a non-uniform electric field is increasing will experience a torque, as well as a force in the direction of its dipole moment, it can be shown that this force will always be parallel to the dipole moment regardless of co- or anti-parallel orientation of the dipole. More for a continuous distribution of charge confined to a volume V, the corresponding expression for the dipole moment is: p = ∫ V ρ d 3 r 0, where r locates the point of observation and d3r0 denotes an elementary volume in V. For an array of point charges, the charge density becomes a sum of Dirac delta functions: ρ = ∑ i = 1 N q i δ, where each ri is a vector from some reference point to the charge qi.
Substitution into the above integration formula provides: p = ∑ i = 1 N q i ∫ V δ d 3 r 0 = ∑ i = 1 N q i. This expression is equivalent to the previous expression in the case of charge neutrality and N = 2. For two opposite charges, denoting the location of the positive charge of the pair as r+ and the location of the negative charge as r−: p = q 1 + q 2 ( r
The octet rule is a chemical rule of thumb that reflects observation that atoms of main-group elements tend to bond in such a way that each atom has eight electrons in its valence shell, giving it the same electron configuration as a noble gas. The rule is applicable to carbon, nitrogen and the halogens, but to metals such as sodium or magnesium; the valence electrons can be counted using a Lewis electron dot diagram as shown at the right for carbon dioxide. The electrons shared by the two atoms in a covalent bond are counted once for each atom. In carbon dioxide each oxygen shares four electrons with the central carbon, two from the oxygen itself and two from the carbon. All four of these electrons are counted in both the oxygen octet. Ionic bonding is common between pairs of atoms, where one of the pair is a metal of low electronegativity and the second a nonmetal of high electronegativity. A chlorine atom has seven electrons in its outer electron shell, the first and second shells being filled with two and eight electrons respectively.
The first electron affinity of chlorine is -328.8 kJ per mole of chlorine atoms. Adding a second electron to chlorine requires energy, energy that cannot be recovered by the formation of a chemical bond; the result is that chlorine will often form a compound in which it has eight electrons in its outer shell. A sodium atom has a single electron in its outermost electron shell, the first and second shells again being full with two and eight electrons respectively. To remove this outer electron requires only the first ionization energy, +495.8 kJ per mole of sodium atoms, a small amount of energy. By contrast, the second electron resides in the deeper second electron shell, the second ionization energy required for its removal is much larger: +4562.4 kJ per mole. Thus sodium will, in most cases, form a compound in which it has lost a single electron and have a full outer shell of eight electrons, or octet; the energy required to transfer an electron from a sodium atom to a chlorine atom is small: +495.8 − 328.8 = +167 kJ mol−1.
This energy is offset by the lattice energy of sodium chloride: −787.3 kJ mol−1. This completes the explanation of the octet rule in this case. In the late 19th century it was known that coordination compounds were formed by the combination of atoms or molecules in such a manner that the valencies of the atoms involved became satisfied. In 1893, Alfred Werner showed that the number of atoms or groups associated with a central atom is 4 or 6. In 1904 Richard Abegg was one of the first to extend the concept of coordination number to a concept of valence in which he distinguished atoms as electron donors or acceptors, leading to positive and negative valence states that resemble the modern concept of oxidation states. Abegg noted that the difference between the maximum positive and negative valences of an element under his model is eight. In 1916, Gilbert N. Lewis referred to this insight as Abegg's rule and used it to help formulate his cubical atom model and the "rule of eight", which began to distinguish between valence and valence electrons.
In 1919 Irving Langmuir refined these concepts further and renamed them the "cubical octet atom" and "octet theory". The "octet theory" evolved into what is now known as the "octet rule". Walther Kossel and Gilbert N. Lewis saw that noble gases did not have the tendency of taking part in chemical reactions under ordinary conditions. On the basis of this observation they concluded that atoms of noble gases are stable and on the basis of this conclusion they proposed a theory of valency known as "Electronic Theory of valency" in 1916: During the formation of a chemical bond, atoms combine together by gaining, losing or sharing electrons in such a way that they acquire nearest noble gas configuration; the quantum theory of the atom explains the eight electrons as a closed shell with an s2p6 electron configuration. A closed-shell configuration is one in which low-lying energy levels are full and higher energy levels are empty. For example, the neon atom ground state has an empty n = 3 shell. According to the octet rule, the atoms before and after neon in the periodic table, tend to attain a similar configuration by gaining, losing, or sharing electrons.
The argon atom has an analogous 3s2 3p6 configuration. There is an empty 3d level, but it is at higher energy than 3s and 3p, so that 3s2 3p6 is still considered a closed shell for chemical purposes; the atoms before and after argon tend to attain this configuration in compounds. There are, some hypervalent molecules in which the 3d level may play a part in the bonding, although this is controversial. For helium there is no 1p level according to the quantum theory, so that 1s2 is a closed shell with no p electrons; the atoms before and after helium follow a duet rule and tend to have the same 1s2 configuration as helium. The octet rule is only applicable to main group elements and there are many molecules that do not obey the octet rule; these molecules can be divided into two types: unstable intermediates that react so as to attain stability, stable molecules that follow other electron counting rules. Although stable odd-electron molecules and hypervalent molecules are taught as v
In chemistry, polarity is a separation of electric charge leading to a molecule or its chemical groups having an electric dipole moment, with a negatively charged end and a positively charged end. Polar molecules must contain polar bonds due to a difference in electronegativity between the bonded atoms. A polar molecule with two or more polar bonds must have a geometry, asymmetric in at least one direction, so that the bond dipoles do not cancel each other. Polar molecules interact through dipole–dipole intermolecular forces and hydrogen bonds. Polarity underlies a number of physical properties including surface tension and melting and boiling points. Not all atoms attract electrons with the same force; the amount of "pull" an atom exerts on its electrons is called its electronegativity. Atoms with high electronegativities – such as fluorine and nitrogen – exert a greater pull on electrons than atoms with lower electronegativities such as alkali metals and alkaline earth metals. In a bond, this leads to unequal sharing of electrons between the atoms, as electrons will be drawn closer to the atom with the higher electronegativity.
Because electrons have a negative charge, the unequal sharing of electrons within a bond leads to the formation of an electric dipole: a separation of positive and negative electric charge. Because the amount of charge separated in such dipoles is smaller than a fundamental charge, they are called partial charges, denoted as δ+ and δ−; these symbols were introduced by Sir Christopher Ingold and Dr. Edith Hilda Ingold in 1926; the bond dipole moment is calculated by multiplying the amount of charge separated and the distance between the charges. These dipoles within molecules can interact with dipoles in other molecules, creating dipole-dipole intermolecular forces. Bonds can fall between one of two extremes – being nonpolar or polar. A nonpolar bond occurs when the electronegativities are identical and therefore possess a difference of zero. A polar bond is more called an ionic bond, occurs when the difference between electronegativities is large enough that one atom takes an electron from the other.
The terms "polar" and "nonpolar" are applied to covalent bonds, that is, bonds where the polarity is not complete. To determine the polarity of a covalent bond using numerical means, the difference between the electronegativity of the atoms is used. Bond polarity is divided into three groups that are loosely based on the difference in electronegativity between the two bonded atoms. According to the Pauling scale: Nonpolar bonds occur when the difference in electronegativity between the two atoms is less than 0.5 Polar bonds occur when the difference in electronegativity between the two atoms is between 0.5 and 2.0 Ionic bonds occur when the difference in electronegativity between the two atoms is greater than 2.0Pauling based this classification scheme on the partial ionic character of a bond, an approximate function of the difference in electronegativity between the two bonded atoms. He estimated that a difference of 1.7 corresponds to 50% ionic character, so that a greater difference corresponds to a bond, predominantly ionic.
As a quantum-mechanical description, Pauling proposed that the wave function for a polar molecule AB is a linear combination of wave functions for covalent and ionic molecules: ψ = aψ + bψ. The amount of covalent and ionic character depends on the values of the squared coefficients a2 and b2. While the molecules can be described as "polar covalent", "nonpolar covalent", or "ionic", this is a relative term, with one molecule being more polar or more nonpolar than another. However, the following properties are typical of such molecules. A molecule is composed of one or more chemical bonds between molecular orbitals of different atoms. A molecule may be polar either as a result of polar bonds due to differences in electronegativity as described above, or as a result of an asymmetric arrangement of nonpolar covalent bonds and non-bonding pairs of electrons known as a full molecular orbital. A polar molecule has a net dipole as a result of the opposing charges from polar bonds arranged asymmetrically.
Water is an example of a polar molecule since it has a slight positive charge on one side and a slight negative charge on the other. The dipoles do not cancel out resulting in a net dipole. Due to the polar nature of the water molecule itself, polar molecules are able to dissolve in water. Other examples include sugars, which have many polar oxygen–hydrogen groups and are overall polar. If the bond dipole moments of the molecule do not cancel, the molecule is polar. For example, the water molecule contains two polar O−H bonds in a bent geometry; the bond dipole moments do not cancel, so that the molecule forms a molecular dipole with its negative pole at the oxygen and its positive pole midway between the two hydrogen atoms. In the figure each bond joins the central O atom with a negative charge to an H atom with a positive charge; the hydrogen fluoride, HF, molecule is polar by virtue of polar covalent bonds – in the covalent bond electrons are displaced toward the more electronegative fluorine atom.
Ammonia, NH3, molecule. The molecule has two lone electrons in an orbital, that points towards the fourth apex of the approximate tetrahedron; this orbital is not participating in covalent bonding.
An amine oxide known as amine-N-oxide and N-oxide, is a chemical compound that contains the functional group R3N+−O−, an N−O coordinate covalent bond with three additional hydrogen and/or hydrocarbon side chains attached to N. Sometimes it is written as R3N→O or, wrongly, as R3N=O. In the strict sense, the term amine oxide applies only to oxides of tertiary amines. Sometimes it is used for the analogous derivatives of primary and secondary amines. Examples of amine oxides include pyridine-N-oxide, a water-soluble crystalline solid with melting point 62–67 °C, N-methylmorpholine N-oxide, an oxidant. Amine oxides are surfactants used in consumer products such as shampoos, conditioners and hard surface cleaners. Alkyl dimethyl amine oxide is the most commercially used amine oxide, they are considered a high production volume class of compounds in more than one member country of the Organisation for Economic Co-operation and Development. In North America, more than 95% of amine oxides are used in home cleaning products.
They serve as stabilizers, emollients and conditioners with active concentrations in the range of 0.1–10%. The remainder is used in personal care, commercial products and for unique patented uses such as photography. Amine oxides are used as protecting group as chemical intermediates. Long-chain alkyl amine oxides are used as amphoteric surfactants and foam stabilizers. Amine oxides are polar molecules and have a polarity close to that of quaternary ammonium salts. Small amine oxides are hydrophilic and have an excellent water solubility and a poor solubility in most organic solvents. Amine oxides are weak bases with a pKb of around 4.5 that form R3N+−OH, cationic hydroxylamines, upon protonation at a pH below their pKb. Amine oxides are prepared by oxidation of tertiary amines or pyridine analogs with hydrogen peroxide, Caro's acid or peracids like mCPBA in N-oxidation. Amine oxides exhibit many kinds of reactions. Pyrolytic elimination. Amine oxides, when heated to 150–200 °C eliminate a hydroxylamine, resulting in an alkene.
This pyrolytic syn-elimination reaction is known under the name Cope reaction. The mechanism is similar to that of the Hofmann elimination. Reduction to amines. Amine oxides are converted to the parent amine by common reduction reagents including lithium aluminium hydride, sodium borohydride, catalytic reduction, zinc / acetic acid, iron / acetic acid. Pyridine N-oxides can be deoxygenated by phosphorus oxychloride Sacrificial catalysis. Oxidants can be regenerated by reduction of N-oxides, as in the case of regeneration of osmium tetroxide by N-methylmorpholine oxide. O-alkylation. Pyridine N-oxides react with alkyl halides to the O-alkylated product Bis-ter-pyridine derivatives adsorbed on silver surfaces are discussed to react with oxygen to bis-ter-pyridine-N-oxide; this reaction can be followed by video-scanning tunneling microscopy with sub-molecular resolution. In the Meisenheimer rearrangement certain N-oxides R1R2R3N+O− rearrange to hydroxylamines R2R3N−O−R1in a 1,2-rearrangement: or a 2,3-rearrangement: In the Polonovski reaction a tertiary N-oxide is cleaved by acetic acid anhydride to the corresponding acetamide and aldehyde: Amine oxides are common metabolites of medication and psychoactive drugs.
Examples include nicotine and morphine. Amine oxides of anti-cancer drugs have been developed as prodrugs that are metabolized in the oxygen-deficient cancer tissue to the active drug. Amine oxides are not known to dermal sensitizers or cause reproductive toxicity, they are metabolized and excreted if ingested. Chronic ingestion by rabbits found lower body weight and lenticular opacities at a lowest observed adverse effect levels in the range of 87–150 mg AO/kw bw/day. Tests of human skin exposure have found. Eye irritation due to amine oxides and other surfactants is moderate and temporary with no lasting effects. Amine oxides with an average chain length of 12.6 have been measured to be water-soluble at ~410 g L−1. They are considered to have low bioaccumulation potential in aquatic species based on log Kow data from chain lengths less than C14. Levels of AO in untreated influent were found to be 2.3–27.8 ug L−1, while in effluent they were found to be 0.4–2.91 ug L−1. The highest effluent concentrations were found in oxidation ditch and trickling filter treatment plants.
On average, over 96% removal has been found with secondary activated sludge treatment. Acute toxicity in fish, as indicated by 96h LC50 tests, is in the range of 1,000–3,000 ug L−1 for carbon chain lengths less than C14. LC50 values for chain lengths greater than C14 range from 600 to 1400 ug L−1. Chronic toxicity data for fish is 420 ug/L; when normalized to C12.9, the NOEC is 310 ug L − 1 for hatchability. Functional group Amine, NR3 Hydroxylamine, NR2OH Phosphine oxide, PR3=O Sulfoxide, R2S=O Azoxy, RN=N+R RN=N+RO− Aminoxyl group Radicals with the general structure R2N–O• Chemistry of amine oxides Surfactants and uses The amine oxides homepage Nomenclature of nitrogen compounds IUPAC definition