In optics, the refractive index or index of refraction of a material is a dimensionless number that describes how fast light propagates through the material. It is defined as n = c v, where c is the speed of light in vacuum and v is the phase velocity of light in the medium. For example, the refractive index of water is 1.333, meaning that light travels 1.333 times as fast in vacuum as in water. The refractive index determines how much the path of light is bent, or refracted, when entering a material; this is described by Snell's law of refraction, n1 sinθ1 = n2 sinθ2, where θ1 and θ2 are the angles of incidence and refraction of a ray crossing the interface between two media with refractive indices n1 and n2. The refractive indices determine the amount of light, reflected when reaching the interface, as well as the critical angle for total internal reflection and Brewster's angle; the refractive index can be seen as the factor by which the speed and the wavelength of the radiation are reduced with respect to their vacuum values: the speed of light in a medium is v = c/n, the wavelength in that medium is λ = λ0/n, where λ0 is the wavelength of that light in vacuum.
This implies that vacuum has a refractive index of 1, that the frequency of the wave is not affected by the refractive index. As a result, the energy of the photon, therefore the perceived color of the refracted light to a human eye which depends on photon energy, is not affected by the refraction or the refractive index of the medium. While the refractive index affects wavelength, it depends on photon frequency and energy so the resulting difference in the bending angle causes white light to split into its constituent colors; this is called dispersion. It can be observed in prisms and rainbows, chromatic aberration in lenses. Light propagation in absorbing materials can be described using a complex-valued refractive index; the imaginary part handles the attenuation, while the real part accounts for refraction. The concept of refractive index applies within the full electromagnetic spectrum, from X-rays to radio waves, it can be applied to wave phenomena such as sound. In this case the speed of sound is used instead of that of light, a reference medium other than vacuum must be chosen.
The refractive index n of an optical medium is defined as the ratio of the speed of light in vacuum, c = 299792458 m/s, the phase velocity v of light in the medium, n = c v. The phase velocity is the speed at which the crests or the phase of the wave moves, which may be different from the group velocity, the speed at which the pulse of light or the envelope of the wave moves; the definition above is sometimes referred to as the absolute refractive index or the absolute index of refraction to distinguish it from definitions where the speed of light in other reference media than vacuum is used. Air at a standardized pressure and temperature has been common as a reference medium. Thomas Young was the person who first used, invented, the name "index of refraction", in 1807. At the same time he changed this value of refractive power into a single number, instead of the traditional ratio of two numbers; the ratio had the disadvantage of different appearances. Newton, who called it the "proportion of the sines of incidence and refraction", wrote it as a ratio of two numbers, like "529 to 396".
Hauksbee, who called it the "ratio of refraction", wrote it as a ratio with a fixed numerator, like "10000 to 7451.9". Hutton wrote it as a ratio with a fixed denominator, like 1.3358 to 1. Young did not use a symbol for the index of refraction, in 1807. In the next years, others started using different symbols: n, m, µ; the symbol n prevailed. For visible light most transparent media have refractive indices between 1 and 2. A few examples are given in the adjacent table; these values are measured at the yellow doublet D-line of sodium, with a wavelength of 589 nanometers, as is conventionally done. Gases at atmospheric pressure have refractive indices close to 1 because of their low density. All solids and liquids have refractive indices above 1.3, with aerogel as the clear exception. Aerogel is a low density solid that can be produced with refractive index in the range from 1.002 to 1.265. Moissanite lies at the other end of the range with a refractive index as high as 2.65. Most plastics have refractive indices in the range from 1.3 to 1.7, but some high-refractive-index polymers can have values as high as 1.76.
For infrared light refractive indices can be higher. Germanium is transparent in the wavelength region from 2 to 14 µm and has a refractive index of about 4. A type of new materials, called topological insulator, was found holding higher refractive index of up to 6 in near to mid infrared frequency range. Moreover, topological insulator material are transparent; these excellent properties make them a type of significant materials for infrared optics. According to the theory of relativity, no information can travel faster than the speed of light in vacuum, but this does not mean that the refractive index cannot be lower than 1; the refractive index measures the phase velocity of light. The phase velocity is the speed at which the crests of the wave move and can be faster than the speed of light in vacuum, thereby give a refractive index below 1; this can occur close to resonance frequencies, for absorbing media, in plasmas, for X-rays. In the X-ray regime the refractive indices are
Antimony oxide is the inorganic compound with the formula Sb2O3. It is the most important commercial compound of antimony, it is found in nature as the minerals senarmontite. Like most polymeric oxides, Sb2O3 dissolves in aqueous solutions with hydrolysis. Global production of antimony oxide in 2012 was 130,000 tonnes, an increase from 112,600 tonnes in 2002. China produces the largest share followed by US/Mexico, Europe and South Africa and other countries; as of 2010, antimony oxide was produced at four sites in EU27. It is produced via two routes, re-volatilizing of crude antimony oxide and by oxidation of antimony metal. Oxidation of antimony metal dominates in Europe. Several processes for the production of crude antimony oxide or metallic antimony from virgin material; the choice of process depends on the composition of other factors. Typical steps include mining and grinding of ore, sometimes followed by froth flotation and separation of the metal using pyrometallurgical processes or in a few cases by hydrometallurgical processes.
These steps do not take place in the EU but closer to the mining location. Step 1) Crude stibnite is oxidized to crude antimony oxide using furnaces operating at 500 to 1,000 °C; the reaction is the following: 2 Sb2S3 + 9 O2 → 2 Sb2O3 + 6 SO2Step 2) The crude antimony oxide is purified by sublimation. Antimony metal is oxidized to antimony oxide in furnaces; the reaction is exothermic. Antimony oxide is recovered in bag filters; the size of the formed particles is controlled by process conditions in gas flow. The reaction can be schematically described by: 4 Sb + 3 O2 → 2 Sb2O3 Antimony oxide is an amphoteric oxide, it dissolves in aqueous sodium hydroxide solution to give the meta-antimonite NaSbO2, which can be isolated as the trihydrate. Antimony oxide dissolves in concentrated mineral acids to give the corresponding salts, which hydrolyzes upon dilution with water. With nitric acid, the trioxide is oxidized to antimony oxide; when heated with carbon, the oxide is reduced to antimony metal. With other reducing agents such as sodium borohydride or lithium aluminium hydride, the unstable and toxic gas stibine is produced.
When heated with potassium bitartrate, a complex salt potassium antimony tartrate, KSb2•C4H2O6 is formed. The structure of Sb2O3 depends on the temperature of the sample. Dimeric Sb4O6 is the high temperature gas. Sb4O6 molecules are bicyclic cages, phosphorus trioxide; the cage structure is retained in a solid. The Sb-O distance is 197.7 pm and the O-Sb-O angle of 95.6°. This form exists in nature as the mineral senarmontite. Above 606 °C, the more stable form is orthorhombic, consisting of pairs of -Sb-O-Sb-O- chains that are linked by oxide bridges between the Sb centers; this form exists in nature as the mineral valentinite. The annual consumption of antimony oxide in the United States and Europe is 10,000 and 25,000 tonnes, respectively; the main application is as flame retardant synergist in combination with halogenated materials. The combination of the halides and the antimony is key to the flame-retardant action for polymers, helping to form less flammable chars; such flame retardants are found in electrical apparatuses, textiles and coatings.
Other applications: Antimony oxide is an opacifying agent for glasses and enamels. Some specialty pigments contain antimony. Antimony oxide is a useful catalyst in the production of polyethylene terephthalate and the vulcanization of rubber. Antimony oxide has suspected carcinogenic potential for humans, its TLV is 0.5 mg/m3, as for most antimony compounds. No other human health hazards were identified for antimony oxide, no risks to human health and the environment were identified from the production and use of antimony trioxide in daily life. Institut national de recherche et de sécurité, Fiche toxicologique nº 198: Trioxyde de diantimoine, 1992; the Oxide Handbook, G. V. Samsonov, 1981, 2nd ed. IFI/Plenum, ISBN 0-306-65177-7 International Antimony Association International Chemical Safety Card 0012 Antimony Market And Price Société industrielle et chimique de l'Aisne
Enthalpy of vaporization
The enthalpy of vaporization known as the heat of vaporization or heat of evaporation, is the amount of energy that must be added to a liquid substance, to transform a quantity of that substance into a gas. The enthalpy of vaporization is a function of the pressure; the enthalpy of vaporization is quoted for the normal boiling temperature of the substance. The heat of vaporization is temperature-dependent, though a constant heat of vaporization can be assumed for small temperature ranges and for reduced temperature T r ≪ 1; the heat of vaporization diminishes with increasing temperature and it vanishes at a certain point called the critical temperature. Above the critical temperature, the liquid and vapor phases are indistinguishable, the substance is called a supercritical fluid. Values are quoted in J/mol or kJ/mol, although kJ/kg or J/g, older units like kcal/mol, cal/g and Btu/lb are sometimes still used, among others; the enthalpy of condensation is by definition equal to the enthalpy of vaporization with the opposite sign: enthalpy changes of vaporization are always positive, whereas enthalpy changes of condensation are always negative.
The enthalpy of vaporization can be written as Δ H v a p = Δ U v a p + p Δ V It is equal to the increased internal energy of the vapor phase compared with the liquid phase, plus the work done against ambient pressure. The increase in the internal energy can be viewed as the energy required to overcome the intermolecular interactions in the liquid. Hence helium has a low enthalpy of vaporization, 0.0845 kJ/mol, as the van der Waals forces between helium atoms are weak. On the other hand, the molecules in liquid water are held together by strong hydrogen bonds, its enthalpy of vaporization, 40.65 kJ/mol, is more than five times the energy required to heat the same quantity of water from 0 °C to 100 °C. Care must be taken, when using enthalpies of vaporization to measure the strength of intermolecular forces, as these forces may persist to an extent in the gas phase, so the calculated value of the bond strength will be too low; this is true of metals, which form covalently bonded molecules in the gas phase: in these cases, the enthalpy of atomization must be used to obtain a true value of the bond energy.
An alternative description is to view the enthalpy of condensation as the heat which must be released to the surroundings to compensate for the drop in entropy when a gas condenses to a liquid. As the liquid and gas are in equilibrium at the boiling point, ΔvG = 0, which leads to: Δ v S = S g a s − S l i q u i d = Δ v H / T b As neither entropy nor enthalpy vary with temperature, it is normal to use the tabulated standard values without any correction for the difference in temperature from 298 K. A correction must be made if the pressure is different from 100 kPa, as the entropy of a gas is proportional to its pressure: the entropies of liquids vary little with pressure, as the compressibility of a liquid is small; these two definitions are equivalent: the boiling point is the temperature at which the increased entropy of the gas phase overcomes the intermolecular forces. As a given quantity of matter always has a higher entropy in the gas phase than in a condensed phase, from Δ G = Δ H − T Δ S,the Gibbs free energy change falls with increasing temperature: gases are favored at higher temperatures, as is observed in practice.
Estimation of the enthalpy of vaporization of electrolyte solutions can be carried out using equations based on the chemical thermodynamic models, such as Pitzer model or TCPC model. The vaporization of metals is a key step in metal vapor synthesis, which exploits the increased reactivity of metal atoms or small particles relative to the bulk elements. Enthalpies of vaporization of common substances, measured at their respective standard boiling points: Enthalpy of fusion Enthalpy of sublimation Joback method CODATA Key Values for Thermodynamics Gmelin, Leopold. Gmelin-Handbuch der anorganischen Chemie / 08 a. Berlin: Springer. Pp. 116–117. ISBN 978-3-540-93516-2. NIST Chemistry WebBook Young, Francis W. Sears, Mark W. Zemansky, Hugh D.. University physics. Read
In molecular geometry, bond length or bond distance is the average distance between nuclei of two bonded atoms in a molecule. It is a transferable property of a bond between atoms of fixed types independent of the rest of the molecule. Bond length is related to bond order: when more electrons participate in bond formation the bond is shorter. Bond length is inversely related to bond strength and the bond dissociation energy: all other factors being equal, a stronger bond will be shorter. In a bond between two identical atoms, half the bond distance is equal to the covalent radius. Bond lengths are measured in the solid phase by means of X-ray diffraction, or approximated in the gas phase by microwave spectroscopy. A bond between a given pair of atoms may vary between different molecules. For example, the carbon to hydrogen bonds in methane are different from those in methyl chloride, it is however possible to make generalizations. A table with experimental single bonds for carbon to other elements is given below.
Bond lengths are given in picometers. By approximation the bond distance between two different atoms is the sum of the individual covalent radii; as a general trend, bond distances decrease across the row in the periodic table and increase down a group. This trend is identical to that of the atomic radius; the bond length between two atoms in a molecule depends not only on the atoms but on such factors as the orbital hybridization and the electronic and steric nature of the substituents. The carbon–carbon bond length in diamond is 154 pm, the largest bond length that exists for ordinary carbon covalent bonds. Since one atomic unit of length is 52.9177 pm, the C–C bond length is 2.91 atomic units, or three Bohr radii long. Unusually long bond lengths do exist. In one compound, tricyclobutabenzene, a bond length of 160 pm is reported; the current record holder is another cyclobutabenzene with length 174 pm based on X-ray crystallography. In this type of compound the cyclobutane ring would force 90° angles on the carbon atoms connected to the benzene ring where they ordinarily have angles of 120°.
The existence of a long C–C bond length of up to 290 pm is claimed in a dimer of two tetracyanoethylene dianions, although this concerns a 2-electron-4-center bond. This type of bonding has been observed in neutral phenalenyl dimers; the bond lengths of these so-called "pancake bonds" are up to 305 pm. Shorter than average C–C bond distances are possible: alkenes and alkynes have bond lengths of 133 and 120 pm due to increased s-character of the sigma bond. In benzene all bonds have the same length: 139 pm. Carbon–carbon single bonds increased s-character is notable in the central bond of diacetylene and that of a certain tetrahedrane dimer. In propionitrile the cyano group withdraws electrons resulting in a reduced bond length. Squeezing a C–C bond is possible by application of strain. An unusual organic compound exists called In-methylcyclophane with a short bond distance of 147 pm for the methyl group being squeezed between a triptycene and a phenyl group. In an in silico experiment a bond distance of 136 pm was estimated for neopentane locked up in fullerene.
The smallest theoretical C–C single bond obtained in this study is 131 pm for a hypothetical tetrahedrane derivative. The same study estimated that stretching or squeezing the C–C bond in an ethane molecule by 5 pm required 2.8 or 3.5 kJ/mol, respectively. Stretching or squeezing the same bond by 15 pm required an estimated 21.9 or 37.7 kJ/mol. Bond length tutorial
Critical point (thermodynamics)
In thermodynamics, a critical point is the end point of a phase equilibrium curve. The most prominent example is the liquid-vapor critical point, the end point of the pressure-temperature curve that designates conditions under which a liquid and its vapor can coexist. At higher temperatures, the gas cannot be liquefied by pressure alone. At the critical point, defined by a critical temperature Tc and a critical pressure pc, phase boundaries vanish. Other examples include the liquid–liquid critical points in mixtures. For simplicity and clarity, the generic notion of critical point is best introduced by discussing a specific example, the liquid-vapor critical point; this was the first critical point to be discovered, it is still the best known and most studied one. The figure to the right shows the schematic PT diagram of a pure substance; the known phases solid and vapor are separated by phase boundaries, i.e. pressure-temperature combinations where two phases can coexist. At the triple point, all three phases can coexist.
However, the liquid-vapor boundary terminates in an endpoint at some critical temperature Tc and critical pressure pc. This is the critical point. In water, the critical point occurs at 22.064 MPa. In the vicinity of the critical point, the physical properties of the liquid and the vapor change with both phases becoming more similar. For instance, liquid water under normal conditions is nearly incompressible, has a low thermal expansion coefficient, has a high dielectric constant, is an excellent solvent for electrolytes. Near the critical point, all these properties change into the exact opposite: water becomes compressible, expandable, a poor dielectric, a bad solvent for electrolytes, prefers to mix with nonpolar gases and organic molecules. At the critical point, only one phase exists; the heat of vaporization is zero. There is a stationary inflection point in the constant-temperature line on a PV diagram; this means that at the critical point: T = 0 T = 0 Above the critical point there exists a state of matter, continuously connected with both the liquid and the gaseous state.
It is called supercritical fluid. The common textbook knowledge that all distinction between liquid and vapor disappears beyond the critical point has been challenged by Fisher and Widom who identified a p,T-line that separates states with different asymptotic statistical properties; the existence of a critical point was first discovered by Charles Cagniard de la Tour in 1822 and named by Dmitri Mendeleev in 1860 and Thomas Andrews in 1869. Cagniard showed that CO2 could be liquefied at 31 °C at a pressure of 73 atm, but not at a higher temperature under pressures as high as 3,000 atm. Solving the above condition T = 0 for the van der Waals equation, one can compute the critical point as T c = 8 a 27 R b, V c = 3 n b, p c = a 27 b 2. However, the van der Waals equation, based on a mean field theory, does not hold near the critical point. In particular, it predicts wrong scaling laws. To analyse properties of fluids near the critical point, reduced state variables are sometimes defined relative to the critical properties T r = T T c, p r = p p c, V r = V R T c / p c.
The principle of corresponding states indicates that substances at equal reduced pressures and temperatures have equal reduced volumes. This relationship is true for many substances, but becomes inaccurate for large values of pr. For some gases, there is an additional correction factor, called Newton's correction, added to the critical temperature and critical pressure calculated in this manner; these vary with the pressure range of interest. The liquid–liquid critical point of a solution, which occurs at the critical solution temperature, occurs at the limit of the two-phase region of the phase diagram. In other words, it is the point at which an infinitesimal change in some thermodynamic variable will lead to separation of the mixture into two distinct liquid phases, as shown in the polymer–solvent phase diagram to the right. Two types of liquid–liquid critical points are the upper critical solution temperature, t
Mass spectrometry is an analytical technique that ionizes chemical species and sorts the ions based on their mass-to-charge ratio. In simpler terms, a mass spectrum measures the masses within a sample. Mass spectrometry is used in many different fields and is applied to pure samples as well as complex mixtures. A mass spectrum is a plot of the ion signal as a function of the mass-to-charge ratio; these spectra are used to determine the elemental or isotopic signature of a sample, the masses of particles and of molecules, to elucidate the chemical structures of molecules and other chemical compounds. In a typical MS procedure, a sample, which may be solid, liquid, or gas, is ionized, for example by bombarding it with electrons; this may cause some of the sample's molecules to break into charged fragments. These ions are separated according to their mass-to-charge ratio by accelerating them and subjecting them to an electric or magnetic field: ions of the same mass-to-charge ratio will undergo the same amount of deflection.
The ions are detected by a mechanism capable of detecting charged particles, such as an electron multiplier. Results are displayed as spectra of the relative abundance of detected ions as a function of the mass-to-charge ratio; the atoms or molecules in the sample can be identified by correlating known masses to the identified masses or through a characteristic fragmentation pattern. In 1886, Eugen Goldstein observed rays in gas discharges under low pressure that traveled away from the anode and through channels in a perforated cathode, opposite to the direction of negatively charged cathode rays. Goldstein called these positively charged anode rays "Kanalstrahlen". Wilhelm Wien found that strong electric or magnetic fields deflected the canal rays and, in 1899, constructed a device with perpendicular electric and magnetic fields that separated the positive rays according to their charge-to-mass ratio. Wien found. English scientist J. J. Thomson improved on the work of Wien by reducing the pressure to create the mass spectrograph.
The word spectrograph had become part of the international scientific vocabulary by 1884. Early spectrometry devices that measured the mass-to-charge ratio of ions were called mass spectrographs which consisted of instruments that recorded a spectrum of mass values on a photographic plate. A mass spectroscope is similar to a mass spectrograph except that the beam of ions is directed onto a phosphor screen. A mass spectroscope configuration was used in early instruments when it was desired that the effects of adjustments be observed. Once the instrument was properly adjusted, a photographic plate was exposed; the term mass spectroscope continued to be used though the direct illumination of a phosphor screen was replaced by indirect measurements with an oscilloscope. The use of the term mass spectroscopy is now discouraged due to the possibility of confusion with light spectroscopy. Mass spectrometry is abbreviated as mass-spec or as MS. Modern techniques of mass spectrometry were devised by Arthur Jeffrey Dempster and F.
W. Aston in 1918 and 1919 respectively. Sector mass spectrometers known as calutrons were developed by Ernest O. Lawrence and used for separating the isotopes of uranium during the Manhattan Project. Calutron mass spectrometers were used for uranium enrichment at the Oak Ridge, Tennessee Y-12 plant established during World War II. In 1989, half of the Nobel Prize in Physics was awarded to Hans Dehmelt and Wolfgang Paul for the development of the ion trap technique in the 1950s and 1960s. In 2002, the Nobel Prize in Chemistry was awarded to John Bennett Fenn for the development of electrospray ionization and Koichi Tanaka for the development of soft laser desorption and their application to the ionization of biological macromolecules proteins. A mass spectrometer consists of three components: an ion source, a mass analyzer, a detector; the ionizer converts a portion of the sample into ions. There is a wide variety of ionization techniques, depending on the phase of the sample and the efficiency of various ionization mechanisms for the unknown species.
An extraction system removes ions from the sample, which are targeted through the mass analyzer and into the detector. The differences in masses of the fragments allows the mass analyzer to sort the ions by their mass-to-charge ratio; the detector measures the value of an indicator quantity and thus provides data for calculating the abundances of each ion present. Some detectors give spatial information, e.g. a multichannel plate. The following example describes the operation of a spectrometer mass analyzer, of the sector type. Consider a sample of sodium chloride. In the ion source, the sample is ionized into sodium and chloride ions. Sodium atoms and ions are monoisotopic, with a mass of about 23 u. Chloride atoms and ions come in two isotopes with masses of 35 u and 37 u; the analyzer part of the spectrometer contains electric and magnetic fields, which exert forces on ions traveling through these fields. The speed of a charged particle may be increased or decreased while passing through the electric field, its direction may be altered by the magnetic field.
The magnitude of the deflection of the moving ion's trajectory depends on its mass-to-charge ratio. L
Nuclear magnetic resonance spectroscopy
Nuclear magnetic resonance spectroscopy, most known as NMR spectroscopy or magnetic resonance spectroscopy, is a spectroscopic technique to observe local magnetic fields around atomic nuclei. The sample is placed in a magnetic field and the NMR signal is produced by excitation of the nuclei sample with radio waves into nuclear magnetic resonance, detected with sensitive radio receivers; the intramolecular magnetic field around an atom in a molecule changes the resonance frequency, thus giving access to details of the electronic structure of a molecule and its individual functional groups. As the fields are unique or characteristic to individual compounds, in modern organic chemistry practice, NMR spectroscopy is the definitive method to identify monomolecular organic compounds. Biochemists use NMR to identify proteins and other complex molecules. Besides identification, NMR spectroscopy provides detailed information about the structure, reaction state, chemical environment of molecules; the most common types of NMR are proton and carbon-13 NMR spectroscopy, but it is applicable to any kind of sample that contains nuclei possessing spin.
NMR spectra are unique, well-resolved, analytically tractable and highly predictable for small molecules. Different functional groups are distinguishable, identical functional groups with differing neighboring substituents still give distinguishable signals. NMR has replaced traditional wet chemistry tests such as color reagents or typical chromatography for identification. A disadvantage is that a large amount, 2–50 mg, of a purified substance is required, although it may be recovered through a workup. Preferably, the sample should be dissolved in a solvent, because NMR analysis of solids requires a dedicated magic angle spinning machine and may not give well-resolved spectra; the timescale of NMR is long, thus it is not suitable for observing fast phenomena, producing only an averaged spectrum. Although large amounts of impurities do show on an NMR spectrum, better methods exist for detecting impurities, as NMR is inherently not sensitive - though at higher frequencies, sensitivity is higher.
Correlation spectroscopy is a development of ordinary NMR. In two-dimensional NMR, the emission is centered around a single frequency, correlated resonances are observed; this allows identifying the neighboring substituents of the observed functional group, allowing unambiguous identification of the resonances. There are more complex 3D and 4D methods and a variety of methods designed to suppress or amplify particular types of resonances. In nuclear Overhauser effect spectroscopy, the relaxation of the resonances is observed; as NOE depends on the proximity of the nuclei, quantifying the NOE for each nucleus allows for construction of a three-dimensional model of the molecule. NMR spectrometers are expensive. Modern NMR spectrometers have a strong and expensive liquid helium-cooled superconducting magnet, because resolution directly depends on magnetic field strength. Less expensive machines using permanent magnets and lower resolution are available, which still give sufficient performance for certain application such as reaction monitoring and quick checking of samples.
There are benchtop nuclear magnetic resonance spectrometers. NMR can be observed than a millitesla. Low-resolution NMR produces broader peaks which can overlap one another causing issues in resolving complex structures; the use of higher strength magnetic fields result in clear resolution of the peaks and is the standard in industry. The Purcell group at Harvard University and the Bloch group at Stanford University independently developed NMR spectroscopy in the late 1940s and early 1950s. Edward Mills Purcell and Felix Bloch shared the 1952 Nobel Prize in Physics for their discoveries; when placed in a magnetic field, NMR active nuclei absorb electromagnetic radiation at a frequency characteristic of the isotope. The resonant frequency, energy of the radiation absorbed, the intensity of the signal are proportional to the strength of the magnetic field. For example, in a 21 Tesla magnetic field, hydrogen atoms resonate at 900 MHz, it is common to refer to a 21 T magnet as a 900 MHz magnet since hydrogen is the most common nucleus detected, however different nuclei will resonate at different frequencies at this field strength in proportion to their nuclear magnetic moments.
An NMR spectrometer consists of a spinning sample-holder inside a strong magnet, a radio-frequency emitter and a receiver with a probe that goes inside the magnet to surround the sample, optionally gradient coils for diffusion measurements, electronics to control the system. Spinning the sample is necessary to average out diffusional motion, however some experiments call for a stationary sample when solution movement is an important variable. For instance, measurements of diffusion constants are done using a stationary sample with spinning off, flow cells can be used for online analysis of process flows; the vast majority of molecules in a solution are solvent molecules, most regular solvents are hydrocarbons and so contain NMR-active protons. In order to avoid detecting only signals from solvent hydrogen atoms, deuterated solvents are used where 99+% of the protons are replaced with deuterium; the most used deuterated solvent is deuterochloroform, although other solvents may be used depending on the solubility of a sample.
Deuterium oxide and deuterated DMSO (DMSO-d