Radium is a chemical element with symbol Ra and atomic number 88. It is the sixth element in group 2 of the periodic table known as the alkaline earth metals. Pure radium is silvery-white, but it reacts with nitrogen on exposure to air, forming a black surface layer of radium nitride. All isotopes of radium are radioactive, with the most stable isotope being radium-226, which has a half-life of 1600 years and decays into radon gas; when radium decays, ionizing radiation is a product, which can excite fluorescent chemicals and cause radioluminescence. Radium, in the form of radium chloride, was discovered by Marie and Pierre Curie in 1898, they extracted the radium compound from uraninite and published the discovery at the French Academy of Sciences five days later. Radium was isolated in its metallic state by Marie Curie and André-Louis Debierne through the electrolysis of radium chloride in 1911. In nature, radium is found in uranium and thorium ores in trace amounts as small as a seventh of a gram per ton of uraninite.
Radium is not necessary for living organisms, adverse health effects are when it is incorporated into biochemical processes because of its radioactivity and chemical reactivity. Other than its use in nuclear medicine, radium has no commercial applications. Today, these former applications are no longer in vogue because radium's toxicity has since become known, less dangerous isotopes are used instead in radioluminescent devices. Radium is the only radioactive member of its group, its physical and chemical properties most resemble its lighter congener barium. Pure radium is a volatile silvery-white metal, although its lighter congeners calcium and barium have a slight yellow tint; this tint vanishes on exposure to air, yielding a black layer of radium nitride. Its melting point is either 700 °C or 960 °C and its boiling point is 1,737 °C. Both of these values are lower than those of barium, confirming periodic trends down the group 2 elements. Like barium and the alkali metals, radium crystallizes in the body-centered cubic structure at standard temperature and pressure: the radium–radium bond distance is 514.8 picometers.
Radium has a density of 5.5 g/cm3, higher than that of barium, again confirming periodic trends. Radium has 33 known isotopes, with mass numbers from 202 to 234: all of them are radioactive. Four of these – 223Ra, 224Ra, 226Ra, 228Ra – occur in the decay chains of primordial thorium-232, uranium-235, uranium-238; these isotopes still have half-lives too short to be primordial radionuclides and only exist in nature from these decay chains. Together with the artificial 225Ra, these are the five most stable isotopes of radium. All other known radium isotopes have half-lives under two hours, the majority have half-lives under a minute. At least 12 nuclear isomers have been reported. In the early history of the study of radioactivity, the different natural isotopes of radium were given different names. In this scheme, 223Ra was named actinium X, 224Ra thorium X, 226Ra radium, 228Ra mesothorium 1; when it was realized that all of these are isotopes of the same element, many of these names fell out of use, "radium" came to refer to all isotopes, not just 226Ra.
Some of radium-226's decay products received historical names including "radium", ranging from radium A to radium G, with the letter indicating how far they were down the chain from their parent 226Ra.226Ra is the most stable isotope of radium and is the last isotope in the decay chain of uranium-238 with a half-life of over a millennium: it makes up all of natural radium. Its immediate decay product is the dense radioactive noble gas radon, responsible for much of the danger of environmental radium, it is 2.7 million times more radioactive than the same molar amount of natural uranium, due to its proportionally shorter half-life. A sample of radium metal maintains itself at a higher temperature than its surroundings because of the radiation it emits – alpha particles, beta particles, gamma rays. More natural radium emits alpha particles, but other steps in its decay chain emit alpha or beta particles, all particle emissions are accompanied by gamma rays. In 2013 it was discovered; this was the first discovery of an asymmetric nucleus.
Radium, like barium, is a reactive metal and always exhibits its group oxidation state of +2. It forms the colorless Ra2+ cation in aqueous solution, basic and does not form complexes readily. Most radium compounds are therefore simple ionic compounds, though participation from the 6s and 6p electrons is expected due to relativistic effects and would enhance the covalent character of radium compounds such as RaF2 and RaAt2. For this reason, the standard electrode potential for the half-reaction Ra2+ + 2e− →
Manganese is a chemical element with symbol Mn and atomic number 25. It is not found as a free element in nature. Manganese is a metal with important industrial metal alloy uses in stainless steels. Manganese is named for pyrolusite and other black minerals from the region of Magnesia in Greece, which gave its name to magnesium and the iron ore magnetite. By the mid-18th century, Swedish-German chemist Carl Wilhelm Scheele had used pyrolusite to produce chlorine. Scheele and others were aware that pyrolusite contained a new element, but they were unable to isolate it. Johan Gottlieb Gahn was the first to isolate an impure sample of manganese metal in 1774, which he did by reducing the dioxide with carbon. Manganese phosphating is used for corrosion prevention on steel. Ionized manganese is used industrially as pigments of various colors, which depend on the oxidation state of the ions; the permanganates of alkali and alkaline earth metals are powerful oxidizers. Manganese dioxide is used as the cathode material in alkaline batteries.
In biology, manganese ions function as cofactors for a large variety of enzymes with many functions. Manganese enzymes are essential in detoxification of superoxide free radicals in organisms that must deal with elemental oxygen. Manganese functions in the oxygen-evolving complex of photosynthetic plants. While the element is a required trace mineral for all known living organisms, it acts as a neurotoxin in larger amounts. Through inhalation, it can cause manganism, a condition in mammals leading to neurological damage, sometimes irreversible. Manganese is a silvery-gray metal, it is hard and brittle, difficult to fuse, but easy to oxidize. Manganese metal and its common ions are paramagnetic. Manganese tarnishes in air and oxidizes like iron in water containing dissolved oxygen. Occurring manganese is composed of one stable isotope, 55Mn. Eighteen radioisotopes have been isolated and described, ranging in atomic weight from 46 u to 65 u; the most stable are 53Mn with a half-life of 3.7 million years, 54Mn with a half-life of 312.3 days, 52Mn with a half-life of 5.591 days.
All of the remaining radioactive isotopes have half-lives of less than three hours, the majority of less than one minute. The primary decay mode before the most abundant stable isotope, 55Mn, is electron capture and the primary mode after is beta decay. Manganese has three meta states. Manganese is part of the iron group of elements, which are thought to be synthesized in large stars shortly before the supernova explosion. 53Mn decays to 53Cr with a half-life of 3.7 million years. Because of its short half-life, 53Mn is rare, produced by cosmic rays impact on iron. Manganese isotopic contents are combined with chromium isotopic contents and have found application in isotope geology and radiometric dating. Mn–Cr isotopic ratios reinforce the evidence from 26Al and 107Pd for the early history of the solar system. Variations in 53Cr/52Cr and Mn/Cr ratios from several meteorites suggest an initial 53Mn/55Mn ratio, which indicates that Mn–Cr isotopic composition must result from in situ decay of 53Mn in differentiated planetary bodies.
Hence, 53Mn provides additional evidence for nucleosynthetic processes before coalescence of the solar system. The most common oxidation states of manganese are +2, +3, +4, +6, +7, though all oxidation states from −3 to +7 have been observed. Mn2+ competes with Mg2+ in biological systems. Manganese compounds where manganese is in oxidation state +7, which are restricted to the unstable oxide Mn2O7, compounds of the intensely purple permanganate anion MnO4−, a few oxyhalides, are powerful oxidizing agents. Compounds with oxidation states +5 and +6 are strong oxidizing agents and are vulnerable to disproportionation; the most stable oxidation state for manganese is +2, which has a pale pink color, many manganese compounds are known, such as manganese sulfate and manganese chloride. This oxidation state is seen in the mineral rhodochrosite. Manganese most exists with a high spin, S = 5/2 ground state because of the high pairing energy for manganese. However, there are a few examples of S = 1/2 manganese.
There are no spin-allowed d–d transitions in manganese, explaining why manganese compounds are pale to colorless. The +3 oxidation state is known in compounds like manganese acetate, but these are quite powerful oxidizing agents and prone to disproportionation in solution, forming manganese and manganese. Solid compounds of manganese are characterized by its strong purple-red color and a preference for distorted octahedral coordination resulting from the Jahn-Teller effect; the oxidation state +5 can be produced by dissolving manganese dioxide in molten sodium nitrite. Manganate salts can be produced by dissolving Mn compounds, such as manganese dioxide, in molten alkali while exposed to air. Permanganate compounds are purple, can give glass a violet color. Potassium permanganate, sodium permanganate, barium permanganate are all potent oxidizers. Potassium permanganate called Condy's crystals, is a used laboratory reagent because of its oxidizing properties. Solutions of potassium permanganate were among the first stains and fixatives to be used in the preparation of biological cells and tissues for electron microscopy
Luminous paint or luminescent paint is paint that exhibits luminescence. In other words, it gives off visible light through fluorescence, phosphorescence, or radioluminescence. There are three types of luminous paints.. Fluorescent paints offer a wide range of pigments and chroma which also'glow' when exposed to the long-wave "ultraviolet" frequencies; these UV frequencies are found in sunlight and some artificial lights, but they—and their glowing-paint applications—are popularly known as black light and'black-light effects', respectively. In fluorescence the visible light component—sometimes known as "white light"—tends to be reflected and perceived as colour. Human eyes perceive these changes as the unusual'glow' of fluorescence; the fluorescent type of luminescence is different from the natural bioluminescence of bacteria and fish such as the case of the firefly, etc. Bio-luminescence involves no reflection at all. There are both invisible fluorescent paints; the visible appear under white light to be any bright color, turning peculiarly brilliant under black lights.
Invisible fluorescent paints appear transparent or pale under daytime lighting, but will glow under UV light in a limited range of colors. Since these can seem to'disappear', they can be used to create a variety of clever effects. Both types of fluorescent painting benefit when used within a contrasting ambiance of clean, matte-black backgrounds and borders; such a "black out" effect will minimize other awareness, so cultivating the peculiar luminescence of UV fluorescence. Both types of paints have extensive application where artistic lighting effects are desired in "black box" entertainments and environments such as theaters, shrines, etc. Out-of-doors, however, UV wavelengths are scattered in space or absorbed by complex natural surfaces, dulling the effect. Furthermore, the complex pigments will degrade in sunlight. Phosphorescent paint is called "glow-in-the-dark" paint, it is made from phosphors such as silver-activated zinc sulfide or doped strontium aluminate, glows a pale green to greenish-blue color.
The mechanism for producing light is similar to that of fluorescent paint, but the emission of visible light persists long after it has been exposed to light. Phosphorescent paints have a sustained glow which lasts for up to 12 hours after exposure to light, fading over time; this type of paint has been used to mark escape paths in aircraft and for decorative use such as "stars" applied to walls and ceilings. It is an alternative to radioluminescent paint. Kenner's Lightning Bug Glo-Juice was a popular non-toxic paint product in 1968, marketed at children, alongside other glow-in-the-dark toys and novelties. Phosphorescent paint is used as body paint, on children's walls and outdoors; when applied as a paint or a more sophisticated coating, phosphorescence can be used for temperature detection or degradation measurements known as phosphor thermometry. Radioluminescent paint is a self-luminous paint that consists of a small amount of a radioactive isotope mixed with a radioluminescent phosphor chemical.
The radioisotope continually decays, emitting radiation particles which strike molecules of the phosphor, exciting them to emit visible light. The isotopes selected are strong emitters of beta radiation, preferred since this radiation will not penetrate an enclosure. Radioluminescent paints will glow without exposure to light until the radioactive isotope has decayed, which may be many years; because of safety concerns and tighter regulation, consumer products such as clocks and watches now use phosphorescent rather than radioluminescent substances. Radioluminescent paint may still be preferred in specialist applications, such as diving watches. Radioluminescent paint was invented in 1908 by Sabin Arnold von Sochocky and incorporated radium-226. Radium paint was used for 40 years on the faces of watches and aircraft instruments, so they could be read in the dark. Radium is a radiological hazard, emitting gamma rays that can penetrate a glass watch dial and into human tissue. During the 1920s and 1930s, the harmful effects of this paint became clear.
A notorious case involved the "Radium Girls", a group of women who painted watchfaces and suffered adverse health effects from ingestion. In 1928, Dr von Sochocky himself died of aplastic anemia as a result of radiation exposure. Radium was banned from this use decades ago by international law, but the thousands of legacy radium dials still owned by the public can be a dangerous source of radioactive contamination. Radium paint used zinc sulfide phosphor trace metal doped with an activator, such as copper and more copper-magnesium; the phosphor degrades fast and the dials lose luminosity in several years to a few decades. However, due to the long 1600 year half-life of the Ra-226 isotope they are still radioactive and can be identified with a Geiger counter; the dials can be renovate
A radionuclide is an atom that has excess nuclear energy, making it unstable. This excess energy can be used in one of three ways: emitted from the nucleus as gamma radiation. During those processes, the radionuclide is said to undergo radioactive decay; these emissions are considered ionizing radiation because they are powerful enough to liberate an electron from another atom. The radioactive decay can produce a stable nuclide or will sometimes produce a new unstable radionuclide which may undergo further decay. Radioactive decay is a random process at the level of single atoms: it is impossible to predict when one particular atom will decay. However, for a collection of atoms of a single element the decay rate, thus the half-life for that collection can be calculated from their measured decay constants; the range of the half-lives of radioactive atoms have no known limits and span a time range of over 55 orders of magnitude. Radionuclides occur or are artificially produced in nuclear reactors, particle accelerators or radionuclide generators.
There are about 730 radionuclides with half-lives longer than 60 minutes. Thirty-two of those are primordial radionuclides. At least another 60 radionuclides are detectable in nature, either as daughters of primordial radionuclides or as radionuclides produced through natural production on Earth by cosmic radiation. More than 2400 radionuclides have half-lives less than 60 minutes. Most of those are only produced artificially, have short half-lives. For comparison, there are about 253 stable nuclides. All chemical elements can exist as radionuclides; the lightest element, has a well-known radionuclide, tritium. Elements heavier than lead, the elements technetium and promethium, exist only as radionuclides. Unplanned exposure to radionuclides has a harmful effect on living organisms including humans, although low levels of exposure occur without harm; the degree of harm will depend on the nature and extent of the radiation produced, the amount and nature of exposure, the biochemical properties of the element.
However, radionuclides with suitable properties are used in nuclear medicine for both diagnosis and treatment. An imaging tracer made with radionuclides is called a radioactive tracer. A pharmaceutical drug made with radionuclides is called a radiopharmaceutical. On Earth occurring radionuclides fall into three categories: primordial radionuclides, secondary radionuclides, cosmogenic radionuclides. Radionuclides are produced in stellar nucleosynthesis and supernova explosions along with stable nuclides. Most decay but can still be observed astronomically and can play a part in understanding astronomic processes. Primordial radionuclides, such as uranium and thorium, exist in the present time because their half-lives are so long that they have not yet decayed; some radionuclides have half-lives so long that decay has only been detected, for most practical purposes they can be considered stable, most notably bismuth-209: detection of this decay meant that bismuth was no longer considered stable.
It is possible decay may be observed in other nuclides adding to this list of primordial radionuclides. Secondary radionuclides are radiogenic isotopes derived from the decay of primordial radionuclides, they have shorter half-lives than primordial radionuclides. They arise in the decay chain of the primordial isotopes thorium-232, uranium-238 and uranium-235. Examples include the natural isotopes of radium. Cosmogenic isotopes, such as carbon-14, are present because they are continually being formed in the atmosphere due to cosmic rays. Many of these radionuclides exist only in trace amounts in nature, including all cosmogenic nuclides. Secondary radionuclides will occur in proportion to their half-lives, so short-lived ones will be rare, thus polonium can be found in uranium ores at about 0.1 mg per metric ton. Further radionunclides may occur in nature in undetectable amounts as a result of rare events such as spontaneous fission or uncommon cosmic ray interactions. Radionuclides are produced as an unavoidable result of nuclear thermonuclear explosions.
The process of nuclear fission creates a wide range of fission products, most of which are radionuclides. Further radionuclides can be created from irradiation of the nuclear fuel and of the surrounding structures, yielding activation products; this complex mixture of radionuclides with different chemistries and radioactivity makes handling nuclear waste and dealing with nuclear fallout problematic. Synthetic radionuclides are deliberately synthesised using nuclear reactors, particle accelerators or radionuclide generators: As well as being extracted from nuclear waste, radioisotopes can be produced deliberately with nuclear reactors, exploiting the high flux of neutrons present; these neutrons activate elements placed within the reactor. A typical product from a nuclear reactor is iridium-
Fluorescence is the emission of light by a substance that has absorbed light or other electromagnetic radiation. It is a form of luminescence. In most cases, the emitted light has a longer wavelength, therefore lower energy, than the absorbed radiation; the most striking example of fluorescence occurs when the absorbed radiation is in the ultraviolet region of the spectrum, thus invisible to the human eye, while the emitted light is in the visible region, which gives the fluorescent substance a distinct color that can be seen only when exposed to UV light. Fluorescent materials cease to glow nearly when the radiation source stops, unlike phosphorescent materials, which continue to emit light for some time after. Fluorescence has many practical applications, including mineralogy, medicine, chemical sensors, fluorescent labelling, biological detectors, cosmic-ray detection, most fluorescent lamps. Fluorescence occurs in nature in some minerals and in various biological states in many branches of the animal kingdom.
An early observation of fluorescence was described in 1560 by Bernardino de Sahagún and in 1565 by Nicolás Monardes in the infusion known as lignum nephriticum. It was derived from the wood of Pterocarpus indicus and Eysenhardtia polystachya; the chemical compound responsible for this fluorescence is matlaline, the oxidation product of one of the flavonoids found in this wood. In 1819, Edward D. Clarke and in 1822 René Just Haüy described fluorescence in fluorites, Sir David Brewster described the phenomenon for chlorophyll in 1833 and Sir John Herschel did the same for quinine in 1845. In his 1852 paper on the "Refrangibility" of light, George Gabriel Stokes described the ability of fluorspar and uranium glass to change invisible light beyond the violet end of the visible spectrum into blue light, he named this phenomenon fluorescence: "I am inclined to coin a word, call the appearance fluorescence, from fluor-spar, as the analogous term opalescence is derived from the name of a mineral." The name was derived from the mineral fluorite, some examples of which contain traces of divalent europium, which serves as the fluorescent activator to emit blue light.
In a key experiment he used a prism to isolate ultraviolet radiation from sunlight and observed blue light emitted by an ethanol solution of quinine exposed by it. Fluorescence occurs when an orbital electron of a molecule, atom, or nanostructure, relaxes to its ground state by emitting a photon from an excited singlet state: Excitation: S 0 + h ν e x → S 1 Fluorescence: S 1 → S 0 + h ν e m + h e a t Here h ν is a generic term for photon energy with h = Planck's constant and ν = frequency of light; the specific frequencies of exciting and emitted lights are depended on the particular system. S0 is called the ground state of the fluorophore, S1 is its first excited singlet state. A molecule in S1 can relax by various competing pathways, it can undergo non-radiative relaxation in which the excitation energy is dissipated as heat to the solvent. Excited organic molecules can relax via conversion to a triplet state, which may subsequently relax via phosphorescence, or by a secondary non-radiative relaxation step.
Relaxation from S1 can occur through interaction with a second molecule through fluorescence quenching. Molecular oxygen is an efficient quencher of fluorescence just because of its unusual triplet ground state. In most cases, the emitted light has a longer wavelength, therefore lower energy, than the absorbed radiation. However, when the absorbed electromagnetic radiation is intense, it is possible for one electron to absorb two photons; the emitted radiation may be of the same wavelength as the absorbed radiation, termed "resonance fluorescence". Molecules that are excited through light absorption or via a different process can transfer energy to a second'sensitized' molecule, converted to its excited state and can fluoresce; the fluorescence quantum yield gives the efficiency of the fluorescence process. It is defined as the ratio of the number of photons emitted to the number of photons absorbed. Φ = Number of photons emitted Number of photons absorbed The maximum possible fluorescence quantum yield is 1.0.
Compounds with quantum yields of 0.10 are still considered quite fluorescent. Another way to define the quantum yield of fluorescence is by the rate of excited state decay: Φ = k f ∑ i k i where k f is the rate constant of spontaneous emission of radiation and ∑ i k i is the sum of all rates of
A spinthariscope is a device for observing individual nuclear disintegrations caused by the interaction of ionizing radiation with a phosphor or scintillator. The spinthariscope was invented by William Crookes in 1903. While observing the uniform fluorescence on a zinc sulfide screen created by the radioactive emissions of a sample of radium bromide, he spilled some of the sample, owing to its extreme rarity and cost, he was eager to find and recover it. Upon inspecting the zinc sulfide screen under a microscope, he noticed separate flashes of light created by individual alpha particle collisions with the screen. Crookes took his discovery a step further and invented a device intended to view these scintillations, it consisted of a small screen coated with zinc sulfide affixed to the end of a tube, with a tiny amount of radium salt suspended a short distance from the screen and a lens on the other end of the tube for viewing the screen. Crookes named his device from Greek σπινθήρ "spark". Spinthariscopes were replaced with more accurate and quantitative devices for measuring radiation in scientific experiments, but enjoyed a modest revival in the mid 20th century as children's educational toys.
In 1947, Kix cereal offered a Lone Ranger atomic bomb ring in exchange for a box top and 0.15 USD that contained a small one. Spinthariscopes can still be bought today as instructional novelties, but they now use americium or thorium. A spinthariscope plays a pivotal role in The Blue Ghost Mystery. Modern spinthariscope Elements of electricity: a practical discussion of the fundamental laws and... by Robert Andrews Millikan, Edwin Sherwood Bishop, American Technical Society
Ernest Rutherford, 1st Baron Rutherford of Nelson, HFRSE LLD, was a New Zealand-born British physicist who came to be known as the father of nuclear physics. Encyclopædia Britannica considers him to be the greatest experimentalist since Michael Faraday. In early work, Rutherford discovered the concept of radioactive half-life, the radioactive element radon, differentiated and named alpha and beta radiation; this work was performed at McGill University in Canada. It is the basis for the Nobel Prize in Chemistry he was awarded in 1908 "for his investigations into the disintegration of the elements, the chemistry of radioactive substances", for which he was the first Canadian and Oceanian Nobel laureate. Rutherford moved in 1907 to the Victoria University of Manchester in the UK, where he and Thomas Royds proved that alpha radiation is helium nuclei. Rutherford performed his most famous work. In 1911, although he could not prove that it was positive or negative, he theorized that atoms have their charge concentrated in a small nucleus, thereby pioneered the Rutherford model of the atom, through his discovery and interpretation of Rutherford scattering by the gold foil experiment of Hans Geiger and Ernest Marsden.
He conducted research that led to the first "splitting" of the atom in 1917 in a nuclear reaction between nitrogen and alpha particles, in which he discovered the proton. Rutherford became Director of the Cavendish Laboratory at the University of Cambridge in 1919. Under his leadership the neutron was discovered by James Chadwick in 1932 and in the same year the first experiment to split the nucleus in a controlled manner was performed by students working under his direction, John Cockcroft and Ernest Walton. After his death in 1937, he was honoured by being interred with the greatest scientists of the United Kingdom, near Sir Isaac Newton's tomb in Westminster Abbey; the chemical element rutherfordium was named after him in 1997. Ernest Rutherford was the son of James Rutherford, a farmer, his wife Martha Thompson from Hornchurch, England. James had emigrated to New Zealand from Perth, Scotland, "to raise a little flax and a lot of children". Ernest was born near Nelson, New Zealand, his first name was mistakenly spelled ` Earnest'.
Rutherford's mother Martha Thompson was a schoolteacher. He studied at Havelock School and Nelson College and won a scholarship to study at Canterbury College, University of New Zealand, where he participated in the debating society and played rugby. After gaining his BA, MA and BSc, doing two years of research during which he invented a new form of radio receiver, in 1895 Rutherford was awarded an 1851 Research Fellowship from the Royal Commission for the Exhibition of 1851, to travel to England for postgraduate study at the Cavendish Laboratory, University of Cambridge, he was among the first of the'aliens' allowed to do research at the university, under the inspiring leadership of J. J. Thomson, which aroused jealousies from the more conservative members of the Cavendish fraternity. With Thomson's encouragement, he managed to detect radio waves at half a mile and held the world record for the distance over which electromagnetic waves could be detected, though when he presented his results at the British Association meeting in 1896, he discovered he had been outdone by another lecturer, by the name of Marconi.
In 1898, Thomson recommended Rutherford for a position at McGill University in Canada. He was to replace Hugh Longbourne Callendar who held the chair of Macdonald Professor of physics and was coming to Cambridge. Rutherford was accepted, which meant that in 1900 he could marry Mary Georgina Newton to whom he had become engaged before leaving New Zealand. In 1901, he gained a DSc from the University of New Zealand. In 1907, Rutherford returned to Britain to take the chair of physics at the Victoria University of Manchester, he was knighted in 1914. During World War I, he worked on a top secret project to solve the practical problems of submarine detection by sonar. In 1916, he was awarded the Hector Memorial Medal. In 1919, he returned to the Cavendish succeeding J. J. Thomson as the Cavendish professor and Director. Under him, Nobel Prizes were awarded to James Chadwick for discovering the neutron, John Cockcroft and Ernest Walton for an experiment, to be known as splitting the atom using a particle accelerator, Edward Appleton for demonstrating the existence of the ionosphere.
In 1925, Rutherford pushed calls to the Government of New Zealand to support education and research, which led to the formation of the Department of Scientific and Industrial Research in the following year. Between 1925 and 1930, he served as President of the Royal Society, as president of the Academic Assistance Council which helped 1,000 university refugees from Germany, he was appointed to the Order of Merit in the 1925 New Year Honours and raised to the peerage as Baron Rutherford of Nelson, of Cambridge in the County of Cambridge in 1931, a title that became extinct upon his unexpected death in 1937. In 1933, Rutherford was one of the two inaugural recipients of the T. K. Sidey Medal, set up by the Royal Society of New Zealand as an award for outstanding scientific research. For some time before his death, Rutherford had a small hernia, which he had neglected to have fixed, it became strangulated, causing him to be violently ill. Despite an emergency operation in Lon