The boiling point of a substance is the temperature at which the vapor pressure of a liquid equals the pressure surrounding the liquid and the liquid changes into a vapor. The boiling point of a liquid varies depending upon the surrounding environmental pressure. A liquid in a partial vacuum has a lower boiling point than when that liquid is at atmospheric pressure. A liquid at high pressure has a higher boiling point than when that liquid is at atmospheric pressure. For example, water at 93.4 °C at 1,905 metres altitude. For a given pressure, different liquids will boil at different temperatures; the normal boiling point of a liquid is the special case in which the vapor pressure of the liquid equals the defined atmospheric pressure at sea level, 1 atmosphere. At that temperature, the vapor pressure of the liquid becomes sufficient to overcome atmospheric pressure and allow bubbles of vapor to form inside the bulk of the liquid; the standard boiling point has been defined by IUPAC since 1982 as the temperature at which boiling occurs under a pressure of 1 bar.
The heat of vaporization is the energy required to transform a given quantity of a substance from a liquid into a gas at a given pressure. Liquids may change to a vapor at temperatures below their boiling points through the process of evaporation. Evaporation is a surface phenomenon in which molecules located near the liquid's edge, not contained by enough liquid pressure on that side, escape into the surroundings as vapor. On the other hand, boiling is a process in which molecules anywhere in the liquid escape, resulting in the formation of vapor bubbles within the liquid. A saturated liquid contains as much thermal energy. Saturation temperature means boiling point; the saturation temperature is the temperature for a corresponding saturation pressure at which a liquid boils into its vapor phase. The liquid can be said to be saturated with thermal energy. Any addition of thermal energy results in a phase transition. If the pressure in a system remains constant, a vapor at saturation temperature will begin to condense into its liquid phase as thermal energy is removed.
A liquid at saturation temperature and pressure will boil into its vapor phase as additional thermal energy is applied. The boiling point corresponds to the temperature at which the vapor pressure of the liquid equals the surrounding environmental pressure. Thus, the boiling point is dependent on the pressure. Boiling points may be published with respect to the NIST, USA standard pressure of 101.325 kPa, or the IUPAC standard pressure of 100.000 kPa. At higher elevations, where the atmospheric pressure is much lower, the boiling point is lower; the boiling point increases with increased pressure up to the critical point, where the gas and liquid properties become identical. The boiling point cannot be increased beyond the critical point; the boiling point decreases with decreasing pressure until the triple point is reached. The boiling point cannot be reduced below the triple point. If the heat of vaporization and the vapor pressure of a liquid at a certain temperature are known, the boiling point can be calculated by using the Clausius–Clapeyron equation, thus: T B = − 1, where: T B is the boiling point at the pressure of interest, R is the ideal gas constant, P is the vapour pressure of the liquid at the pressure of interest, P 0 is some pressure where the corresponding T 0 is known, Δ H vap is the heat of vaporization of the liquid, T 0 is the boiling temperature, ln is the natural logarithm.
Saturation pressure is the pressure for a corresponding saturation temperature at which a liquid boils into its vapor phase. Saturation pressure and saturation temperature have a direct relationship: as saturation pressure is increased, so is saturation temperature. If the temperature in a system remains constant, vapor at saturation pressure and temperature will begin to condense into its liquid phase as the system pressure is increased. A liquid at saturation pressure and temperature will tend to flash into its vapor phase as system pressure is decreased. There are two conventions regarding the standard boiling point of water: The normal boiling point is 99.97 °C at a pressure of 1 atm. The IUPAC recommended standard boiling point of water at a standard pressure of 100 kPa is 99.61 °C. For comparison, on top of Mount Everest, at 8,848 m elevation, the pressure is about 34 kPa and the boiling point of water is 71 °C; the Celsius temperature scale was defined until 1954 by two points: 0 °C being defined by the wate
Spectroscopy is the study of the interaction between matter and electromagnetic radiation. Spectroscopy originated through the study of visible light dispersed according to its wavelength, by a prism; the concept was expanded to include any interaction with radiative energy as a function of its wavelength or frequency, predominantly in the electromagnetic spectrum, though matter waves and acoustic waves can be considered forms of radiative energy. Spectroscopic data are represented by an emission spectrum, a plot of the response of interest as a function of wavelength or frequency. Spectroscopy in the electromagnetic spectrum, is a fundamental exploratory tool in the fields of physics and astronomy, allowing the composition, physical structure and electronic structure of matter to be investigated at atomic scale, molecular scale, macro scale, over astronomical distances. Important applications arise from biomedical spectroscopy in the areas of tissue analysis and medical imaging. Spectroscopy and spectrography are terms used to refer to the measurement of radiation intensity as a function of wavelength and are used to describe experimental spectroscopic methods.
Spectral measurement devices are referred to as spectrometers, spectrophotometers, spectrographs or spectral analyzers. Daily observations of color can be related to spectroscopy. Neon lighting is a direct application of atomic spectroscopy. Neon and other noble gases have characteristic emission frequencies. Neon lamps use collision of electrons with the gas to excite these emissions. Inks and paints include chemical compounds selected for their spectral characteristics in order to generate specific colors and hues. A encountered molecular spectrum is that of nitrogen dioxide. Gaseous nitrogen dioxide has a characteristic red absorption feature, this gives air polluted with nitrogen dioxide a reddish-brown color. Rayleigh scattering is a spectroscopic scattering phenomenon. Spectroscopic studies were central to the development of quantum mechanics and included Max Planck's explanation of blackbody radiation, Albert Einstein's explanation of the photoelectric effect and Niels Bohr's explanation of atomic structure and spectra.
Spectroscopy is used in physical and analytical chemistry because atoms and molecules have unique spectra. As a result, these spectra can be used to detect and quantify information about the atoms and molecules. Spectroscopy is used in astronomy and remote sensing on Earth. Most research telescopes have spectrographs; the measured spectra are used to determine the chemical composition and physical properties of astronomical objects. One of the central concepts in spectroscopy is its corresponding resonant frequency. Resonances were first characterized in mechanical systems such as pendulums. Mechanical systems that vibrate or oscillate will experience large amplitude oscillations when they are driven at their resonant frequency. A plot of amplitude vs. excitation frequency will have a peak centered at the resonance frequency. This plot is one type of spectrum, with the peak referred to as a spectral line, most spectral lines have a similar appearance. In quantum mechanical systems, the analogous resonance is a coupling of two quantum mechanical stationary states of one system, such as an atom, via an oscillatory source of energy such as a photon.
The coupling of the two states is strongest when the energy of the source matches the energy difference between the two states. The energy of a photon is related to its frequency by E = h ν where h is Planck's constant, so a spectrum of the system response vs. photon frequency will peak at the resonant frequency or energy. Particles such as electrons and neutrons have a comparable relationship, the de Broglie relations, between their kinetic energy and their wavelength and frequency and therefore can excite resonant interactions. Spectra of atoms and molecules consist of a series of spectral lines, each one representing a resonance between two different quantum states; the explanation of these series, the spectral patterns associated with them, were one of the experimental enigmas that drove the development and acceptance of quantum mechanics. The hydrogen spectral series in particular was first explained by the Rutherford-Bohr quantum model of the hydrogen atom. In some cases spectral lines are well separated and distinguishable, but spectral lines can overlap and appear to be a single transition if the density of energy states is high enough.
Named series of lines include the principal, sharp and fundamental series. Spectroscopy is a sufficiently broad field that many sub-disciplines exist, each with numerous implementations of specific spectroscopic techniques; the various implementations and techniques can be classified in several ways. The types of spectroscopy are distinguished by the type of radiative energy involved in the interaction. In many applications, the spectrum is determined by measuring changes in the intensity or frequency of this energy; the types of radiative energy studied include: Electromagnetic radiation was the first source of energy used for spectroscopic studies. Techniques that employ electromagnetic radiation are classified by the wavelength region of the spectrum and include microwave, terahe
Comet Hale–Bopp is a comet, the most observed of the 20th century, one of the brightest seen for many decades. Hale–Bopp was discovered on July 23, 1995, separately by Alan Hale and Thomas Bopp prior to it becoming naked-eye visible on Earth. Although predicting the maximum apparent brightness of new comets with any degree of certainty is difficult, Hale–Bopp met or exceeded most predictions when it passed perihelion on April 1, 1997, it was visible to the naked eye for a record 18 months, twice as long as the previous record holder, the Great Comet of 1811. Accordingly, Hale–Bopp was dubbed the Great Comet of 1997; the comet was discovered independently on July 23, 1995, by two observers, Alan Hale and Thomas Bopp, both in the United States. Hale had spent many hundreds of hours searching for comets without success, was tracking known comets from his driveway in New Mexico when he chanced upon Hale–Bopp just after midnight; the comet had an apparent magnitude of 10.5 and lay near the globular cluster M70 in the constellation of Sagittarius.
Hale first established that there was no other deep-sky object near M70, consulted a directory of known comets, finding that none were known to be in this area of the sky. Once he had established that the object was moving relative to the background stars, he emailed the Central Bureau for Astronomical Telegrams, the clearing house for astronomical discoveries. Bopp did not own a telescope, he was out with friends near Stanfield, observing star clusters and galaxies when he chanced across the comet while at the eyepiece of his friend's telescope. He realized he might have spotted something new when, like Hale, he checked his star maps to determine if any other deep-sky objects were known to be near M70, found that there were none, he alerted the Central Bureau for Astronomical Telegrams through a Western Union telegram. Brian G. Marsden, who had run the bureau since 1968, laughed, "Nobody sends telegrams anymore. I mean, by the time that telegram got here, Alan Hale had e-mailed us three times with updated coordinates."The following morning, it was confirmed that this was a new comet, it was given the designation C/1995 O1.
The discovery was announced in International Astronomical Union circular 6187. Hale–Bopp's orbital position was calculated as 7.2 astronomical units from the Sun, placing it between Jupiter and Saturn and by far the greatest distance from Earth at which a comet had been discovered by amateurs. Most comets at this distance are faint, show no discernible activity, but Hale–Bopp had an observable coma. A precovery image taken at the Anglo-Australian Telescope in 1993 was found to show the then-unnoticed comet some 13 AU from the Sun, a distance at which most comets are unobservable. Analysis indicated that its comet nucleus was 60±20 kilometres in diameter six times the size of Halley, its great distance and surprising activity indicated that comet Hale–Bopp might become bright indeed when it reached perihelion in 1997. However, comet scientists were wary – comets can be unpredictable, many have large outbursts at great distance only to diminish in brightness later. Comet Kohoutek in 1973 had been touted as a'comet of the century' and turned out to be unspectacular.
Hale–Bopp became visible to the naked eye in May 1996, although its rate of brightening slowed during the latter half of that year, scientists were still cautiously optimistic that it would become bright. It was too aligned with the Sun to be observable during December 1996, but when it reappeared in January 1997 it was bright enough to be seen by anyone who looked for it from large cities with light-polluted skies; the Internet was a growing phenomenon at the time, numerous websites that tracked the comet's progress and provided daily images from around the world became popular. The Internet played a large role in encouraging the unprecedented public interest in comet Hale–Bopp; as the comet approached the Sun, it continued to brighten, shining at 2nd magnitude in February, showing a growing pair of tails, the blue gas tail pointing straight away from the Sun and the yellowish dust tail curving away along its orbit. On March 9, a solar eclipse in China and eastern Siberia allowed observers there to see the comet in the daytime.
Hale–Bopp had its closest approach to Earth on March 22, 1997, at a distance of 1.315 AU. As it passed perihelion on April 1, 1997, the comet developed into a spectacular sight, it shone brighter than any star in the sky except Sirius, its dust tail stretched 40–45 degrees across the sky. The comet was visible well before the sky got dark each night, while many great comets are close to the Sun as they pass perihelion, comet Hale–Bopp was visible all night to northern hemisphere observers. After its perihelion passage, the comet moved into the southern celestial hemisphere; the comet was much less impressive to southern hemisphere observers than it had been in the northern hemisphere, but southerners were able to see the comet fade from view during the second half of 1997. The last naked-eye observations were reported in December 1997, which meant that the comet had remained visible without aid for 569 days, or about 18 and a half months; the previous record had been set by the Great Comet of 1811, visible to the naked eye for about 9 months.
The comet continued to fade as it is still being tracked by astronomers. In October 2007, 10 years after the perihelion and at distance of 25.7 AU from Sun, the comet was still active as indicated by the detection of the CO-driven coma. H
Hydrogen is a chemical element with symbol H and atomic number 1. With a standard atomic weight of 1.008, hydrogen is the lightest element in the periodic table. Hydrogen is the most abundant chemical substance in the Universe, constituting 75% of all baryonic mass. Non-remnant stars are composed of hydrogen in the plasma state; the most common isotope of hydrogen, termed protium, has no neutrons. The universal emergence of atomic hydrogen first occurred during the recombination epoch. At standard temperature and pressure, hydrogen is a colorless, tasteless, non-toxic, nonmetallic combustible diatomic gas with the molecular formula H2. Since hydrogen forms covalent compounds with most nonmetallic elements, most of the hydrogen on Earth exists in molecular forms such as water or organic compounds. Hydrogen plays a important role in acid–base reactions because most acid-base reactions involve the exchange of protons between soluble molecules. In ionic compounds, hydrogen can take the form of a negative charge when it is known as a hydride, or as a positively charged species denoted by the symbol H+.
The hydrogen cation is written as though composed of a bare proton, but in reality, hydrogen cations in ionic compounds are always more complex. As the only neutral atom for which the Schrödinger equation can be solved analytically, study of the energetics and bonding of the hydrogen atom has played a key role in the development of quantum mechanics. Hydrogen gas was first artificially produced in the early 16th century by the reaction of acids on metals. In 1766–81, Henry Cavendish was the first to recognize that hydrogen gas was a discrete substance, that it produces water when burned, the property for which it was named: in Greek, hydrogen means "water-former". Industrial production is from steam reforming natural gas, less from more energy-intensive methods such as the electrolysis of water. Most hydrogen is used near the site of its production, the two largest uses being fossil fuel processing and ammonia production for the fertilizer market. Hydrogen is a concern in metallurgy as it can embrittle many metals, complicating the design of pipelines and storage tanks.
Hydrogen gas is flammable and will burn in air at a wide range of concentrations between 4% and 75% by volume. The enthalpy of combustion is −286 kJ/mol: 2 H2 + O2 → 2 H2O + 572 kJ Hydrogen gas forms explosive mixtures with air in concentrations from 4–74% and with chlorine at 5–95%; the explosive reactions may be triggered by heat, or sunlight. The hydrogen autoignition temperature, the temperature of spontaneous ignition in air, is 500 °C. Pure hydrogen-oxygen flames emit ultraviolet light and with high oxygen mix are nearly invisible to the naked eye, as illustrated by the faint plume of the Space Shuttle Main Engine, compared to the visible plume of a Space Shuttle Solid Rocket Booster, which uses an ammonium perchlorate composite; the detection of a burning hydrogen leak may require a flame detector. Hydrogen flames in other conditions are blue; the destruction of the Hindenburg airship was a notorious example of hydrogen combustion and the cause is still debated. The visible orange flames in that incident were the result of a rich mixture of hydrogen to oxygen combined with carbon compounds from the airship skin.
H2 reacts with every oxidizing element. Hydrogen can react spontaneously and violently at room temperature with chlorine and fluorine to form the corresponding hydrogen halides, hydrogen chloride and hydrogen fluoride, which are potentially dangerous acids; the ground state energy level of the electron in a hydrogen atom is −13.6 eV, equivalent to an ultraviolet photon of 91 nm wavelength. The energy levels of hydrogen can be calculated accurately using the Bohr model of the atom, which conceptualizes the electron as "orbiting" the proton in analogy to the Earth's orbit of the Sun. However, the atomic electron and proton are held together by electromagnetic force, while planets and celestial objects are held by gravity; because of the discretization of angular momentum postulated in early quantum mechanics by Bohr, the electron in the Bohr model can only occupy certain allowed distances from the proton, therefore only certain allowed energies. A more accurate description of the hydrogen atom comes from a purely quantum mechanical treatment that uses the Schrödinger equation, Dirac equation or the Feynman path integral formulation to calculate the probability density of the electron around the proton.
The most complicated treatments allow for the small effects of special relativity and vacuum polarization. In the quantum mechanical treatment, the electron in a ground state hydrogen atom has no angular momentum at all—illustrating how the "planetary orbit" differs from electron motion. There exist two different spin isomers of hydrogen diatomic molecules that differ by the relative spin of their nuclei. In the orthohydrogen form, the spins of the two protons are parallel and form a triplet state with a molecular spin quantum number of 1. At standard temperature and pressure, hydrogen gas contains about 25% of the para form and 75% of the ortho form known as the "normal form"; the equilibrium ratio of orthohydrogen to parahydrogen depends on temperature, but because the ortho form is an excited state and has a higher energy
Nitric oxide is a colorless gas with the formula NO. It is one of the principal oxides of nitrogen. Nitric oxide is a free radical, i.e. it has an unpaired electron, sometimes denoted by a dot in its chemical formula, i.e. ·NO. Nitric oxide is a heteronuclear diatomic molecule, a historic class that drew researches which spawned early modern theories of chemical bonding. An important intermediate in chemical industry, nitric oxide forms in combustion systems and can be generated by lightning in thunderstorms. In mammals, including humans, nitric oxide is a signaling molecule in many physiological and pathological processes, it was proclaimed the "Molecule of the Year" in 1992. The 1998 Nobel Prize in Physiology or Medicine was awarded for discovering nitric oxide's role as a cardiovascular signalling molecule. Nitric oxide should not be confused with nitrous oxide, an anesthetic, or with nitrogen dioxide, a brown toxic gas and a major air pollutant. Upon condensing to a liquid, nitric oxide dimerizes to dinitrogen dioxide, but the association is weak and reversible.
The N–N distance in crystalline NO is 218 pm, nearly twice the N–O distance. Since the heat of formation of ·NO is endothermic, NO can be decomposed to the elements. Catalytic converters in cars exploit this reaction: 2 NO → O2 + N2; when exposed to oxygen, nitric oxide converts into nitrogen dioxide: 2 NO + O2 → 2 NO2. This conversion has been speculated as occurring via the ONOONO intermediate. In water, nitric oxide reacts with water to form nitrous acid; the reaction is thought to proceed via the following stoichiometry: 4 NO + O2 + 2 H2O → 4 HNO2. Nitric oxide reacts with fluorine and bromine to form the nitrosyl halides, such as nitrosyl chloride: 2 NO + Cl2 → 2 NOCl. With NO2 a radical, NO combines to form the intensely blue dinitrogen trioxide: NO + NO2 ⇌ ON−NO2; the addition of a nitric oxide moiety to another molecule is referred to as nitrosylation. Nitric oxide reacts with acetone and an alkoxide to a diazeniumdiolate or nitrosohydroxylamine and methyl acetate: This reaction, discovered around 1898, remains of interest in nitric oxide prodrug research.
Nitric oxide can react directly with sodium methoxide, forming sodium formate and nitrous oxide. Nitric oxide reacts with transition metals to give complexes called metal nitrosyls; the most common bonding mode of nitric oxide is the terminal linear type. Alternatively, nitric oxide can serve as a one-electron pseudohalide. In such complexes, the M−N−O group is characterized by an angle between 120° and 140°; the NO group can bridge between metal centers through the nitrogen atom in a variety of geometries. In commercial settings, nitric oxide is produced by the oxidation of ammonia at 750–900 °C with platinum as catalyst: 4 NH3 + 5 O2 → 4 NO + 6 H2OThe uncatalyzed endothermic reaction of oxygen and nitrogen, effected at high temperature by lightning has not been developed into a practical commercial synthesis: N2 + O2 → 2 NO In the laboratory, nitric oxide is conveniently generated by reduction of dilute nitric acid with copper: 8 HNO3 + 3 Cu → 3 Cu2 + 4 H2O + 2 NOAn alternative route involves the reduction of nitrous acid in the form of sodium nitrite or potassium nitrite: 2 NaNO2 + 2 NaI + 2 H2SO4 → I2 + 4 NaHSO4 + 2 NO 2 NaNO2 + 2 FeSO4 + 3 H2SO4 → Fe23 + 2 NaHSO4 + 2 H2O + 2 NO 3 KNO2 + KNO3 + Cr2O3 → 2 K2CrO4 + 4 NOThe iron sulfate route is simple and has been used in undergraduate laboratory experiments.
So-called NONOate compounds are used for nitric oxide generation. Nitric oxide concentration can be determined using a chemiluminescent reaction involving ozone. A sample containing nitric oxide is mixed with a large quantity of ozone; the nitric oxide reacts with the ozone to produce oxygen and nitrogen dioxide, accompanied with emission of light: NO + O3 → NO2 + O2 + hνwhich can be measured with a photodetector. The amount of light produced is proportional to the amount of nitric oxide in the sample. Other methods of testing include electroanalysis, where ·NO reacts with an electrode to induce a current or voltage change; the detection of NO radicals in biological tissues is difficult due to the short lifetime and concentration of these radicals in tissues. One of the few practical methods is spin trapping of nitric oxide with iron-dithiocarbamate complexes and subsequent detection of the mono-nitrosyl-iron complex with electron paramagnetic resonance. A group of fluorescent dye indicators that are available in acetylated form for intracellular measurements exist.
The most common compound is 4,5-diaminofluorescein. Nitric oxide reacts with the hydroperoxy radical to form nitrogen dioxide, which can react with a hydroxyl radical to produce nitric acid: ·NO + HO2•→ •NO2 + •OH ·NO2 + •OH → HNO3Nitric acid, along with sulfuric acid, contribute acid rain deposition. Furthermore, ·NO participates in ozone layer depletion. In this process, nitric oxide reacts with stratospheric ozone to form O2 and nitrogen dioxide: ·NO + O3 → NO2 + O2As seen in the Concentration Measurement section, this reaction is utilized to measure concentrations of ·NO in control volumes; as seen in the Acid deposition section, nitric oxide can transform into nitrogen dioxide. Symptoms of short-term nitrogen dioxide exposure include nausea and headache. Long-term effects could include impaired respiratory function. NO is a gaseous signaling molecule, it is a key vertebrate biological messenger. It is
A molecule is an electrically neutral group of two or more atoms held together by chemical bonds. Molecules are distinguished from ions by their lack of electrical charge. However, in quantum physics, organic chemistry, biochemistry, the term molecule is used less also being applied to polyatomic ions. In the kinetic theory of gases, the term molecule is used for any gaseous particle regardless of its composition. According to this definition, noble gas atoms are considered molecules as they are monatomic molecules. A molecule may be homonuclear, that is, it consists of atoms of one chemical element, as with oxygen. Atoms and complexes connected by non-covalent interactions, such as hydrogen bonds or ionic bonds, are not considered single molecules. Molecules as components of matter are common in organic substances, they make up most of the oceans and atmosphere. However, the majority of familiar solid substances on Earth, including most of the minerals that make up the crust and core of the Earth, contain many chemical bonds, but are not made of identifiable molecules.
No typical molecule can be defined for ionic crystals and covalent crystals, although these are composed of repeating unit cells that extend either in a plane or three-dimensionally. The theme of repeated unit-cellular-structure holds for most condensed phases with metallic bonding, which means that solid metals are not made of molecules. In glasses, atoms may be held together by chemical bonds with no presence of any definable molecule, nor any of the regularity of repeating units that characterizes crystals; the science of molecules is called molecular chemistry or molecular physics, depending on whether the focus is on chemistry or physics. Molecular chemistry deals with the laws governing the interaction between molecules that results in the formation and breakage of chemical bonds, while molecular physics deals with the laws governing their structure and properties. In practice, this distinction is vague. In molecular sciences, a molecule consists of a stable system composed of two or more atoms.
Polyatomic ions may sometimes be usefully thought of as electrically charged molecules. The term unstable molecule is used for reactive species, i.e. short-lived assemblies of electrons and nuclei, such as radicals, molecular ions, Rydberg molecules, transition states, van der Waals complexes, or systems of colliding atoms as in Bose–Einstein condensate. According to Merriam-Webster and the Online Etymology Dictionary, the word "molecule" derives from the Latin "moles" or small unit of mass. Molecule – "extremely minute particle", from French molécule, from New Latin molecula, diminutive of Latin moles "mass, barrier". A vague meaning at first; the definition of the molecule has evolved. Earlier definitions were less precise, defining molecules as the smallest particles of pure chemical substances that still retain their composition and chemical properties; this definition breaks down since many substances in ordinary experience, such as rocks and metals, are composed of large crystalline networks of chemically bonded atoms or ions, but are not made of discrete molecules.
Molecules are held together by ionic bonding. Several types of non-metal elements exist only as molecules in the environment. For example, hydrogen only exists as hydrogen molecule. A molecule of a compound is made out of two or more elements. A covalent bond is a chemical bond; these electron pairs are termed shared pairs or bonding pairs, the stable balance of attractive and repulsive forces between atoms, when they share electrons, is termed covalent bonding. Ionic bonding is a type of chemical bond that involves the electrostatic attraction between oppositely charged ions, is the primary interaction occurring in ionic compounds; the ions are atoms that have lost one or more electrons and atoms that have gained one or more electrons. This transfer of electrons is termed electrovalence in contrast to covalence. In the simplest case, the cation is a metal atom and the anion is a nonmetal atom, but these ions can be of a more complicated nature, e.g. molecular ions like NH4+ or SO42−. An ionic bond is the transfer of electrons from a metal to a non-metal for both atoms to obtain a full valence shell.
Most molecules are far too small to be seen with the naked eye. DNA, a macromolecule, can reach macroscopic sizes, as can molecules of many polymers. Molecules used as building blocks for organic synthesis have a dimension of a few angstroms to several dozen Å, or around one billionth of a meter. Single molecules cannot be observed by light, but small molecules and the outlines of individual atoms may be traced in some circumstances by use of an atomic force microscope; some of the largest molecules are supermolecules. The smallest molecule is the diatomic hydrogen, with a bond length of 0.74 Å. Effective molecular radius is the size; the table of permselectivity for different substances contains examples. The chemical formula for a molecule uses one line of chemical element symbols and sometimes al
Diatomic molecules are molecules composed of only two atoms, of the same or different chemical elements. The prefix di- is of Greek origin, meaning "two". If a diatomic molecule consists of two atoms of the same element, such as hydrogen or oxygen it is said to be homonuclear. Otherwise, if a diatomic molecule consists of two different atoms, such as carbon monoxide or nitric oxide, the molecule is said to be heteronuclear; the only chemical elements that form stable homonuclear diatomic molecules at standard temperature and pressure are the gases hydrogen, oxygen and chlorine. The noble gases are gases at STP, but they are monatomic; the homonuclear diatomic gases and noble gases together are called "elemental gases" or "molecular gases", to distinguish them from other gases that are chemical compounds. At elevated temperatures, the halogens bromine and iodine form diatomic gases. All halogens have been observed as diatomic molecules, except for astatine, uncertain; the mnemonics BrINClHOF, pronounced "Brinklehof", HONClBrIF, pronounced "Honkelbrif", HOFBrINCl have been coined to aid recall of the list of diatomic elements.
Other elements form diatomic molecules when evaporated, but these diatomic species repolymerize when cooled. Heating elemental phosphorus gives diphosphorus, P2. Sulfur vapor is disulfur. Dilithium is known in the gas phase. Ditungsten and dimolybdenum form with sextuple bonds in the gas phase; the bond in a homonuclear diatomic molecule is non-polar. Dirubidium is diatomic. All other diatomic molecules are chemical compounds of two different elements. Many elements can combine to form heteronuclear diatomic molecules, depending on temperature and pressure; some examples include, gases carbon monoxide, nitric oxide, hydrogen chloride. Many 1:1 binary compounds are not considered diatomic because they are polymeric at room temperature, but they form diatomic molecules when evaporated, for example gaseous MgO, SiO, many others. Hundreds of diatomic molecules have been identified in the environment of the Earth, in the laboratory, in interstellar space. About 99% of the Earth's atmosphere is composed of two species of diatomic molecules: nitrogen and oxygen.
The natural abundance of hydrogen in the Earth's atmosphere is only of the order of parts per million, but H2 is the most abundant diatomic molecule in the universe. The interstellar medium is, dominated by hydrogen atoms. Diatomic elements played an important role in the elucidation of the concepts of element and molecule in the 19th century, because some of the most common elements, such as hydrogen and nitrogen, occur as diatomic molecules. John Dalton's original atomic hypothesis assumed that all elements were monatomic and that the atoms in compounds would have the simplest atomic ratios with respect to one another. For example, Dalton assumed water's formula to be HO, giving the atomic weight of oxygen as eight times that of hydrogen, instead of the modern value of about 16; as a consequence, confusion existed regarding atomic weights and molecular formulas for about half a century. As early as 1805, Gay-Lussac and von Humboldt showed that water is formed of two volumes of hydrogen and one volume of oxygen, by 1811 Amedeo Avogadro had arrived at the correct interpretation of water's composition, based on what is now called Avogadro's law and the assumption of diatomic elemental molecules.
However, these results were ignored until 1860 due to the belief that atoms of one element would have no chemical affinity toward atoms of the same element, partly due to apparent exceptions to Avogadro's law that were not explained until in terms of dissociating molecules. At the 1860 Karlsruhe Congress on atomic weights, Cannizzaro resurrected Avogadro's ideas and used them to produce a consistent table of atomic weights, which agree with modern values; these weights were an important prerequisite for the discovery of the periodic law by Dmitri Mendeleev and Lothar Meyer. Diatomic molecules are in their lowest or ground state, which conventionally is known as the X state; when a gas of diatomic molecules is bombarded by energetic electrons, some of the molecules may be excited to higher electronic states, as occurs, for example, in the natural aurora. Such excitation can occur when the gas absorbs light or other electromagnetic radiation; the excited states are unstable and relax back to the ground state.
Over various short time scales after the excitation, transitions occur from higher to lower electronic states and to the ground state, in each transition results a photon is emitted. This emission is known as fluorescence. Successively higher electronic states are conventionally named A, B, C, etc.. The excitation energy must be greater than or equal to the energy of the electronic state in order for the excitation to occur. In quantum theory, an electronic state of a diatomic molecule is represented by 2 S + 1 Λ ( v