1.
Chemistry
–
Chemistry is a branch of physical science that studies the composition, structure, properties and change of matter. Chemistry is sometimes called the science because it bridges other natural sciences, including physics. For the differences between chemistry and physics see comparison of chemistry and physics, the history of chemistry can be traced to alchemy, which had been practiced for several millennia in various parts of the world. The word chemistry comes from alchemy, which referred to a set of practices that encompassed elements of chemistry, metallurgy, philosophy, astrology, astronomy, mysticism. An alchemist was called a chemist in popular speech, and later the suffix -ry was added to this to describe the art of the chemist as chemistry, the modern word alchemy in turn is derived from the Arabic word al-kīmīā. In origin, the term is borrowed from the Greek χημία or χημεία and this may have Egyptian origins since al-kīmīā is derived from the Greek χημία, which is in turn derived from the word Chemi or Kimi, which is the ancient name of Egypt in Egyptian. Alternately, al-kīmīā may derive from χημεία, meaning cast together, in retrospect, the definition of chemistry has changed over time, as new discoveries and theories add to the functionality of the science. The term chymistry, in the view of noted scientist Robert Boyle in 1661, in 1837, Jean-Baptiste Dumas considered the word chemistry to refer to the science concerned with the laws and effects of molecular forces. More recently, in 1998, Professor Raymond Chang broadened the definition of chemistry to mean the study of matter, early civilizations, such as the Egyptians Babylonians, Indians amassed practical knowledge concerning the arts of metallurgy, pottery and dyes, but didnt develop a systematic theory. Greek atomism dates back to 440 BC, arising in works by such as Democritus and Epicurus. In 50 BC, the Roman philosopher Lucretius expanded upon the theory in his book De rerum natura, unlike modern concepts of science, Greek atomism was purely philosophical in nature, with little concern for empirical observations and no concern for chemical experiments. Work, particularly the development of distillation, continued in the early Byzantine period with the most famous practitioner being the 4th century Greek-Egyptian Zosimos of Panopolis. He formulated Boyles law, rejected the four elements and proposed a mechanistic alternative of atoms. Before his work, though, many important discoveries had been made, the Scottish chemist Joseph Black and the Dutchman J. B. English scientist John Dalton proposed the theory of atoms, that all substances are composed of indivisible atoms of matter. Davy discovered nine new elements including the alkali metals by extracting them from their oxides with electric current, british William Prout first proposed ordering all the elements by their atomic weight as all atoms had a weight that was an exact multiple of the atomic weight of hydrogen. The inert gases, later called the noble gases were discovered by William Ramsay in collaboration with Lord Rayleigh at the end of the century, thereby filling in the basic structure of the table. Organic chemistry was developed by Justus von Liebig and others, following Friedrich Wöhlers synthesis of urea which proved that organisms were, in theory
2.
Rule of thumb
–
A rule of thumb is a principle with broad application that is not intended to be strictly accurate or reliable for every situation. It is a learned and easily applied procedure for approximately calculating or recalling some value. It is based not on theory but on practical experience, compare this to heuristic, a similar concept used in mathematical discourse, psychology, and computer science, particularly in algorithm design. The exact origin of the phrase is uncertain, the earliest known citation comes from J. Durham’s Heaven upon Earth,1685, ii. 217, Many profest Christians are like to foolish builders, who build by guess, sir William Hope wrote in his The Compleat Fencing Master,1692, What he doth, he doth by rule of Thumb, and not by Art. James Kellys The Complete Collection of Scottish Proverbs,1721, pg 257, includes, No Rule so good as Rule of Thumb, if it hit. According to phrases. org. uk, The phrase joins the whole nine yards as one that derives from some form of measurement. It is often claimed that the rule of thumb is derived from a law that limited the maximum thickness of a stick with which it was permissible for a man to beat his wife. English common law before the reign of Charles II permitted a man to give his wife moderate correction, but no rule of thumb has ever been the law in England. Belief in the existence of a rule of thumb to excuse spousal abuse can be traced as far back as 1782, in the United States, where the rule of thumb was mentioned in case law, it was usually to reject it as a legal standard. Legal decisions in Mississippi and North Carolina make reference to—and reject—an unnamed old doctrine or ancient law by which a man was allowed to beat his wife with a stick no wider than his thumb. In 1976, womens rights activist Del Martin used the phrase rule of thumb as a reference to describe such a doctrine. She was interpreted by many as claiming the doctrine as an origin of the phrase, and the connection gained currency in 1982. Commission on Civil Rights issued a report on wife abuse titled Under the Rule of Thumb, right-hand rule Oersteds law Rule of 72 Pattern language Heuristic argument Common sense Back-of-the-envelope calculation Modulor Satisficing Rules of Thumb
3.
Atom
–
An atom is the smallest constituent unit of ordinary matter that has the properties of a chemical element. Every solid, liquid, gas, and plasma is composed of neutral or ionized atoms, Atoms are very small, typical sizes are around 100 picometers. Atoms are small enough that attempting to predict their behavior using classical physics - as if they were billiard balls, through the development of physics, atomic models have incorporated quantum principles to better explain and predict the behavior. Every atom is composed of a nucleus and one or more bound to the nucleus. The nucleus is made of one or more protons and typically a number of neutrons. Protons and neutrons are called nucleons, more than 99. 94% of an atoms mass is in the nucleus. The protons have an electric charge, the electrons have a negative electric charge. If the number of protons and electrons are equal, that atom is electrically neutral, if an atom has more or fewer electrons than protons, then it has an overall negative or positive charge, respectively, and it is called an ion. The electrons of an atom are attracted to the protons in a nucleus by this electromagnetic force. The number of protons in the nucleus defines to what chemical element the atom belongs, for example, the number of neutrons defines the isotope of the element. The number of influences the magnetic properties of an atom. Atoms can attach to one or more other atoms by chemical bonds to form compounds such as molecules. The ability of atoms to associate and dissociate is responsible for most of the changes observed in nature. The idea that matter is made up of units is a very old idea, appearing in many ancient cultures such as Greece. The word atom was coined by ancient Greek philosophers, however, these ideas were founded in philosophical and theological reasoning rather than evidence and experimentation. As a result, their views on what look like. They also could not convince everybody, so atomism was but one of a number of competing theories on the nature of matter. It was not until the 19th century that the idea was embraced and refined by scientists, in the early 1800s, John Dalton used the concept of atoms to explain why elements always react in ratios of small whole numbers
4.
Main-group element
–
The main group includes the elements in groups 1 and 2, and groups 13 to 18. In older nomenclature the main-group elements are groups IA and IIA, group 12 is labelled as group IIB in both systems. Group 3 is labelled as group IIIA in the older nomenclature, main-group elements are the most abundant elements on Earth, in the Solar System, and in the Universe. They are sometimes called the representative elements. An organometallic compound consisting of a carbon attached with METAL ATTĜķاﺁﺁئ Ralf Steudel, Chemie der Nichtmetalle, 2nd Edition
5.
Electron
–
The electron is a subatomic particle, symbol e− or β−, with a negative elementary electric charge. Electrons belong to the first generation of the lepton particle family, the electron has a mass that is approximately 1/1836 that of the proton. Quantum mechanical properties of the include a intrinsic angular momentum of a half-integer value, expressed in units of the reduced Planck constant. As it is a fermion, no two electrons can occupy the same state, in accordance with the Pauli exclusion principle. Like all elementary particles, electrons exhibit properties of particles and waves, they can collide with other particles and can be diffracted like light. Since an electron has charge, it has an electric field. Electromagnetic fields produced from other sources will affect the motion of an electron according to the Lorentz force law, electrons radiate or absorb energy in the form of photons when they are accelerated. Laboratory instruments are capable of trapping individual electrons as well as electron plasma by the use of electromagnetic fields, special telescopes can detect electron plasma in outer space. Electrons are involved in applications such as electronics, welding, cathode ray tubes, electron microscopes, radiation therapy, lasers, gaseous ionization detectors. Interactions involving electrons with other particles are of interest in fields such as chemistry. The Coulomb force interaction between the positive protons within atomic nuclei and the negative electrons without, allows the composition of the two known as atoms, ionization or differences in the proportions of negative electrons versus positive nuclei changes the binding energy of an atomic system. The exchange or sharing of the electrons between two or more atoms is the cause of chemical bonding. In 1838, British natural philosopher Richard Laming first hypothesized the concept of a quantity of electric charge to explain the chemical properties of atoms. Irish physicist George Johnstone Stoney named this charge electron in 1891, electrons can also participate in nuclear reactions, such as nucleosynthesis in stars, where they are known as beta particles. Electrons can be created through beta decay of isotopes and in high-energy collisions. The antiparticle of the electron is called the positron, it is identical to the electron except that it carries electrical, when an electron collides with a positron, both particles can be totally annihilated, producing gamma ray photons. The ancient Greeks noticed that amber attracted small objects when rubbed with fur, along with lightning, this phenomenon is one of humanitys earliest recorded experiences with electricity. In his 1600 treatise De Magnete, the English scientist William Gilbert coined the New Latin term electricus, both electric and electricity are derived from the Latin ēlectrum, which came from the Greek word for amber, ἤλεκτρον
6.
Electron shell
–
In chemistry and atomic physics, an electron shell, or a principal energy level, may be thought of as an orbit followed by electrons around an atoms nucleus. The closest shell to the nucleus is called the 1 shell, followed by the 2 shell, then the 3 shell, the shells correspond with the principal quantum numbers or are labeled alphabetically with letters used in the X-ray notation. Each shell can contain only a number of electrons, The first shell can hold up to two electrons, the second shell can hold up to eight electrons, the third shell can hold up to 18. The general formula is that the nth shell can in principle hold up to 2 electrons, since electrons are electrically attracted to the nucleus, an atoms electrons will generally occupy outer shells only if the more inner shells have already been completely filled by other electrons. However, this is not a requirement, atoms may have two or even three incomplete outer shells. For an explanation of why electrons exist in these shells see electron configuration, the electrons in the outermost occupied shell determine the chemical properties of the atom, it is called the valence shell. Each shell consists of one or more subshells, and each consists of one or more atomic orbitals. The shell terminology comes from Arnold Sommerfelds modification of the Bohr model, sommerfeld retained Bohrs planetary model, but added mildly elliptical orbits to explain the fine spectroscopic structure of some elements. The multiple electrons with the principal quantum number had close orbits that formed a shell of positive thickness instead of the infinitely thin circular orbit of Bohrs model. The existence of electron shells was first observed experimentally in Charles Barklas, barkla labeled them with the letters K, L, M, N, O, P, and Q. The origin of this terminology was alphabetic, a J series was also suspected, though later experiments indicated that the K absorption lines are produced by the innermost electrons. These letters were found to correspond to the n values 1,2,3. They are used in the spectroscopic Siegbahn notation, the physical chemist Gilbert Lewis was responsible for much of the early development of the theory of the participation of valence shell electrons in chemical bonding. Linus Pauling later generalized and extended the theory while applying insights from quantum mechanics. The electron shells are labeled K, L, M, N, O, P, and Q, or 1,2,3,4,5,6, and 7, going from innermost shell outwards. Electrons in outer shells have higher energy and travel farther from the nucleus than those in inner shells. This makes them important in determining how the atom reacts chemically and behaves as a conductor, because the pull of the atoms nucleus upon them is weaker. In this way, a given elements reactivity is highly dependent upon its electronic configuration, each shell is composed of one or more subshells, which are themselves composed of atomic orbitals
7.
Electron configuration
–
In atomic physics and quantum chemistry, the electron configuration is the distribution of electrons of an atom or molecule in atomic or molecular orbitals. For example, the configuration of the neon atom is 1s2 2s2 2p6. Electronic configurations describe electrons as each moving independently in an orbital, mathematically, configurations are described by Slater determinants or configuration state functions. Knowledge of the configuration of different atoms is useful in understanding the structure of the periodic table of elements. This is also useful for describing the chemical bonds that hold atoms together, in bulk materials, this same idea helps explain the peculiar properties of lasers and semiconductors. An electron shell is the set of allowed states that share the principal quantum number, n. An atoms nth electron shell can accommodate 2n2 electrons, e. g. the first shell can accommodate 2 electrons, the second shell 8 electrons, a subshell is the set of states defined by a common azimuthal quantum number, ℓ, within a shell. The values ℓ =0,1,2,3 correspond to the s, p, d, for example the 3d subshell has n =3 and ℓ =2. The maximum number of electrons that can be placed in a subshell is given by 2 and this gives two electrons in an s subshell, six electrons in a p subshell, ten electrons in a d subshell and fourteen electrons in an f subshell. Physicists and chemists use a notation to indicate the electron configurations of atoms. For atoms, the notation consists of a sequence of atomic subshell labels with the number of electrons assigned to each subshell placed as a superscript, for example, hydrogen has one electron in the s-orbital of the first shell, so its configuration is written 1s1. Lithium has two electrons in the 1s-subshell and one in the 2s-subshell, so its configuration is written 1s2 2s1, phosphorus is as follows, 1s2 2s2 2p6 3s2 3p3. For atoms with many electrons, this notation can become lengthy, phosphorus, for instance, is in the third period. It differs from the neon, whose configuration is 1s2 2s2 2p6. This convention is useful as it is the electrons in the outermost shell that most determine the chemistry of the element, for a given configuration, the order of writing the orbitals is not completely fixed since only the orbital occupancies have physical significance. For example, the configuration of the titanium ground state can be written as either 4s2 3d2 or 3d2 4s2. The first notation follows the order based on the Madelung rule for the configurations of atoms, 4s is filled before 3d in the sequence Ar, K, Ca, Sc. The superscript 1 for a singly occupied subshell is not compulsory and it is quite common to see the letters of the orbital labels written in an italic or slanting typeface, although the International Union of Pure and Applied Chemistry recommends a normal typeface
8.
Noble gas
–
The noble gases make a group of chemical elements with similar properties, under standard conditions, they are all odorless, colorless, monatomic gases with very low chemical reactivity. The six noble gases that occur naturally are helium, neon, argon, krypton, xenon, oganesson is predicted to be a noble gas as well, but its chemistry has not yet been investigated. For the first six periods of the table, the noble gases are exactly the members of group 18 of the periodic table. Noble gases are typically highly unreactive except when under particular extreme conditions, the inertness of noble gases makes them very suitable in applications where reactions are not wanted. The melting and boiling points for a noble gas are close together, differing by less than 10 °C. Neon, argon, krypton, and xenon are obtained from air in an air separation unit using the methods of liquefaction of gases, Noble gases have several important applications in industries such as lighting, welding, and space exploration. After the risks caused by the flammability of hydrogen became apparent, it was replaced with helium in blimps, Noble gas is translated from the German noun Edelgas, first used in 1898 by Hugo Erdmann to indicate their extremely low level of reactivity. The name makes an analogy to the noble metals, which also have low reactivity. The noble gases have also referred to as inert gases. Rare gases is another term that was used, but this is inaccurate because argon forms a fairly considerable part of the Earths atmosphere due to decay of radioactive potassium-40. Pierre Janssen and Joseph Norman Lockyer discovered a new element on August 18,1868 while looking at the chromosphere of the Sun, no chemical analysis was possible at the time, but helium was later found to be a noble gas. Before them, in 1784, the English chemist and physicist Henry Cavendish had discovered that air contains a proportion of a substance less reactive than nitrogen. A century later, in 1895, Lord Rayleigh discovered that samples of nitrogen from the air were of a different density than nitrogen resulting from chemical reactions, with this discovery, they realized an entire class of gases was missing from the periodic table. During his search for argon, Ramsay also managed to isolate helium for the first time while heating cleveite, Ramsay continued to search for these gases using the method of fractional distillation to separate liquid air into several components. In 1898, he discovered the elements krypton, neon, and xenon, and named them after the Greek words κρυπτός, νέος, and ξένος, respectively. Rayleigh and Ramsay received the 1904 Nobel Prizes in Physics and in Chemistry, respectively, for their discovery of the noble gases, the discovery of the noble gases aided in the development of a general understanding of atomic structure. In 1895, French chemist Henri Moissan attempted to form a reaction between fluorine, the most electronegative element, and argon, one of the noble gases, but failed. Scientists were unable to prepare compounds of argon until the end of the 20th century, in 1962, Neil Bartlett discovered the first chemical compound of a noble gas, xenon hexafluoroplatinate
9.
Carbon
–
Carbon is a chemical element with symbol C and atomic number 6. It is nonmetallic and tetravalent—making four electrons available to form covalent chemical bonds, three isotopes occur naturally, 12C and 13C being stable, while 14C is a radioactive isotope, decaying with a half-life of about 5,730 years. Carbon is one of the few elements known since antiquity, Carbon is the 15th most abundant element in the Earths crust, and the fourth most abundant element in the universe by mass after hydrogen, helium, and oxygen. It is the second most abundant element in the body by mass after oxygen. The atoms of carbon can bond together in different ways, termed allotropes of carbon, the best known are graphite, diamond, and amorphous carbon. The physical properties of carbon vary widely with the allotropic form, for example, graphite is opaque and black while diamond is highly transparent. Graphite is soft enough to form a streak on paper, while diamond is the hardest naturally occurring material known, graphite is a good electrical conductor while diamond has a low electrical conductivity. Under normal conditions, diamond, carbon nanotubes, and graphene have the highest thermal conductivities of all known materials, all carbon allotropes are solids under normal conditions, with graphite being the most thermodynamically stable form. They are chemically resistant and require high temperature to react even with oxygen, the most common oxidation state of carbon in inorganic compounds is +4, while +2 is found in carbon monoxide and transition metal carbonyl complexes. The largest sources of carbon are limestones, dolomites and carbon dioxide, but significant quantities occur in organic deposits of coal, peat, oil. For this reason, carbon has often referred to as the king of the elements. The allotropes of carbon graphite, one of the softest known substances, and diamond. It bonds readily with other small atoms including other carbon atoms, Carbon is known to form almost ten million different compounds, a large majority of all chemical compounds. Carbon also has the highest sublimation point of all elements, although thermodynamically prone to oxidation, carbon resists oxidation more effectively than elements such as iron and copper that are weaker reducing agents at room temperature. Carbon is the element, with a ground-state electron configuration of 1s22s22p2. Its first four ionisation energies,1086.5,2352.6,4620.5 and 6222.7 kJ/mol, are higher than those of the heavier group 14 elements. Carbons covalent radii are normally taken as 77.2 pm,66.7 pm and 60.3 pm, although these may vary depending on coordination number, in general, covalent radius decreases with lower coordination number and higher bond order. Carbon compounds form the basis of all life on Earth
10.
Nitrogen
–
Nitrogen is a chemical element with symbol N and atomic number 7. It was first discovered and isolated by Scottish physician Daniel Rutherford in 1772, although Carl Wilhelm Scheele and Henry Cavendish had independently done so at about the same time, Rutherford is generally accorded the credit because his work was published first. Nitrogen is the lightest member of group 15 of the periodic table, the name comes from the Greek πνίγειν to choke, directly referencing nitrogens asphyxiating properties. It is an element in the universe, estimated at about seventh in total abundance in the Milky Way. At standard temperature and pressure, two atoms of the element bind to form dinitrogen, a colourless and odorless diatomic gas with the formula N2, dinitrogen forms about 78% of Earths atmosphere, making it the most abundant uncombined element. Nitrogen occurs in all organisms, primarily in amino acids, in the nucleic acids, the human body contains about 3% nitrogen by mass, the fourth most abundant element in the body after oxygen, carbon, and hydrogen. The nitrogen cycle describes movement of the element from the air, into the biosphere and organic compounds, many industrially important compounds, such as ammonia, nitric acid, organic nitrates, and cyanides, contain nitrogen. The extremely strong bond in elemental nitrogen, the second strongest bond in any diatomic molecule. Synthetically produced ammonia and nitrates are key industrial fertilisers, and fertiliser nitrates are key pollutants in the eutrophication of water systems. Apart from its use in fertilisers and energy-stores, nitrogen is a constituent of organic compounds as diverse as Kevlar used in high-strength fabric, Nitrogen is a constituent of every major pharmacological drug class, including antibiotics. Many notable nitrogen-containing drugs, such as the caffeine and morphine or the synthetic amphetamines. Nitrogen compounds have a long history, ammonium chloride having been known to Herodotus. They were well known by the Middle Ages, alchemists knew nitric acid as aqua fortis, as well as other nitrogen compounds such as ammonium salts and nitrate salts. The mixture of nitric and hydrochloric acids was known as aqua regia, celebrated for its ability to dissolve gold, the discovery of nitrogen is attributed to the Scottish physician Daniel Rutherford in 1772, who called it noxious air. Though he did not recognise it as a different chemical substance, he clearly distinguished it from Joseph Blacks fixed air. The fact that there was a component of air that does not support combustion was clear to Rutherford, Nitrogen was also studied at about the same time by Carl Wilhelm Scheele, Henry Cavendish, and Joseph Priestley, who referred to it as burnt air or phlogisticated air. Nitrogen gas was inert enough that Antoine Lavoisier referred to it as air or azote, from the Greek word άζωτικός. In an atmosphere of nitrogen, animals died and flames were extinguished
11.
Oxygen
–
Oxygen is a chemical element with symbol O and atomic number 8. It is a member of the group on the periodic table and is a highly reactive nonmetal. By mass, oxygen is the third-most abundant element in the universe, after hydrogen, at standard temperature and pressure, two atoms of the element bind to form dioxygen, a colorless and odorless diatomic gas with the formula O2. This is an important part of the atmosphere and diatomic oxygen gas constitutes 20. 8% of the Earths atmosphere, additionally, as oxides the element makes up almost half of the Earths crust. Most of the mass of living organisms is oxygen as a component of water, conversely, oxygen is continuously replenished by photosynthesis, which uses the energy of sunlight to produce oxygen from water and carbon dioxide. Oxygen is too reactive to remain a free element in air without being continuously replenished by the photosynthetic action of living organisms. Another form of oxygen, ozone, strongly absorbs ultraviolet UVB radiation, but ozone is a pollutant near the surface where it is a by-product of smog. At low earth orbit altitudes, sufficient atomic oxygen is present to cause corrosion of spacecraft, the name oxygen was coined in 1777 by Antoine Lavoisier, whose experiments with oxygen helped to discredit the then-popular phlogiston theory of combustion and corrosion. One of the first known experiments on the relationship between combustion and air was conducted by the 2nd century BCE Greek writer on mechanics, Philo of Byzantium. In his work Pneumatica, Philo observed that inverting a vessel over a burning candle, Philo incorrectly surmised that parts of the air in the vessel were converted into the classical element fire and thus were able to escape through pores in the glass. Many centuries later Leonardo da Vinci built on Philos work by observing that a portion of air is consumed during combustion and respiration, Oxygen was discovered by the Polish alchemist Sendivogius, who considered it the philosophers stone. In the late 17th century, Robert Boyle proved that air is necessary for combustion, English chemist John Mayow refined this work by showing that fire requires only a part of air that he called spiritus nitroaereus. From this he surmised that nitroaereus is consumed in both respiration and combustion, Mayow observed that antimony increased in weight when heated, and inferred that the nitroaereus must have combined with it. Accounts of these and other experiments and ideas were published in 1668 in his work Tractatus duo in the tract De respiratione. Robert Hooke, Ole Borch, Mikhail Lomonosov, and Pierre Bayen all produced oxygen in experiments in the 17th and the 18th century but none of them recognized it as a chemical element. This may have been in part due to the prevalence of the philosophy of combustion and corrosion called the phlogiston theory, which was then the favored explanation of those processes. Established in 1667 by the German alchemist J. J. Becher, one part, called phlogiston, was given off when the substance containing it was burned, while the dephlogisticated part was thought to be its true form, or calx. The fact that a substance like wood gains overall weight in burning was hidden by the buoyancy of the combustion products
12.
Halogen
–
The halogens or halogen elements are a group in the periodic table consisting of five chemically related elements, fluorine, chlorine, bromine, iodine, and astatine. The artificially created element 117 may also be a halogen, in the modern IUPAC nomenclature, this group is known as group 17. The symbol X is often used generically to refer to any halogen, when halogens react with metals they produce a wide range of salts, including calcium fluoride, sodium chloride, silver bromide and potassium iodide. The group of halogens is the periodic table group that contains elements in three of the four main states of matter at standard temperature and pressure. All of the halogens form acids when bonded to hydrogen, most halogens are typically produced from minerals or salts. The middle halogens, that is chlorine, bromine and iodine, are used as disinfectants. Organobromides are the most important class of flame retardants, elemental halogens are dangerously to potentially lethally toxic. The fluorine mineral fluorospar was known as early as 1529, early chemists realized that fluorine compounds contain an undiscovered element, but were unable to isolate it. In 1869, George Gore, an English chemist, ran a current of electricity through hydrofluoric acid and discovered fluorine, in 1886, Henri Moissan, a chemist in Paris, performed electrolysis on potassium bifluoride dissolved in waterless hydrofluoric acid, and successfully produced fluorine. Hydrochloric acid was known to alchemists and early chemists, However, elemental chlorine was not produced until 1774, when Carl Wilhelm Scheele heated hydrochloric acid with manganese dioxide. Scheele called the element dephlogisticated muriatic acid, which is how chlorine was known for 33 years, in 1807, Humphry Davy investigated chlorine and discovered that it is an actual element. Chlorine was used as a poison gas during World War I, bromine was discovered in the 1820s by Antoine-Jérôme Balard. Balard discovered bromine by passing gas through a sample of brine. He originally proposed the name muride for the new element, iodine was discovered by Bernard Courtois, who was using seaweed ash as part of a process for saltpeter manufacture. Courtois typically boiled the seaweed ash with water to generate potassium chloride, However, in 1811, Courtois added sulfuric acid to his process, and found that his process produced purple fumes that condensed into black crystals. Suspecting that these crystals were a new element, Courtois sent samples to other chemists for investigation, iodine was proven to be a new element by Joseph Gay-Lussac. In 1931, Fred Allison claimed to have discovered element 85 with a machine, and named the element Alabamine. In 1937, Jajendralal De claimed to have discovered element 85 in minerals, and called the element dakine, element 85, now named astatine, was produced successfully in 1940 by Dale R. Corson, K. R
13.
Sodium
–
Sodium is a chemical element with symbol Na and atomic number 11. It is a soft, silvery-white, highly reactive metal, Sodium is an alkali metal, being in group 1 of the periodic table, because it has a single electron in its outer shell that it readily donates, creating a positively charged atom—the Na+ cation. Its only stable isotope is 23Na, the free metal does not occur in nature, but must be prepared from compounds. Sodium is the sixth most abundant element in the Earths crust, Sodium was first isolated by Humphry Davy in 1807 by the electrolysis of sodium hydroxide. Among many other useful compounds, sodium hydroxide is used in soap manufacture, and sodium chloride is a de-icing agent. Sodium is an element for all animals and some plants. Sodium ions are the major cation in the fluid and as such are the major contributor to the ECF osmotic pressure. Loss of water from the ECF compartment increases the sodium concentration, isotonic loss of water and sodium from the ECF compartment decreases the size of that compartment in a condition called ECF hypovolemia. In nerve cells, the charge across the cell membrane enables transmission of the nerve impulse—an action potential—when the charge is dissipated. The melting and boiling points of sodium are lower than those of lithium but higher than those of the alkali metals potassium, rubidium. All of these high-pressure allotropes are insulators and electrides.3 nm, spin-orbit interactions involving the electron in the 3p orbital split the D line into two, at 589.0 and 589.6 nm, hyperfine structures involving both orbitals cause many more lines. Twenty isotopes of sodium are known, but only 23Na is stable, 23Na is created in the carbon-burning process in stars by fusing two carbon atoms together, this requires temperatures above 600 megakelvins and a star of at least three solar masses. Two nuclear isomers have been discovered, the one being 24mNa with a half-life of around 20.2 milliseconds. Sodium atoms have 11 electrons, one more than the stable configuration of the noble gas neon. This process requires so little energy that sodium is oxidized by giving up its 11th electron. In contrast, the ionization energy is very high, because the 10th electron is closer to the nucleus than the 11th electron. As a result, sodium usually forms ionic compounds involving the Na+ cation, the most common oxidation state for sodium is +1. It is generally less reactive than potassium and more reactive than lithium, Sodium metal is highly reducing, with the reduction of sodium ions requiring −2.71 volts, though potassium and lithium have even more negative potentials
14.
Magnesium
–
Magnesium is a chemical element with symbol Mg and atomic number 12. Magnesium is the ninth most abundant element in the universe and it is produced in large, aging stars from the sequential addition of three helium nuclei to a carbon nucleus. When such stars explode as supernovas, much of the magnesium is expelled into the medium where it may recycle into new star systems. Magnesium is the eighth most abundant element in the Earths crust and the fourth most common element in the Earth, making up 13% of the planets mass and it is the third most abundant element dissolved in seawater, after sodium and chlorine. Magnesium occurs naturally only in combination with elements, where it invariably has a +2 oxidation state. The free element can be produced artificially, and is highly reactive, the free metal burns with a characteristic brilliant-white light. The metal is now obtained mainly by electrolysis of magnesium salts obtained from brine, Magnesium is less dense than aluminium, and the alloy is prized for its combination of lightness and strength. Magnesium is the eleventh most abundant element by mass in the body and is essential to all cells. Magnesium ions interact with polyphosphate compounds such as ATP, DNA, hundreds of enzymes require magnesium ions to function. Magnesium compounds are used medicinally as common laxatives, antacids, elemental magnesium is a gray-white lightweight metal, two-thirds the density of aluminium. Magnesium has the lowest melting and the lowest boiling point 1,363 K of all the alkaline earth metals, Magnesium reacts with water at room temperature, though it reacts much more slowly than calcium, a similar group 2 metal. When submerged in water, hydrogen bubbles form slowly on the surface of the metal—though, if powdered, the reaction occurs faster with higher temperatures. Magnesiums reversible reaction with water can be harnessed to store energy, Magnesium also reacts exothermically with most acids such as hydrochloric acid, producing the metal chloride and hydrogen gas, similar to the HCl reaction with aluminium, zinc, and many other metals. Magnesium is highly flammable, especially when powdered or shaved into thin strips, flame temperatures of magnesium and magnesium alloys can reach 3,100 °C, although flame height above the burning metal is usually less than 300 mm. Once ignited, such fires are difficult to extinguish, with combustion continuing in nitrogen, carbon dioxide, Magnesium may also be used as an igniter for thermite, a mixture of aluminium and iron oxide powder that ignites only at a very high temperature. When burning in air, magnesium produces a light that includes strong ultraviolet wavelengths. Magnesium powder was used for illumination in the early days of photography. Later, magnesium filament was used in electrically ignited single-use photography flashbulbs, Magnesium powder is used in fireworks and marine flares where a brilliant white light is required
15.
Lewis structure
–
Lewis structures are diagrams that show the bonding between atoms of a molecule and the lone pairs of electrons that may exist in the molecule. A Lewis structure can be drawn for any covalently bonded molecule, the Lewis structure was named after Gilbert N. Lewis, who introduced it in his 1916 article The Atom and the Molecule. Lewis structures extend the concept of the electron dot diagram by adding lines between atoms to represent shared pairs in a chemical bond, Lewis structures show each atom and its position in the structure of the molecule using its chemical symbol. Lines are drawn between atoms that are bonded to one another, excess electrons that form lone pairs are represented as pairs of dots, and are placed next to the atoms. Hydrogen can only form bonds which share just two electrons, while transition metals often conform to a duodectet rule, the total number of electrons represented in a Lewis structure is equal to the sum of the numbers of valence electrons on each individual atom. Non-valence electrons are not represented in Lewis structures, once the total number of available electrons has been determined, electrons must be placed into the structure. They should be placed initially as lone pairs, one pair of dots for each pair of electrons available. Lone pairs should initially be placed on outer atoms until each outer atom has eight electrons in bonding pairs and lone pairs, when in doubt, lone pairs should be placed on more electronegative atoms first. Once all lone pairs are placed, atoms—especially the central atoms—may not have an octet of electrons, in this case, the atoms must form a double bond, a lone pair of electrons is moved to form a second bond between the two atoms. As the bonding pair is shared between the two atoms, the atom that originally had the pair still has an octet, the other atom now has two more electrons in its valence shell. Lewis structures for polyatomic ions may be drawn by the same method, when counting electrons, negative ions should have extra electrons placed in their Lewis structures, positive ions should have fewer electrons than an uncharged molecule. When the Lewis structure of an ion is written, the structure is placed in brackets. The rest of the electrons just go to all the other atoms octets. Another simple and general procedure to write Lewis structures and resonance forms has been proposed and it has uses in determining possible electron re-configuration when referring to reaction mechanisms, and often results in the same sign as the partial charge of the atom, with exceptions. In general, the charge of an atom can be calculated using the following formula, assuming non-standard definitions for the markup used, C f = N v − U e − B n 2 where. N v represents the number of electrons in a free atom of the element. U e represents the number of unshared electrons on the atom, B n represents the total number of electrons in bonds the atom has with another. The formal charge of an atom is computed as the difference between the number of electrons that a neutral atom would have and the number of electrons that belong to it in the Lewis structure
16.
Covalent bond
–
A covalent bond, also called a molecular bond, is a chemical bond that involves the sharing of electron pairs between atoms. These electron pairs are known as shared pairs or bonding pairs, for many molecules, the sharing of electrons allows each atom to attain the equivalent of a full outer shell, corresponding to a stable electronic configuration. Covalent bonding includes many kinds of interactions, including σ-bonding, π-bonding, metal-to-metal bonding, agostic interactions, bent bonds, the term covalent bond dates from 1939. In the molecule H2, the atoms share the two electrons via covalent bonding. Covalency is greatest between atoms of similar electronegativities, thus, covalent bonding does not necessarily require that the two atoms be of the same elements, only that they be of comparable electronegativity. Covalent bonding that entails sharing of electrons more than two atoms is said to be delocalized. The term covalence in regard to bonding was first used in 1919 by Irving Langmuir in a Journal of the American Chemical Society article entitled The Arrangement of Electrons in Atoms and Molecules. Langmuir wrote that we shall denote by the term covalence the number of pairs of electrons that an atom shares with its neighbors. The idea of covalent bonding can be traced several years before 1919 to Gilbert N. Lewis and he introduced the Lewis notation or electron dot notation or Lewis dot structure, in which valence electrons are represented as dots around the atomic symbols. Pairs of electrons located between atoms represent covalent bonds, multiple pairs represent multiple bonds, such as double bonds and triple bonds. An alternative form of representation, not shown here, has bond-forming electron pairs represented as solid lines, Lewis proposed that an atom forms enough covalent bonds to form a full outer electron shell. In the diagram of methane shown here, the atom has a valence of four and is, therefore, surrounded by eight electrons, four from the carbon itself. Each hydrogen has a valence of one and is surrounded by two electrons – its own one electron plus one from the carbon, walter Heitler and Fritz London are credited with the first successful quantum mechanical explanation of a chemical bond in 1927. Their work was based on the valence bond model, which assumes that a bond is formed when there is good overlap between the atomic orbitals of participating atoms. Atomic orbitals have specific directional properties leading to different types of covalent bonds, sigma bonds are the strongest covalent bonds and are due to head-on overlapping of orbitals on two different atoms. A single bond is usually a σ bond, pi bonds are weaker and are due to lateral overlap between p orbitals. A double bond between two given atoms consists of one σ and one π bond, and a bond is one σ. Covalent bonds are also affected by the electronegativity of the atoms which determines the chemical polarity of the bond
17.
Ionic bonding
–
Ionic bonding is a type of chemical bond that involves the electrostatic attraction between oppositely charged ions, and is the primary interaction occurring in ionic compounds. The ions are atoms that have gained one or more electrons and this transfer of electrons is known as electrovalence in contrast to covalence. In the simplest case, the cation is an atom and the anion is a nonmetal atom. In simpler words, a bond is the transfer of electrons from a metal to a non-metal in order for both atoms to obtain a full valence shell. Bonds with partially ionic and partially covalent character are called polar covalent bonds, Ionic compounds conduct electricity when molten or in solution, typically as a solid. Ionic compounds generally have a melting point, depending on the charge of the ions they consist of. The higher the charges the stronger the forces and the higher the melting point. They also tend to be soluble in water, here, the opposite trend roughly holds, the weaker the cohesive forces, the greater the solubility. Atoms that have an almost full or almost empty valence shell tend to be very reactive, atoms that are strongly electronegative often have only one or two empty orbitals in their valence shell, and frequently bond with other molecules or gain electrons to form anions. Atoms that are weakly electronegative have relatively few valence electrons that can easily be shared with atoms that are strongly electronegative, as a result, weakly electronegative atoms tend to distort their electrons cloud and form cations. Ionic bonding can result from a reaction when atoms of an element, whose ionization energy is low. In doing so, cations are formed, the atom of another element, whose electron affinity is positive, then accepts the electron, again to attain a stable electron configuration, and after accepting electron the atom becomes an anion. Typically, the electron configuration is one of the noble gases for elements in the s-block and the p-block. The electrostatic attraction between the anions and cations leads to the formation of a solid with a lattice in which the ions are stacked in an alternating fashion. In such a lattice, it is not possible to distinguish discrete molecular units. However, the ions themselves can be complex and form molecular ions like the acetate anion or the ammonium cation, for example, common table salt is sodium chloride. When sodium and chlorine are combined, the sodium atoms each lose an electron, forming cations, and these ions are then attracted to each other in a 1,1 ratio to form sodium chloride. For compounds that are transitional to the alloys and possess mixed ionic and metallic bonding, many sulfides, e. g. do form non-stoichiometric compounds
18.
Metal
–
A metal is a material that is typically hard, opaque, shiny, and has good electrical and thermal conductivity. Metals are generally malleable—that is, they can be hammered or pressed permanently out of shape without breaking or cracking—as well as fusible and ductile, about 91 of the 118 elements in the periodic table are metals, the others are nonmetals or metalloids. Some elements appear in both metallic and non-metallic forms, astrophysicists use the term metal to collectively describe all elements other than hydrogen and helium, the simplest two, in a star. The star fuses smaller atoms, mostly hydrogen and helium, to larger ones over its lifetime. In that sense, the metallicity of an object is the proportion of its matter made up of all chemical elements. Many elements and compounds that are not normally classified as metals become metallic under high pressures, the atoms of metallic substances are typically arranged in one of three common crystal structures, namely body-centered cubic, face-centered cubic, and hexagonal close-packed. In bcc, each atom is positioned at the center of a cube of eight others, in fcc and hcp, each atom is surrounded by twelve others, but the stacking of the layers differs. Some metals adopt different structures depending on the temperature, atoms of metals readily lose their outer shell electrons, resulting in a free flowing cloud of electrons within their otherwise solid arrangement. This provides the ability of metallic substances to easily transmit heat, while this flow of electrons occurs, the solid characteristic of the metal is produced by electrostatic interactions between each atom and the electron cloud. This type of bond is called a metallic bond, Metals are usually inclined to form cations through electron loss, reacting with oxygen in the air to form oxides over various timescales. Examples,4 Na + O2 →2 Na2O2 Ca + O2 →2 CaO4 Al +3 O2 →2 Al2O3, the transition metals are slower to oxidize because they form a passivating layer of oxide that protects the interior. Others, like palladium, platinum and gold, do not react with the atmosphere at all, some metals form a barrier layer of oxide on their surface which cannot be penetrated by further oxygen molecules and thus retain their shiny appearance and good conductivity for many decades. The oxides of metals are generally basic, as opposed to those of nonmetals, exceptions are largely oxides with very high oxidation states such as CrO3, Mn2O7, and OsO4, which have strictly acidic reactions. Painting, anodizing or plating metals are good ways to prevent their corrosion, however, a more reactive metal in the electrochemical series must be chosen for coating, especially when chipping of the coating is expected. Water and the two form an electrochemical cell, and if the coating is less reactive than the coatee. Metals in general have high conductivity, high thermal conductivity. Typically they are malleable and ductile, deforming under stress without cleaving, in terms of optical properties, metals are shiny and lustrous. Sheets of metal beyond a few micrometres in thickness appear opaque, although most metals have higher densities than most nonmetals, there is wide variation in their densities, lithium being the least dense solid element and osmium the densest
19.
Nonmetal
–
In chemistry, a nonmetal is a chemical element that mostly lacks metallic attributes. Seventeen elements are classified as nonmetals, most are gases, one is a liquid. Moving rightward across the standard form of the table, nonmetals adopt structures that have progressively fewer nearest neighbours. Polyatomic nonmetals have structures with either three nearest neighbours, as is the case with carbon, or two nearest neighbours in the case of sulfur, diatomic nonmetals, such as hydrogen, have one nearest neighbour, and the monatomic noble gases, such as helium, have none. This gradual fall in the number of nearest neighbours is associated with a reduction in metallic character, the distinction between the three categories of nonmetals, in terms of receding metallicity is not absolute. Boundary overlaps occur as outlying elements in each category show less-distinct, living organisms are also composed almost entirely of nonmetals, and nonmetals form many more compounds than metals. There is no definition of a nonmetal. They show more variability in their properties than metals do, the following are some of the chief characteristics of nonmetals. The elements generally classified as nonmetals include one element in group 1, one in group 14, the distinction between nonmetals and metals is by no means clear. Nonmetals have structures in each atom usually forms bonds with nearest neighbours. Each atom is able to complete its valence shell and attain a stable noble gas configuration. Exceptions to the rule occur with hydrogen, carbon, nitrogen and oxygen, atoms of the latter three elements are sufficiently small such that they are able to form alternative bonding structures, with fewer nearest neighbours. Thus, carbon is able to form its layered graphite structure, the larger size of the remaining non-noble nonmetals weakens their capacity to form multiple bonds and they instead form two or more single bonds to two or more different atoms. Sulfur, for example, forms an eight-membered molecule in which the atoms are arranged in a ring, a similar pattern occurs more generally, at the level of the entire periodic table, in comparing metals and nonmetals. There is a transition from metallic bonding among the metals on the left of the table through to covalent or Van der Waals bonding among the nonmetals on the right of the table, metallic bonding tends to involve close-packed centrosymmetric structures with a high number of nearest neighbours. Post-transition metals and metalloids, sandwiched between the metals and the nonmetals, tend to have more complex structures with an intermediate number of nearest neighbours. Nonmetallic bonding, towards the right of the table, features open-packed directional structures with fewer or zero nearest neighbours, as is the case with the major categories of metals, metalloids and nonmetals, there is some variation and overlapping of properties within and across each category of nonmetal. Among the polyatomic nonmetals, carbon, phosphorus and selenium—which border the metalloids—begin to show some metallic character, sulfur, is the least metallic of the polyatomic nonmetals but even here shows some discernible metal-like character
20.
Chlorine
–
Chlorine is a chemical element with symbol Cl and atomic number 17. The second-lightest of the halogens, it appears between fluorine and bromine in the table and its properties are mostly intermediate between them. Chlorine is a gas at room temperature. It is an extremely reactive element and a strong oxidising agent, among the elements, it has the highest electron affinity, the most common compound of chlorine, sodium chloride, has been known since ancient times. Around 1630, chlorine gas was first synthesised in a chemical reaction, Carl Wilhelm Scheele wrote a description of chlorine gas in 1774, supposing it to be an oxide of a new element. In 1809, chemists suggested that the gas might be an element, and this was confirmed by Sir Humphry Davy in 1810. Because of its reactivity, all chlorine in the Earths crust is in the form of ionic chloride compounds. It is the second-most abundant halogen and twenty-first most abundant chemical element in Earths crust and these crustal deposits are nevertheless dwarfed by the huge reserves of chloride in seawater. Elemental chlorine is produced from brine by electrolysis. The high oxidising potential of chlorine led to the development of commercial bleaches and disinfectants. As a common disinfectant, elemental chlorine and chlorine-generating compounds are used directly in swimming pools to keep them clean. Elemental chlorine at high concentrations is extremely dangerous and poisonous for all living organisms, in the form of chloride ions, chlorine is necessary to all known species of life. Other types of compounds are rare in living organisms. In the upper atmosphere, chlorine-containing organic molecules such as chlorofluorocarbons have been implicated in ozone depletion, small quantities of elemental chlorine are generated by oxidation of chloride to hypochlorite in neutrophils as part of the immune response against bacteria. Its importance in food was very well known in antiquity and was sometimes used as payment for services for Roman generals. Around 1630, chlorine was recognized as a gas by the Flemish chemist, the element was first studied in detail in 1774 by Swedish chemist Carl Wilhelm Scheele, and he is credited with the discovery. He called it dephlogisticated muriatic acid air since it is a gas and he failed to establish chlorine as an element, mistakenly thinking that it was the oxide obtained from the hydrochloric acid. He named the new element within this oxide as muriaticum, in 1809, Joseph Louis Gay-Lussac and Louis-Jacques Thénard tried to decompose dephlogisticated muriatic acid air by reacting it with charcoal to release the free element muriaticum
21.
Electron affinity
–
X + e− → X− + energy In solid state physics, the electron affinity for a surface is defined somewhat differently. This property is measured for atoms and molecules in the state only. A list of the electron affinities was used by Robert S. Mulliken to develop an electronegativity scale for atoms, equal to the average of the electron affinity, other theoretical concepts that use electron affinity include electronic chemical potential and chemical hardness. Another example, a molecule or atom that has a positive value of electron affinity than another is often called an electron acceptor. Together they may undergo charge-transfer reactions, to use electron affinities properly, it is essential to keep track of sign. For any reaction that releases energy, the change ΔE in total energy has a negative value, electron capture for almost all non-noble gas atoms involves the release of energy and thus are exothermic. The positive values that are listed in tables of Eea are amounts or magnitudes and it is the word, released within the definition energy released that supplies the negative sign to ΔE. Confusion arises in mistaking Eea for a change in energy, ΔE, the relation between the two is Eea = −ΔE. However, if the assigned to Eea is negative, the negative sign implies a reversal of direction. In this case, the capture is an endothermic process. Negative values typically arise for the capture of a second electron, since almost all detachments an amount of energy listed on the table, those detachment reactions are endothermic, or ΔE >0. Although Eea varies greatly across the table, some patterns emerge. Generally, nonmetals have more positive Eea than metals, atoms whose anions are more stable than neutral atoms have a greater Eea. Chlorine most strongly attracts extra electrons, mercury most weakly attracts an extra electron, the electron affinities of the noble gases have not been conclusively measured, so they may or may not have slightly negative values. Eea generally increases across a period in the periodic table, a trend of decreasing Eea going down the groups in the periodic table might be expected. The additional electron will be entering an orbital farther away from the nucleus, since this electron is farther from the nucleus it is less attracted to the nucleus and would release less energy when added. However, a counterexample to this trend can be found in Group 2. Thus, electron affinity follows the trend of electronegativity but not the up-down trend
22.
Ionization energy
–
The ionization energy is qualitatively defined as the amount of energy required to remove the most loosely bound electron, the valence electron, of an isolated gaseous atom to form a cation. Generally, the closer the electrons are to the nucleus of the atom, the units for ionization energy are different in physics and chemistry. In physics, the unit is the amount of required to remove a single electron from a single atom or molecule. Comparison of IEs of atoms in the periodic table reveals two patterns, IEs generally increase as one moves from left to right within a given period, IEs generally decrease as one moves down a given group. The nth ionization energy refers to the amount of required to remove an electron from the species with a charge of. However this term is now considered obsolete, type of orbital ionized as the atom having stable electronic configuration has less tendency to lose electrons and consequently has high ionization energy. Occupancy of the orbital matters as if the orbital is half or completely filled then it is harder to remove electrons Further, thus, the net energy may equals energy of attraction minus the repulsion energy. In aqueous solution perhaps the latent heat of the electron should be taken into account, generally, the th ionization energy is larger than the nth ionization energy. Electrons removed from more highly charged ions of a particular element experience greater forces of attraction, thus. Both of these factors increase the ionization energy. Some values for elements of the period are given in the following table. That electron is closer to the nucleus than the previous 3s electron. Ionization energy is also a trend within the periodic table organization. Moving left to right within a period, or upward within a group, as the nuclear charge of the nucleus increases across the period, the atomic radius decreases and the electron cloud becomes closer towards the nucleus. Atomic ionization energy can be predicted by an analysis using electrostatic potential, consider an electron of charge -e and an atomic nucleus with charge +Ze, where Z is the number of protons in the nucleus. According to the Bohr model, if the electron were to approach and bond with the atom, it would come to rest at a certain radius a. The energy required for the electron to climb out and leave the atom is, E = e V = Z e 2 a This analysis is incomplete, as it leaves the distance a as an unknown variable. It can be more rigorous by assigning to each electron of every chemical element a characteristic distance
23.
Mole (unit)
–
The mole is the unit of measurement in the International System of Units for amount of substance. This number is expressed by the Avogadro constant, which has a value of 6. 022140857×1023 mol−1, the mole is one of the base units of the SI, and has the unit symbol mol. The mole is used in chemistry as a convenient way to express amounts of reactants and products of chemical reactions. For example, the chemical equation 2 H2 + O2 →2 H2O implies that 2 moles of dihydrogen and 1 mole of dioxygen react to form 2 moles of water. The mole may also be used to express the number of atoms, ions, the concentration of a solution is commonly expressed by its molarity, defined as the number of moles of the dissolved substance per litre of solution. For example, the relative molecular mass of natural water is about 18.015, therefore. The term gram-molecule was formerly used for essentially the same concept, the term gram-atom has been used for a related but distinct concept, namely a quantity of a substance that contains Avogadros number of atoms, whether isolated or combined in molecules. Thus, for example,1 mole of MgBr2 is 1 gram-molecule of MgBr2 but 3 gram-atoms of MgBr2, in honor of the unit, some chemists celebrate October 23, which is a reference to the 1023 scale of the Avogadro constant, as Mole Day. Some also do the same for February 6 and June 2, thus, by definition, one mole of pure 12C has a mass of exactly 12 g. It also follows from the definition that X moles of any substance will contain the number of molecules as X moles of any other substance. The mass per mole of a substance is called its molar mass, the number of elementary entities in a sample of a substance is technically called its amount. Therefore, the mole is a convenient unit for that physical quantity, one can determine the chemical amount of a known substance, in moles, by dividing the samples mass by the substances molar mass. Other methods include the use of the volume or the measurement of electric charge. The mass of one mole of a substance depends not only on its molecular formula, since the definition of the gram is not mathematically tied to that of the atomic mass unit, the number NA of molecules in a mole must be determined experimentally. The value adopted by CODATA in 2010 is NA =6. 02214129×1023 ±0. 00000027×1023, in 2011 the measurement was refined to 6. 02214078×1023 ±0. 00000018×1023. The number of moles of a sample is the sample mass divided by the mass of the material. The history of the mole is intertwined with that of mass, atomic mass unit, Avogadros number. The first table of atomic mass was published by John Dalton in 1805
24.
Lattice energy
–
The lattice energy of a crystalline solid is usually defined as the energy of formation of the crystal from infinitely-separated ions and as such is invariably positive. However, the value of the energy may either be derived theoretically from electrostatics or from a thermodynamic cycling reaction. The concept of energy was originally developed for rocksalt-structured and sphalerite-structured compounds like NaCl and ZnS. In the case of NaCl, lattice energy is the released by the reaction Na+ + Cl− → NaCl which would amount to -786 kJ/mol. Some older textbooks define lattice energy with the sign, i. e. the energy required to convert the crystal into infinitely separated gaseous ions in vacuum. Following this convention, the energy of NaCl would be +786 kJ/mol. Moreover, factors such as packing of ions doesnt matter efficiently, the Born–Landé equation gives a reasonable fit to the lattice energy. The Kapustinskii equation can be used as a way of deriving lattice energies where high precision is not required. For ionic compounds with ions occupying lattice sites with crystallographic point groups C1, C1h, Cn or Cnv the concept of the lattice energy, as an example, one may consider the case of iron-pyrite FeS2, where sulfur ions occupy lattice site of point symmetry group C3. The lattice energy defining reaction then reads Fe2+ +2 pol S− → FeS2 where pol S− stands for the polarized and it is the gaseous ions that combine 2. By the current definition, the energy is always exothermic. Bond energy Born–Haber cycle Chemical bond Madelung constant Ionic conductivity
25.
Alfred Werner
–
Alfred Werner was a Swiss chemist who was a student at ETH Zurich and a professor at the University of Zurich. He won the Nobel Prize in Chemistry in 1913 for proposing the octahedral configuration of transition metal complexes, Werner developed the basis for modern coordination chemistry. He was the first inorganic chemist to win the Nobel prize, Werner was born in 1866 in Mulhouse, Alsace. He was raised as Roman Catholic and he went to Switzerland to study chemistry at the Swiss Federal Institute in Zurich where he obtained his doctorate in 1890 at the same institution. After postdoctoral study in Paris, he returned to the Swiss Federal Institute to teach, the same year he became a Swiss citizen. In 1893, Werner was the first to propose correct structures for coordination compounds containing complex ions, for example, it was known that cobalt forms a complex hexamminecobalt chloride, with formula CoCl3•6NH3, but the nature of the association indicated by the dot was mysterious. Werner proposed the structure Cl3, with the Co3+ ion surrounded by six NH3 at the vertices of an octahedron, later, magnetic susceptibility analysis was also used to confirm Werners proposal for the chemical nature of CoCl3•6NH3. For complexes with more than one type of ligand, Werner succeeded in explaining the number of isomers observed, for example, he explained the existence of two tetrammine isomers, Co4Cl3, one green and one purple. Werner proposed that these are two isomers of formula Cl, with one Cl− ion dissociated as confirmed by conductivity measurements. The Co atom is surrounded by four NH3 and two Cl ligands at the vertices of an octahedron, the green isomer is trans with the two Cl ligands at opposite vertices, and the purple is cis with the two Cl at adjacent vertices. Werner also prepared complexes with optical isomers, and in 1914 he reported the first synthetic chiral compound lacking carbon, before Werner, chemists defined the valence of an element as the number of its bonds without distinguishing different types of bond. This secondary valence of 6 he referred to as the number which he defined as the number of molecules directly linked to the central metal atom. In other complexes he found coordination numbers of 4 or 8 and this rule was used later in 1916 when Gilbert N. Lewis formulated the “octet rule” in his cubical atom theory. In modern terminology Werners primary valence corresponds to the oxidation state, the Co-Cl bonds are now classed as ionic, and each Co-N bond is a coordinate covalent bond between the Lewis acid Co3+ and the Lewis base NH3. Fischer, Jena 1904 Digital edition by the University and State Library Düsseldorf Biography at Nobelprize
26.
Coordination number
–
In chemistry, crystallography, and materials science the coordination number of a central atom in a molecule or crystal is the number of its near neighbors. This number is determined differently for molecules than for crystals. For molecules and polyatomic ions the coordination number of an atom is determined by counting the other atoms to which it is bonded. For example, − has Cr3+ as its central cation, and has a number of 6. However the solid-state structures of crystals often have less clearly defined bonds, the simplest method is one used in materials science. The usual value of the number for a given structure refers to an atom in the interior of a crystal lattice with neighbors in all directions. In chemistry, coordination number, defined originally in 1893 by Alfred Werner, is the number of neighbors of a central atom in a molecule or ion. Although a carbon atom has four chemical bonds in most stable molecules, in effect we count the first bond to each neighboring atom, but not the other bonds. In coordination complexes, only the first or sigma bond between each ligand and the central atom counts, but not any pi bonds to the same ligands. In tungsten hexacarbonyl, W6, the number of tungsten is counted as six although pi as well as sigma bonding is important in such metal carbonyls. The most common number for d-block transition metal complexes is 6. The observed range is 2 to 9, metals in the f-block can accommodate higher coordination number due to their greater ionic radii and availability of more orbitals for bonding. Coordination numbers of 8 to 12 are commonly observed for f-block elements, for example, with bidentate nitrate ions as ligands, CeIV and ThIV form the 12-coordinate ions 2− and 2−. When the surrounding ligands are much smaller than the central atom, one computational chemistry study predicted a particularly stable PbHe2+15 ion composed of a central lead ion coordinated with no fewer than 15 helium atoms. At the opposite extreme, steric shielding can give rise to low coordination numbers. An extremely rare instance of a metal adopting a number of 1 occurs in the terphenyl-based arylthallium complex 2, 6-Tipp2C6H3Tl. In ferrocene the hapticity, η, of each cyclopentadienide anion is five, there are various ways of assigning the contribution made to the coordination number of the central iron atom by each cyclopentadienide ligand. The contribution could be assigned as one since there is one ligand, or as five since there are five neighbouring atoms, normally the count of electron pairs is taken
27.
Richard Abegg
–
Richard Wilhelm Heinrich Abegg was a German chemist and pioneer of valence theory. He proposed that the difference of the positive and negative valence of an element tends to be eight. This has come to be known as Abeggs rule and he was a gas balloon enthusiast, which caused his death at the age of 41 when he crashed in his balloon in Silesia. Abegg received his PhD on July 19,1891 as the student of August Wilhelm von Hofmann at the University of Berlin. Abegg learned organic chemistry from Hofmann, but one year after finishing his PhD degree he began researching physical chemistry while studying with Friedrich Wilhelm Ostwald in Leipzig, Germany. Abegg later served as assistant to Walther Nernst at the University of Göttingen. Abegg discovered the theory of freezing-point depression and anticipated Gilbert Newton Lewiss octet rule by revealing that the lowest and highest oxidation states of elements differ by eight. He researched many topics in chemistry, including freezing points, the dielectric constant of ice, osmotic pressures, oxidation potentials. Richard Abegg was the son of Wilhelm Abegg and Margarete Friedenthal and he had a brother, Wilhelm Abegg, who became the Prussian Secretary of State. After attending Wilhelm High School in Berlin, Abegg studied organic chemistry at the University of Kiel and he then attended the University of Berlin, from which he received his doctorate as the student of August Wilhelm von Hofmann. In 1895, he married Line Simon, who also a ballooning enthusiast. Abegg occupied himself with photography and balloon excursions and he was the initiator and chairperson of the Silesian Club for Aeronautics in Breslau. Furthermore, he had a function with the presidency of the German Air Sailors Association. During school, Abegg fulfilled his duties in the military, in 1891, he became an officer of the German Reserves. In 1900, he became an Oberleutnant in the Reserves in the 9th Regiment of Hussars, during this year, he made his first flight in a balloon, for military purposes. Balloon flights became a frequent pastime of both Abegg and his wife and he made many scientific observations during his subsequent flights, which were never published. In 1894, Abegg worked as an assistant to Walther Nernst, one of the founders of chemistry and, at the time. In 1897, he took a position as a professor of chemistry at the University of Breslau, two years later, Abegg became a Privatdozent at the Wrocław University of Technology in Wrocław, Poland
28.
Valence (chemistry)
–
In chemistry, the valence or valency of an element is a measure of its combining power with other atoms when it forms chemical compounds or molecules. The concept of valence was developed in the half of the 19th century and was successful in explaining the molecular structure of inorganic and organic compounds. The combining power or affinity of an atom of an element was determined by the number of atoms that it combined with. In methane, carbon has a valence of 4, in ammonia, nitrogen has a valence of 3, in water, oxygen has a valence of 2, and in hydrogen chloride, chlorine has a valence of 1. Chlorine, as it has a valence of one, can be substituted for hydrogen, so phosphorus has a valence of 5 in phosphorus pentachloride, PCl5. Valence diagrams of a compound represent the connectivity of the elements, examples are, Valence only describes connectivity, it does not describe the geometry of molecular compounds, or what are now known to be ionic compounds or giant covalent structures. A line between atoms does not represent a pair of electrons as it does in Lewis diagrams and this definition differs from the IUPAC definition as an element can be said to have more than one valence. It is in manner, according to Frankland, that their affinities are best satisfied. Following these examples and postulates, Frankland declares how obvious it is that This “combining power” was afterwards called quantivalence or valency. In 1857 August Kekulé proposed fixed valences for many elements, such as 4 for carbon, and used them to propose structural formulas for many organic molecules, which are still accepted today. Most 19th-century chemists defined the valence of an element as the number of its bonds without distinguishing different types of valence or of bond. For main-group elements, in 1904 Richard Abegg considered positive and negative valences, in 1916, Gilbert N. Lewis explained valence and chemical bonding in terms of a tendency of atoms to achieve a stable octet of 8 valence-shell electrons. According to Lewis, covalent bonding leads to octets by the sharing of electrons, the term covalence is attributed to Irving Langmuir, who stated in 1919 that the number of pairs of electrons which any given atom shares with the adjacent atoms is called the covalence of that atom. The prefix co- means together, so that a co-valent bond means that the share a valence. Pauling also considered hypervalent molecules, in which main-group elements have apparent valences greater than the maximal of 4 allowed by the octet rule. Similar calculations on transition-metal molecules show that the role of p orbitals is minor, so that one s, for elements in the main groups of the periodic table, the valence can vary between 1 and 7. Many elements have a common valence related to their position in the periodic table, the Latin/Greek prefixes uni-/mono-, bi-/di-, ter-/tri-, quadri-/tetra- and quinque-/penta- are used to describe ions in the charge states 1,2,3,4,5 respectively. Polyvalence or multivalence refers to species that are not restricted to a number of valence bonds
29.
Oxidation state
–
The oxidation state, often called the oxidation number, is an indicator of the degree of oxidation of an atom in a chemical compound. This is never exactly true for real bonds, the term oxidation was first used by Lavoisier to signify reaction of a substance with oxygen. Much later, it was realized that the substance, upon being oxidized, loses electrons, oxidation states are typically represented by integers. In some cases, the oxidation state of an element is a fraction. It is predicted that even a +10 oxidation state may be achieveable by platinum in the platinum tetroxide dication, the increase in oxidation state of an atom, through a chemical reaction, is known as an oxidation, a decrease in oxidation state is known as a reduction. Such reactions involve the transfer of electrons, a net gain in electrons being a reduction. For pure elements, the state is zero. There are various methods for determining oxidation states, in inorganic nomenclature, the oxidation state is determined and expressed as an oxidation number, and is represented by a Roman numeral placed after the element name. In coordination chemistry, oxidation number is defined differently from oxidation state, a “Comprehensive definition of oxidation state ” has been published with free access. This article is a distillation of an IUPAC technical report Toward a comprehensive definition of oxidation state. Exceptions to this are that hydrogen has a state of −1 in hydrides of active metals, e. g. LiH. There are two different methods for determining the state of elements in chemical compounds. First a method based on the rules in the IUPAC definition, second, a method based on the relative electronegativity of the elements in the compound, where the more electronegative element is assumed to take the negative charge. Any pure element—even if it forms diatomic molecules like chlorine —has an oxidation state of zero, examples of this are Cu or O2. For monatomic ions, the state is the same as the charge of the ion. For example, the anion has an oxidation state of −2. The sum of oxidation states for all atoms in a molecule or polyatomic ion is equal to the charge of the molecule or ion, thus, the oxidation state of one element can be calculated from the oxidation states of the other elements. An application of this rule is that the sum of the states of all atoms in a neutral molecule must be zero
30.
Chemical element
–
A chemical element or element is a species of atoms having the same number of protons in their atomic nuclei. There are 118 elements that have identified, of which the first 94 occur naturally on Earth with the remaining 24 being synthetic elements. There are 80 elements that have at least one stable isotope and 38 that have exclusively radioactive isotopes, iron is the most abundant element making up Earth, while oxygen is the most common element in the Earths crust. Chemical elements constitute all of the matter of the universe. The two lightest elements, hydrogen and helium, were formed in the Big Bang and are the most common elements in the universe. The next three elements were formed mostly by cosmic ray spallation, and are rarer than those that follow. Formation of elements with from 6 to 26 protons occurred and continues to occur in main sequence stars via stellar nucleosynthesis, the high abundance of oxygen, silicon, and iron on Earth reflects their common production in such stars. The term element is used for atoms with a number of protons as well as for a pure chemical substance consisting of a single element. A single element can form multiple substances differing in their structure, when different elements are chemically combined, with the atoms held together by chemical bonds, they form chemical compounds. Only a minority of elements are found uncombined as relatively pure minerals, among the more common of such native elements are copper, silver, gold, carbon, and sulfur. All but a few of the most inert elements, such as gases and noble metals, are usually found on Earth in chemically combined form. While about 32 of the elements occur on Earth in native uncombined forms. For example, atmospheric air is primarily a mixture of nitrogen, oxygen, and argon, the history of the discovery and use of the elements began with primitive human societies that found native elements like carbon, sulfur, copper and gold. Later civilizations extracted elemental copper, tin, lead and iron from their ores by smelting, using charcoal, alchemists and chemists subsequently identified many more, almost all of the naturally occurring elements were known by 1900. Save for unstable radioactive elements with short half-lives, all of the elements are available industrially, almost all other elements found in nature were made by various natural methods of nucleosynthesis. On Earth, small amounts of new atoms are produced in nucleogenic reactions, or in cosmogenic processes. Of the 94 naturally occurring elements, those with atomic numbers 1 through 82 each have at least one stable isotope, Isotopes considered stable are those for which no radioactive decay has yet been observed. Elements with atomic numbers 83 through 94 are unstable to the point that radioactive decay of all isotopes can be detected, the very heaviest elements undergo radioactive decay with half-lives so short that they are not found in nature and must be synthesized
31.
Gilbert N. Lewis
–
Lewis successfully contributed to thermodynamics, photochemistry, and isotope separation, and is also known for his concept of acids and bases. G. N. Lewis was born in 1875 in Weymouth, several years later, he became the Dean of the college of Chemistry at Berkeley, where he spent the rest of his life. As a professor, he incorporated thermodynamic principles into the chemistry curriculum and he began measuring the free energy values related to several chemical processes, both organic and inorganic. In 1916, he proposed his theory of bonding and added information about electrons in the periodic table of the chemical elements. In 1933, he started his research on isotope separation, Lewis worked with hydrogen and managed to purify a sample of heavy water. He then came up with his theory of acids and bases, in 1926, Lewis coined the term photon for the smallest unit of radiant energy. He was a brother in Alpha Chi Sigma, the chemistry fraternity. Though he was nominated 41 times, G. N. Lewis never won the Nobel Prize in Chemistry. On March 23,1946, Lewis was found dead in his Berkeley laboratory where he had working with hydrogen cyanide. After Lewis death, his children followed their fathers career in chemistry, Lewis was born in 1875 and raised in Weymouth, Massachusetts, where there exists a street named for him, G. N. Additionally, the wing of the new Weymouth High School Chemistry department has named in his honor. Lewis received his education at home from his parents, Frank Wesley Lewis, a lawyer of independent character. He read at age three and was intellectually precocious, in 1884 his family moved to Lincoln, Nebraska, and in 1889 he received his first formal education at the university preparatory school. In 1893, after two years at the University of Nebraska, Lewis transferred to Harvard University, where he obtained his B. S. in 1896. After a year of teaching at Harvard, Lewis took a fellowship to Germany, the center of physical chemistry. While working in Nernsts lab, Nernst and Lewis apparently developed a lifelong enmity, in the following years, Lewis started to criticize and denounce his former teacher on many occasions, calling Nernsts work on his heat theorem a regrettable episode in the history of chemistry. A friend of Nernsts, Wilhelm Palmær, was a member of the Nobel Chemistry Committee, after his stay in Nernsts lab, Lewis returned to Harvard in 1901 as an instructor for three more years. He was appointed instructor in thermodynamics and electrochemistry, in 1904 Lewis was granted a leave of absence and became Superintendent of Weights and Measures for the Bureau of Science in Manila, Philippines
32.
Cubical atom
–
The cubical atom was an early atomic model in which electrons were positioned at the eight corners of a cube in a non-polar atom or molecule. This theory was developed in 1902 by Gilbert N. Lewis and published in 1916 in the article The Atom, lewiss theory was based on Abeggs rule. It was further developed in 1919 by Irving Langmuir as the cubical octet atom, the figure below shows structural representations for elements of the second row of the periodic table. The 1916 article by Lewis also introduced the concept of the pair in the covalent bond, the octet rule. Single covalent bonds are formed when two atoms share an edge, as in structure C below and this results in the sharing of two electrons. Ionic bonds are formed by the transfer of an electron from one cube to another, an intermediate state B where only one corner is shared was also postulated by Lewis. Double bonds are formed by sharing a face between two cubic atoms and this results in sharing four electrons, Triple bonds could not be accounted for by the cubical atom model, because there is no way of having two cubes share three parallel edges. Lewis suggested that the pairs in atomic bonds have a special attraction. This allows the formation of a bond by sharing a corner, a double bond by sharing an edge. It also accounts for the rotation around single bonds and for the tetrahedral geometry of methane
33.
Valence electron
–
In a transition metal, a valence electron can also be in an inner shell. An atom with a shell of valence electrons tends to be chemically inert. 2) Because of their tendency either to gain the missing valence electrons, similar to an electron in an inner shell, a valence electron has the ability to absorb or release energy in the form of a photon. An energy gain can trigger an electron to move to an outer shell, or the electron can even break free from its associated atoms valence shell, this is ionization to form a positive ion. When an electron energy, then it can move to an inner shell which is not fully occupied. Valence energy levels correspond to the quantum numbers or are labeled alphabetically with letters used in the X-ray notation. The number of electrons of an element can be determined by the periodic table group in which the element is categorized. With the exception of groups 3–12, the digit of the group number identifies how many valence electrons are associated with a neutral atom of an element listed under that particular column. * Consists of ns and d electrons, alternatively, the d electron count is used. ** Except for helium, which has two valence electrons. The electrons that determine how an atom reacts chemically are those whose average distance from the nucleus is greatest, that is, for a main group element, the valence electrons are defined as those electrons residing in the electronic shell of highest principal quantum number n. Thus, the number of electrons that it may have depends on the electron configuration in a simple way. However, transition elements have partially filled d energy levels, that are close in energy to the ns level. So as opposed to group elements, a valence electron for a transition metal is defined as an electron that resides outside a noble-gas core. Thus, generally, the d electrons in transition metals behave as valence electrons although they are not in the valence shell. For example, manganese has configuration 1s2 2s2 2p6 3s2 3p6 4s2 3d5, this is abbreviated to 4s2 3d5, in this atom, a 3d electron has energy similar to that of a 4s electron, and much higher than that of a 3s or 3p electron. In effect, there are possibly seven valence electrons outside the argon-like core, the farther right in each transition metal series, the lower the energy of an electron in a d subshell and the less such an electron has the properties of a valence electron. Thus, although a nickel atom has, in principle, ten valence electrons, for zinc, the 3d subshell is complete and behaves similarly to core electrons
34.
Irving Langmuir
–
Irving Langmuir /ˈlæŋmjʊr/ was an American chemist and physicist. The Langmuir Laboratory for Atmospheric Research near Socorro, New Mexico, was named in his honor as was the American Chemical Society journal for Surface Science, Irving Langmuir was born in Brooklyn, New York, on January 31,1881. He was the third of the four children of Charles Langmuir and Sadie, during his childhood, Langmuirs parents encouraged him to carefully observe nature and to keep a detailed record of his various observations. When Irving was eleven, it was discovered that he had poor eyesight, when this problem was corrected, details that had previously eluded him were revealed, and his interest in the complications of nature was heightened. During his childhood, Langmuir was influenced by his older brother, Arthur was a research chemist who encouraged Irving to be curious about nature and how things work. Arthur helped Irving set up his first chemistry lab in the corner of his bedroom, Langmuirs hobbies included mountaineering, skiing, piloting his own plane, and classical music. In addition to his professional interest in the politics of atomic energy, Langmuir attended his early education at various schools and institutes in America and Paris. Langmuir graduated high school from Chestnut Hill Academy, a private school located in the affluent Chestnut Hill area in Philadelphia. He graduated with a Bachelor of Science degree in engineering from the Columbia University School of Mines in 1903. He earned his Ph. D. degree in 1906 under Nobel laureate Walther Nernst in Göttingen, for research using the Nernst glower. His doctoral thesis was entitled “On the Partial Recombination of Dissolved Gases During Cooling. ”He later did work in chemistry. Langmuir then taught at Stevens Institute of Technology in Hoboken, New Jersey, until 1909 and his initial contributions to science came from his study of light bulbs. His first major development was the improvement of the diffusion pump and he also discovered that twisting the filament into a tight coil improved its efficiency. These were important developments in the history of the incandescent light bulb and his assistant in vacuum tube research was his cousin William Comings White. As he continued to study filaments in vacuum and different gas environments and he was one of the first scientists to work with plasmas and was the first to call these ionized gases by that name, because they reminded him of blood plasma. Langmuir and Tonks discovered electron density waves in plasmas that are now known as Langmuir waves, the current of a biased probe tip is measured as a function of bias voltage to determine the local plasma temperature and density. He also discovered atomic hydrogen, which he put to use by inventing the atomic hydrogen welding process, plasma welding has since been developed into gas tungsten arc welding. In 1917, he published a paper on the chemistry of oil films that became the basis for the award of the 1932 Nobel Prize in chemistry
35.
Neon
–
Neon is a chemical element with symbol Ne and atomic number 10 and is a noble gas. Neon is a colorless, odorless, inert gas under standard conditions. It was discovered in 1898 as one of the three residual rare inert elements remaining in dry air, after nitrogen, oxygen, argon and carbon dioxide were removed. Neon was the second of three rare gases to be discovered, and was immediately recognized as a new element from its bright red emission spectrum. The name neon is derived from the Greek word, νέον, neuter singular form of νέος, Neon is chemically inert and forms no uncharged chemical compounds. The compounds of neon include ionic molecules, molecules held together by van der Waals forces and clathrates, during cosmic nucleogenesis of the elements, large amounts of neon are built up from the alpha-capture fusion process in stars. Although neon is a common element in the universe and solar system. It composes about 18.2 ppm of air by volume, the reason for neons relative scarcity on Earth and the inner planets is that neon is highly volatile and forms no compounds to fix it to solids. As a result, it escaped from the planetesimals under the warmth of the newly ignited Sun in the early Solar System, even the atmosphere of Jupiter is somewhat depleted of neon, presumably for this reason. It is also lighter than air, causing it to escape even from Earths atmosphere, Neon gives a distinct reddish-orange glow when used in low-voltage neon glow lamps, high-voltage discharge tubes and neon advertising signs. The red emission line from neon also causes the well known red light of helium–neon lasers, Neon is used in some plasma tube and refrigerant applications but has few other commercial uses. It is commercially extracted by the distillation of liquid air. Since air is the source, it is considerably more expensive than helium. Neon, was discovered in 1898 by the British chemists Sir William Ramsay, Neon was discovered when Ramsay chilled a sample of air until it became a liquid, then warmed the liquid and captured the gases as they boiled off. The gases nitrogen, oxygen, and argon had been identified, first to be identified was krypton. The next, after krypton had been removed, was a gas which gave a brilliant red light under spectroscopic discharge and this gas, identified in June, was named neon, the Greek analogue of novum, suggested by Ramsays son. A second gas was reported along with neon, having approximately the same density as argon but with a different spectrum – Ramsay. However, subsequent spectroscopic analysis revealed it to be contaminated with carbon monoxide
36.
Argon
–
Argon is a chemical element with symbol Ar and atomic number 18. It is in group 18 of the table and is a noble gas. Argon is the third-most abundant gas in the Earths atmosphere, at 0. 934% and it is more than twice as abundant as water vapor,23 times as abundant as carbon dioxide, and more than 500 times as abundant as neon. Argon is the most abundant noble gas in Earths crust, comprising 0. 00015% of the crust, nearly all of the argon in Earths atmosphere is radiogenic argon-40, derived from the decay of potassium-40 in the Earths crust. In the universe, argon-36 is by far the most common argon isotope, the name argon is derived from the Greek word ἀργόν, neuter singular form of ἀργός meaning lazy or inactive, as a reference to the fact that the element undergoes almost no chemical reactions. The complete octet in the atomic shell makes argon stable. Its triple point temperature of 83.8058 K is a fixed point in the International Temperature Scale of 1990. Argon is produced industrially by the distillation of liquid air. Argon is also used in incandescent, fluorescent lighting, and other gas-discharge tubes, Argon makes a distinctive blue-green gas laser. Argon is also used in fluorescent glow starters, Argon has approximately the same solubility in water as oxygen and is 2.5 times more soluble in water than nitrogen. Argon is colorless, odorless, nonflammable and nontoxic as a solid, liquid or gas, Argon is chemically inert under most conditions and forms no confirmed stable compounds at room temperature. Although argon is a gas, it can form some compounds under extreme conditions. Argon fluorohydride, a compound of argon with fluorine and hydrogen that is stable below 17 K, has been demonstrated. Although the neutral ground-state chemical compounds of argon are presently limited to HArF, ions, such as ArH+, and excited-state complexes, such as ArF, have been demonstrated. Theoretical calculation predicts several more argon compounds that should be stable but have not yet been synthesized, Argon was suspected to be a component of air by Henry Cavendish in 1785. Argon was first isolated from air in 1894 by Lord Rayleigh and Sir William Ramsay at University College London by removing oxygen, carbon dioxide, water and they had determined that nitrogen produced from chemical compounds was 0. 5% lighter than nitrogen from the atmosphere. The difference was slight, but it was important enough to attract their attention for many months and they concluded that there was another gas in the air mixed in with the nitrogen. Argon was also encountered in 1882 through independent research of H. F. Newall, each observed new lines in the color spectrum of air that did not match known elements
37.
Hypervalent molecule
–
A hypervalent molecule is a molecule that contains one or more main group elements apparently bearing more than eight electrons in their valence shells. Phosphorus pentachloride, sulfur hexafluoride, chlorine trifluoride, and the ion are examples of hypervalent molecules. Hypervalent molecules were first formally defined by Jeremy I, Musher in 1969 as molecules having central atoms of group 15–18 in any valence other than the lowest. Lewis and Irving Langmuir and the debate over the nature of the bond in the 1920s. Lewis maintained the importance of the two-center two-electron bond in describing hypervalence, Langmuir, on the other hand, upheld the dominance of the octet rule and preferred the use of ionic bonds to account for hypervalence without violating the rule. The theoretical basis for hypervalency was not delineated until J. I, in 1990, Magnusson published a seminal work definitively excluding the role of d-orbital hybridization in bonding in hypervalent compounds of second-row elements. This had long been a point of contention and confusion in describing these molecules using molecular orbital theory and these facts were historically interpreted to mean that d-orbitals must be involved in bonding. However, Magnusson concludes in his work that d-orbital involvement is not implicated in hypervalency, both the term and concept of hypervalency still fall under criticism. Consistent with this view is the finding that hypercoordinated molecules based on fluorine ligands. The ionic model holds up well in thermochemical calculations and it predicts favorable exothermic formation of PF4+F− from phosphorus trifluoride PF3 and fluorine F2 whereas a similar reaction forming PH4+H− is not favorable. Durrant has proposed a definition of hypervalency, based on the analysis of atomic charge maps obtained from Atoms in molecules theory. For any particular atom X, if the value of γ is greater than 8, examples of γ calculations for phosphate PO43− and orthonitrate NO43− are shown below. Early considerations of the geometry of hypervalent molecules returned familiar arrangements that were explained by the VSEPR model for atomic bonding. Accordingly, AB5 and AB6 type molecules would possess a trigonal bi-pyramidal and octahedral geometry, however, in order to account for the observed bond angles, bond lengths and apparent violation of the Lewis octet rule, several alternative models have been proposed. Additional modifications to the rule have been attempted to involve ionic characteristics in hypervalent bonding. As one of these modifications, in 1951, the concept of the 3-center 4-electron bond and this model in which the octet rule is preserved was also advocated by Musher. A complete description of hypervalent molecules arises from consideration of molecular orbital theory through quantum mechanical methods and this is a stable configuration only for SX6 molecules containing electronegative ligand atoms like fluorine, which explains why SH6 doesnt form. In the bonding model, the two non-bonding MOs are localized equally on all six fluorine atoms, for a hexacoordinate molecule such as sulfur hexafluoride, each of the six bonds is the same length
38.
Helium
–
Helium is a chemical element with symbol He and atomic number 2. It is a colorless, odorless, tasteless, non-toxic, inert, monatomic gas and its boiling point is the lowest among all the elements. Its abundance is similar to figure in the Sun and in Jupiter. This is due to the high nuclear binding energy of helium-4 with respect to the next three elements after helium. This helium-4 binding energy also accounts for why it is a product of nuclear fusion and radioactive decay. Most helium in the universe is helium-4, and is believed to have formed during the Big Bang. Large amounts of new helium are being created by fusion of hydrogen in stars. Helium is named for the Greek god of the Sun, Helios and it was first detected as an unknown yellow spectral line signature in sunlight during a solar eclipse in 1868 by French astronomer Jules Janssen. Janssen is jointly credited with detecting the element along with Norman Lockyer, Janssen observed during the solar eclipse of 1868 while Lockyer observed from Britain. Lockyer was the first to propose that the line was due to a new element, the formal discovery of the element was made in 1895 by two Swedish chemists, Per Teodor Cleve and Nils Abraham Langlet, who found helium emanating from the uranium ore cleveite. In 1903, large reserves of helium were found in gas fields in parts of the United States. Liquid helium is used in cryogenics, particularly in the cooling of superconducting magnets, a well-known but minor use is as a lifting gas in balloons and airships. As with any gas whose density differs from that of air, inhaling a small volume of helium temporarily changes the timbre, on Earth it is relatively rare—5.2 ppm by volume in the atmosphere. Most terrestrial helium present today is created by the radioactive decay of heavy radioactive elements. Previously, terrestrial helium—a non-renewable resource, because once released into the atmosphere it readily escapes into space—was thought to be in short supply. The first evidence of helium was observed on August 18,1868, the line was detected by French astronomer Jules Janssen during a total solar eclipse in Guntur, India. This line was assumed to be sodium. He concluded that it was caused by an element in the Sun unknown on Earth, Lockyer and English chemist Edward Frankland named the element with the Greek word for the Sun, ἥλιος
39.
Radical (chemistry)
–
In chemistry, a radical is an atom, molecule, or ion that has an unpaired valence electron. Most radicals are reasonably stable only at low concentrations in inert media or in a vacuum. A notable example of a radical is the hydroxyl radical. Two other examples are triplet oxygen and triplet carbene which have two unpaired electrons, free radicals may be created in a number of ways, including synthesis with very dilute or rarefied reagents, reactions at very low temperatures, or breakup of larger molecules. The latter can be affected by any process that puts energy into the parent molecule, such as ionizing radiation, heat, electrical discharges, electrolysis. Radicals are intermediate stages in many chemical reactions, free radicals play an important role in combustion, atmospheric chemistry, polymerization, plasma chemistry, biochemistry, and many other chemical processes. In living organisms, the free radicals superoxide and nitric oxide and their reaction products regulate many processes, such as control of vascular tone and they also play a key role in the intermediary metabolism of various biological compounds. Such radicals can even be messengers in a process dubbed redox signaling, a radical may be trapped within a solvent cage or be otherwise bound. The qualifier free was then needed to specify the unbound case, following recent nomenclature revisions, a part of a larger molecule is now called a functional group or substituent, and radical now implies free. However, the old nomenclature may still appear in some books, the term radical was already in use when the now obsolete radical theory was developed. Louis-Bernard Guyton de Morveau introduced the phrase radical in 1785 and the phrase was employed by Antoine Lavoisier in 1789 in his Traité Élémentaire de Chimie, a radical was then identified as the root base of certain acids. Historically, the radical in radical theory was also used for bound parts of the molecule. These are now called functional groups, for example, methyl alcohol was described as consisting of a methyl radical and a hydroxyl radical. In a modern context the first organic free radical identified was triphenylmethyl radical and this species was discovered by Moses Gomberg in 1900 at the University of Michigan USA. In 1933 Morris Kharash and Frank Mayo proposed that free radicals were responsible for anti-Markovnikov addition of hydrogen bromide to allyl bromide. It should be noted that the electron of the breaking bond also moves to pair up with the attacking radical electron. Free radicals also take part in addition and radical substitution as reactive intermediates. Chain reactions involving free radicals can usually be divided into three distinct processes and these are initiation, propagation, and termination
40.
Chlorofluorocarbon
–
A chlorofluorocarbon is an organic compound that contains only carbon, chlorine, and fluorine, produced as volatile derivative of methane, ethane, and propane. They are also known by the DuPont brand name Freon. The most common representative is dichlorodifluoromethane, many CFCs have been widely used as refrigerants, propellants, and solvents. As in simpler alkanes, carbon in the CFCs bonds with tetrahedral symmetry, because the fluorine and chlorine atoms differ greatly in size and effective charge from hydrogen and from each other, the methane-derived CFCs deviate from perfect tetrahedral symmetry. The physical properties of CFCs and HCFCs are tunable by changes in the number, in general they are volatile, but less so than their parent alkanes. The decreased volatility is attributed to the molecular polarity induced by the halides, thus, methane boils at −161 °C whereas the fluoromethanes boil between −51.7 and −128 °C. The CFCs have still higher boiling points because the chloride is even more polarizable than fluoride, because of their polarity, the CFCs are useful solvents, and their boiling points make them suitable as refrigerants. The densities of CFCs are higher than their corresponding alkanes, in general the density of these compounds correlates with the number of chlorides. CFCs and HCFCs are usually produced by halogen exchange starting from chlorinated methanes and ethanes, the chlorine atom, written often as Cl. behaves very differently from the chlorine molecule. The radical Cl. is long-lived in the atmosphere, where it catalyzes the conversion of ozone into O2. Ozone absorbs UV-B radiation, so its depletion allows more of high energy radiation to reach the Earths surface. Bromine atoms are more efficient catalysts, hence brominated CFCs are also regulated. Applications exploit the low toxicity, low reactivity, and low flammability of the CFCs and HCFCs, every permutation of fluorine, chlorine, and hydrogen based on methane and ethane has been examined and most have been commercialized. Furthermore, many examples are known for higher numbers of carbon as well as related compounds containing bromine, uses include refrigerants, blowing agents, propellants in medicinal applications, and degreasing solvents. Billions of kilograms of chlorodifluoromethane are produced annually as precursor to tetrafluoroethylene, chlorofluorocarbons, when derived from methane and ethane these compounds have the formulae CClmF4−m and C2ClmF6−m, where m is nonzero. Hydrofluorocarbons, when derived from methane, ethane, propane, and butane, these compounds have the respective formulae CFmH4−m, C2FmH6−m, C3FmH8−m, and C4FmH10−m, a numbering system is used for fluorinated alkanes, prefixed with Freon-, R-, CFC-, and HCFC-. The rightmost value indicates the number of atoms, the next value to the left is the number of hydrogen atoms plus 1. Thus, Freon-12 indicates a methane derivative containing two fluorine atoms and no hydrogen, another, easier equation that can be applied to get the correct molecular formula of the CFC/R/Freon class compounds is this to take the numbering and add 90 to it
41.
Carbene
–
In chemistry, a carbene is a molecule containing a neutral carbon atom with a valence of two and two unshared valence electrons. The general formula is R--R or R=C, the term carbene may also refer to the specific compound H2C, also called methylene, the parent hydride from which all other carbene compounds are formally derived. Carbenes are classified as either singlets or triplets, depending upon their electronic structure, Most carbenes are very short lived, although persistent carbenes are known. One well-studied carbene is dichlorocarbene Cl2C, which can be generated in situ from chloroform, the two classes of carbenes are singlet and triplet carbenes. In the language of valence bond theory, the molecule adopts an sp2 hybrid structure, triplet carbenes have two unpaired electrons. Most carbenes have a triplet ground state, except for those with nitrogen, oxygen, or sulfur atoms. Carbenes are called singlet or triplet depending on the electronic spins they possess, triplet carbenes are paramagnetic and may be observed by electron spin resonance spectroscopy if they persist long enough. The total spin of singlet carbenes is zero while that of triplet carbenes is one, bond angles are 125-140° for triplet methylene and 102° for singlet methylene. Triplet carbenes are generally stable in the state, while singlet carbenes occur more often in aqueous media. For simple hydrocarbons, triplet carbenes usually have energies 8 kcal/mol lower than singlet carbenes, thus, in general, triplet is the stable state. Substituents that can donate electron pairs may stabilize the state by delocalizing the pair into an empty p-orbital. If the energy of the state is sufficiently reduced it will actually become the ground state. No viable strategies exist for triplet stabilization, the carbene called 9-fluorenylidene has been shown to be a rapidly equilibrating mixture of singlet and triplet states with an approximately 1.1 kcal/mol energy difference. It is, however, debatable whether diaryl carbenes such as the fluorene carbene are true carbenes because the electrons can delocalize to such an extent that they become in fact biradicals. In silico experiments suggest that triplet carbenes can be stabilized with electropositive heteroatoms such as in silyl and silyloxy carbenes. Singlet and triplet carbenes exhibit divergent reactivity, singlet carbenes generally participate in cheletropic reactions as either electrophiles or nucleophiles. Singlet carbenes with unfilled p-orbital should be electrophilic, triplet carbenes can be considered to be diradicals, and participate in stepwise radical additions. Triplet carbenes have to go through an intermediate with two unpaired electrons whereas singlet carbene can react in a concerted step
42.
Diborane
–
Diborane is the chemical compound consisting of boron and hydrogen with the formula B2H6. It is a colorless and highly unstable gas at room temperature with a sweet odor. Diborane mixes well with air, easily forming explosive mixtures, diborane will ignite spontaneously in moist air at room temperature. Synonyms include boroethane, boron hydride, and diboron hexahydride, diborane is a key boron compound with a variety of applications. The compound is classified as endothermic, meaning that its heat of formation, diborane adopts a D2h structure containing four terminal and two bridging hydrogen atoms. The model determined by molecular orbital theory indicates that the bonds between boron and the hydrogen atoms are conventional 2-center, 2-electron covalent bonds. The bonding between the atoms and the bridging hydrogen atoms is, however, different from that in molecules such as hydrocarbons. Having used two electrons in bonding to the hydrogen atoms, each boron has one valence electron remaining for additional bonding. The bridging hydrogen atoms provide one electron each, thus the B2H2 ring is held together by four electrons, an example of 3-center 2-electron bonding. This type of bond is called a banana bond. The weakness of the B-Hbridge vs B-Hterminal bonds is indicated by their signatures in the infrared spectrum, being ~2100 and 2500 cm−1. The structure is isoelectronic with C2H62+, which would arise from the diprotonation of the planar molecule ethene, diborane is one of many compounds with such unusual bonding. Of the other elements in Group IIIA, gallium is known to form a compound, digallane. Aluminium forms a hydride, n, although unstable Al2H6 has been isolated in solid hydrogen and is isostructural with diborane. Extensive studies of diborane have led to the development of multiple syntheses, most preparations entail reactions of hydride donors with boron halides or alkoxides. 2 BH4− +2 H+ →2 H2 + B2H6 Similarly, oxidation of borohydride salts has been demonstrated and remains convenient for small scale preparations. For example, using iodine as an oxidizer,2 NaBH4 + I2 →2 NaI + B 2H6 + H2 Another small-scale synthesis uses potassium hydroborate, diborane is a highly reactive and versatile reagent that has a large number of applications. Its dominating reaction pattern involves formation of adducts with Lewis bases, often such initial adducts proceed rapidly to give other products
43.
Hydrogen
–
Hydrogen is a chemical element with chemical symbol H and atomic number 1. With a standard weight of circa 1.008, hydrogen is the lightest element on the periodic table. Its monatomic form is the most abundant chemical substance in the Universe, non-remnant stars are mainly composed of hydrogen in the plasma state. The most common isotope of hydrogen, termed protium, has one proton, the universal emergence of atomic hydrogen first occurred during the recombination epoch. At standard temperature and pressure, hydrogen is a colorless, odorless, tasteless, non-toxic, nonmetallic, since hydrogen readily 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 an 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 charge when it is known as a hydride. The hydrogen cation is written as though composed of a bare proton, Hydrogen gas was first artificially produced in the early 16th century by the reaction of acids on metals. Industrial production is mainly from steam reforming natural gas, and less often 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, mostly 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 very 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%, the explosive reactions may be triggered by spark, heat, or sunlight. The hydrogen autoignition temperature, the temperature of spontaneous ignition in air, is 500 °C, the detection of a burning hydrogen leak may require a flame detector, such leaks can be very dangerous. Hydrogen flames in other conditions are blue, resembling blue natural gas flames, 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 mixture of hydrogen to oxygen combined with carbon compounds from the airship skin. H2 reacts with every oxidizing element, the ground state energy level of the electron in a hydrogen atom is −13.6 eV, which is equivalent to an ultraviolet photon of roughly 91 nm wavelength. The energy levels of hydrogen can be calculated fairly accurately using the Bohr model of the atom, however, the atomic electron and proton are held together by electromagnetic force, while planets and celestial objects are held by gravity. The most complicated treatments allow for the effects of special relativity
44.
Lithium
–
Lithium is a chemical element with the symbol Li and atomic number 3. It is a soft, silver-white metal belonging to the metal group of chemical elements. Under standard conditions, it is the lightest metal and the least dense solid element, like all alkali metals, lithium is highly reactive and flammable. For this reason, it is stored in mineral oil. When cut open, it exhibits a metallic luster, but contact with moist air corrodes the surface quickly to a silvery gray. Because of its reactivity, lithium never occurs freely in nature, and instead, appears only in compounds. Lithium occurs in a number of minerals, but due to its solubility as an ion, is present in ocean water and is commonly obtained from brines. On a commercial scale, lithium is isolated electrolytically from a mixture of lithium chloride, the nucleus of the lithium atom verges on instability, since the two stable lithium isotopes found in nature have among the lowest binding energies per nucleon of all stable nuclides. Because of its relative instability, lithium is less common in the solar system than 25 of the first 32 chemical elements even though the nuclei are very light in atomic weight. For related reasons, lithium has important links to nuclear physics, the transmutation of lithium atoms to helium in 1932 was the first fully man-made nuclear reaction, and lithium-6 deuteride serves as a fusion fuel in staged thermonuclear weapons. These uses consume more than three quarters of lithium production, Lithium is found in variable amounts in foods, primary food sources are grains and vegetables, in some areas, the drinking water also provides significant amounts of the element. Human dietary lithium intakes depend on location and the type of foods consumed, traces of lithium were detected in human organs and fetal tissues already in the late 19th century, leading to early suggestions as to possible specific functions in the organism. However, it took another century until evidence for the essentiality of lithium became available, in studies conducted from the 1970s to the 1990s, rats and goats maintained on low-lithium rations were shown to exhibit higher mortalities as well as reproductive and behavioral abnormalities. Lithium appears to play an important role during the early fetal development as evidenced by the high lithium contents of the embryo during the early gestational period. The available experimental evidence now appears to be sufficient to accept lithium as essential, the lithium ion Li+ administered as any of several lithium salts has proven to be useful as a mood-stabilizing drug in the treatment of bipolar disorder in humans. Like the other metals, lithium has a single valence electron that is easily given up to form a cation. Because of this, lithium is a conductor of heat and electricity as well as a highly reactive element. Lithiums low reactivity is due to the proximity of its electron to its nucleus
