1.
Atom
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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
2.
Ion
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An ion is an atom or a molecule in which the total number of electrons is not equal to the total number of protons, giving the atom or molecule a net positive or negative electrical charge. Ions can be created, by chemical or physical means. In chemical terms, if an atom loses one or more electrons. If an atom gains electrons, it has a net charge and is known as an anion. Ions consisting of only a single atom are atomic or monatomic ions, because of their electric charges, cations and anions attract each other and readily form ionic compounds, such as salts. In the case of ionization of a medium, such as a gas, which are known as ion pairs are created by ion impact, and each pair consists of a free electron. The word ion comes from the Greek word ἰόν, ion, going and this term was introduced by English physicist and chemist Michael Faraday in 1834 for the then-unknown species that goes from one electrode to the other through an aqueous medium. Faraday also introduced the words anion for a charged ion. In Faradays nomenclature, cations were named because they were attracted to the cathode in a galvanic device, arrhenius explanation was that in forming a solution, the salt dissociates into Faradays ions. Arrhenius proposed that ions formed even in the absence of an electric current, ions in their gas-like state are highly reactive, and do not occur in large amounts on Earth, except in flames, lightning, electrical sparks, and other plasmas. These gas-like ions rapidly interact with ions of charge to give neutral molecules or ionic salts. These stabilized species are commonly found in the environment at low temperatures. A common example is the present in seawater, which are derived from the dissolved salts. Electrons, due to their mass and thus larger space-filling properties as matter waves, determine the size of atoms. Thus, anions are larger than the parent molecule or atom, as the excess electron repel each other, as such, in general, cations are smaller than the corresponding parent atom or molecule due to the smaller size of its electron cloud. One particular cation contains no electrons, and thus consists of a single proton - very much smaller than the parent hydrogen atom. Since the electric charge on a proton is equal in magnitude to the charge on an electron, an anion, from the Greek word ἄνω, meaning up, is an ion with more electrons than protons, giving it a net negative charge. A cation, from the Greek word κατά, meaning down, is an ion with fewer electrons than protons, there are additional names used for ions with multiple charges
3.
Molecule
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A molecule is an electrically neutral group of two or more atoms held together by chemical bonds. Molecules are distinguished from ions by their lack of electrical charge, however, in quantum physics, organic chemistry, and biochemistry, the term molecule is often used less strictly, also being applied to polyatomic ions. In the kinetic theory of gases, the molecule is often used for any gaseous particle regardless of its composition. According to this definition, noble gas atoms are considered molecules as they are in fact monoatomic molecules. A molecule may be homonuclear, that is, it consists of atoms of one element, as with oxygen, or it may be heteronuclear. Atoms and complexes connected by non-covalent interactions, such as hydrogen bonds or ionic bonds, are not considered single molecules. Molecules as components of matter are common in organic substances and they also make up most of the oceans and atmosphere. Also, no typical molecule can be defined for ionic crystals and covalent crystals, the theme of repeated unit-cellular-structure also holds for most condensed phases with metallic bonding, which means that solid metals are also not made of molecules. In glasses, atoms may also be together by chemical bonds with no presence of any definable molecule. The science of molecules is called molecular chemistry or molecular physics, in practice, however, this distinction is vague. In molecular sciences, a molecule consists of a system composed of two or more atoms. Polyatomic ions may sometimes be thought of as electrically charged molecules. The term unstable molecule is used for very reactive species, i. e, according to Merriam-Webster and the Online Etymology Dictionary, the word molecule derives from the Latin moles or small unit of mass. Molecule – extremely minute particle, from French molécule, from New Latin molecula, diminutive of Latin moles mass, a vague meaning at first, the vogue for the word can be traced to the philosophy of Descartes. The definition of the molecule has evolved as knowledge of the structure of molecules has increased, earlier definitions were less precise, defining molecules as the smallest particles of pure chemical substances that still retain their composition and chemical properties. Molecules are held together by covalent bonding or ionic bonding. Several types of non-metal elements exist only as molecules in the environment, for example, hydrogen only exists as hydrogen molecule. A molecule of a compound is made out of two or more elements, a covalent bond is a chemical bond that involves the sharing of electron pairs between atoms
4.
Chemical compound
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A chemical compound is an entity consisting of two or more atoms, at least two from different elements, which associate via chemical bonds. Many chemical compounds have a numerical identifier assigned by the Chemical Abstracts Service. For example, water is composed of two atoms bonded to one oxygen atom, the chemical formula is H2O. A compound can be converted to a different chemical composition by interaction with a chemical compound via a chemical reaction. In this process, bonds between atoms are broken in both of the compounds, and then bonds are reformed so that new associations are made between atoms. Schematically, this reaction could be described as AB + CD → AC + BD, where A, B, C, and D are each unique atoms, and AB, CD, AC, and BD are each unique compounds. A chemical element bonded to a chemical element is not a chemical compound since only one element. Examples are the diatomic hydrogen and the polyatomic molecule sulfur. Chemical compounds have a unique and defined chemical structure held together in a spatial arrangement by chemical bonds. Pure chemical elements are not considered chemical compounds, failing the two or more atom requirement, though they often consist of molecules composed of multiple atoms. There is varying and sometimes inconsistent nomenclature differentiating substances, which include truly non-stoichiometric examples, from chemical compounds, other compounds regarded as chemically identical may have varying amounts of heavy or light isotopes of the constituent elements, which changes the ratio of elements by mass slightly. Characteristic properties of compounds include that elements in a compound are present in a definite proportion, for example, the molecule of the compound water is composed of hydrogen and oxygen in a ratio of 2,1. In addition, compounds have a set of properties. The physical and chemical properties of compounds differ from those of their constituent elements, however, mixtures can be created by mechanical means alone, but a compound can be created only by a chemical reaction. Some mixtures are so combined that they have some properties similar to compounds. Other examples of compound-like mixtures include intermetallic compounds and solutions of metals in a liquid form of ammonia. Compounds may be described using formulas in various formats, for compounds that exist as molecules, the formula for the molecular unit is shown. For polymeric materials, such as minerals and many metal oxides, the elements in a chemical formula are normally listed in a specific order, called the Hill system
5.
Coulomb's law
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Coulombs law, or Coulombs inverse-square law, is a law of physics that describes force interacting between static electrically charged particles. The force of interaction between the charges is attractive if the charges have opposite signs and repulsive if like-signed, the law was first published in 1784 by French physicist Charles Augustin de Coulomb and was essential to the development of the theory of electromagnetism. It is analogous to Isaac Newtons inverse-square law of universal gravitation, Coulombs law can be used to derive Gausss law, and vice versa. The law has been tested extensively, and all observations have upheld the laws principle, ancient cultures around the Mediterranean knew that certain objects, such as rods of amber, could be rubbed with cats fur to attract light objects like feathers. Thales was incorrect in believing the attraction was due to a magnetic effect and he coined the New Latin word electricus to refer to the property of attracting small objects after being rubbed. This association gave rise to the English words electric and electricity, however, he did not generalize or elaborate on this. In 1767, he conjectured that the force between charges varied as the square of the distance. In 1769, Scottish physicist John Robison announced that, according to his measurements, in the early 1770s, the dependence of the force between charged bodies upon both distance and charge had already been discovered, but not published, by Henry Cavendish of England. Finally, in 1785, the French physicist Charles-Augustin de Coulomb published his first three reports of electricity and magnetism where he stated his law and this publication was essential to the development of the theory of electromagnetism. The torsion balance consists of a bar suspended from its middle by a thin fiber, the fiber acts as a very weak torsion spring. In Coulombs experiment, the balance was an insulating rod with a metal-coated ball attached to one end. The ball was charged with a charge of static electricity. The two charged balls repelled one another, twisting the fiber through an angle, which could be read from a scale on the instrument. By knowing how much force it took to twist the fiber through a given angle, the force is along the straight line joining them. If the two charges have the sign, the electrostatic force between them is repulsive, if they have different signs, the force between them is attractive. Coulombs law can also be stated as a mathematical expression. The vector form of the equation calculates the force F1 applied on q1 by q2, if r12 is used instead, then the effect on q2 can be found. It can be calculated using Newtons third law, F2 = −F1
6.
Covalent bond
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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
7.
Ionic bonding
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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
8.
Intermolecular force
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Intermolecular forces are weak relative to intramolecular forces – the forces which hold a molecule together. For example, the covalent bond, involving sharing electron pairs between atoms, is stronger than the forces present between neighboring molecules. Both sets of forces are essential parts of force fields used in molecular mechanics. The investigation of intermolecular forces starts from macroscopic observations which indicate the existence and these observations include non-ideal-gas thermodynamic behavior reflected by virial coefficients, vapor pressure, viscosity, superficial tension, and absorption data. The first reference to the nature of microscopic forces is found in Alexis Clairauts work Theorie de la Figure de la Terre, other scientists who have contributed to the investigation of microscopic forces include, Laplace, Gauss, Maxwell and Boltzmann. The link to microscopic aspects is given by virial coefficients and Lennard-Jones potentials, dipole-dipole interactions are electrostatic interactions between permanent dipoles in molecules. These interactions tend to align the molecules to increase attraction, an example of a dipole-dipole interaction can be seen in hydrogen chloride, the positive end of a polar molecule will attract the negative end of the other molecule and influence its position. Polar molecules have a net attraction between them, examples of polar molecules include hydrogen chloride and chloroform. H δ + − Cl δ − ⋯ H δ + − Cl δ − Often molecules contain dipolar groups and this occurs if there is symmetry within the molecule that causes the dipoles to cancel each other out. This occurs in such as tetrachloromethane and carbon dioxide. The dipole-dipole interaction between two atoms is usually zero, since atoms rarely carry a permanent dipole. Ion-dipole and ion-induced dipole forces are similar to dipole-dipole and induced-dipole interactions but involve ions, ion-dipole and ion-induced dipole forces are stronger than dipole-dipole interactions because the charge of any ion is much greater than the charge of a dipole moment. Ion-dipole bonding is stronger than hydrogen bonding, an ion-dipole force consists of an ion and a polar molecule interacting. They align so that the positive and negative groups are next to one another, an ion-induced dipole force consists of an ion and a non-polar molecule interacting. Like a dipole-induced dipole force, the charge of the ion causes distortion of the cloud on the non-polar molecule. A hydrogen bond is the attraction between the pair of an electronegative atom and a hydrogen atom that is bonded to either nitrogen, oxygen. The hydrogen bond is described as a strong electrostatic dipole-dipole interaction. Intermolecular hydrogen bonding is responsible for the boiling point of water compared to the other group 16 hydrides
9.
London dispersion force
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London dispersion forces are a type of force acting between atoms and molecules. They are part of the van der Waals forces, the LDF is named after the German-American physicist Fritz London. The LDF is a weak intermolecular force arising from quantum-induced instantaneous polarization multipoles in molecules and they can therefore act between molecules without permanent multipole moments. London forces are exhibited by non-polar molecules because of the movements of the electrons in interacting molecules. Because the electrons in adjacent molecules flee as they repel each other and this is frequently described as the formation of instantaneous dipoles that attract each other. London forces become stronger as the atom in question becomes larger and this is due to the increased polarizability of molecules with larger, more dispersed electron clouds. This trend is exemplified by the halogens, fluorine and chlorine are gases at room temperature, bromine is a liquid, and iodine is a solid. The London forces also become stronger with larger amounts of surface contact, greater surface area means closer interaction between different molecules. The first explanation of the attraction between noble gas atoms was given by Fritz London in 1930 and he used a quantum-mechanical theory based on second-order perturbation theory. The perturbation is because of the Coulomb interaction between the electrons and nuclei of the two moieties, the second-order perturbation expression of the interaction energy contains a sum over states. The states appearing in this sum are simple products of the electronic states of the monomers. Thus, no intermolecular antisymmetrization of the states is included. London wrote Taylor series expansion of the perturbation in 1 R and this expansion is known as the multipole expansion because the terms in this series can be regarded as energies of two interacting multipoles, one on each monomer. Substitution of the form of V into the second-order energy yields an expression that resembles somewhat an expression describing the interaction between instantaneous multipoles. Additionally, an approximation, named after Albrecht Unsöld, must be introduced in order to obtain a description of London dispersion in terms of dipole polarizabilities and ionization potentials. In this manner, the approximation is obtained for the dispersion interaction E A B d i s p between two atoms A and B. Here α A and α B are the dipole polarizabilities of the respective atoms, the quantities I A and I B are the first ionization potentials of the atoms, and R is the intermolecular distance. E A B d i s p ≈ −32 I A I B I A + I B α A α B R6 Note that this final London equation does not contain instantaneous dipoles
10.
Hydrogen bond
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Hydrogen bonds can occur between molecules or within different parts of a single molecule. Depending on geometry and environment, the hydrogen bond free energy content is between 1 and 5 kcal/mol and this makes it stronger than a van der Waals interaction, but weaker than covalent or ionic bonds. This type of bond can occur in molecules such as water and in organic molecules like DNA. Intermolecular hydrogen bonding is responsible for the boiling point of water compared to the other group 16 hydrides that have much weaker hydrogen bonds. Intramolecular hydrogen bonding is responsible for the secondary and tertiary structures of proteins. It also plays an important role in the structure of polymers, in 2011, an IUPAC Task Group recommended a modern evidence-based definition of hydrogen bonding, which was published in the IUPAC journal Pure and Applied Chemistry. An accompanying detailed technical report provides the rationale behind the new definition, a hydrogen atom attached to a relatively electronegative atom will play the role of the hydrogen bond donor. This electronegative atom is usually fluorine, oxygen, or nitrogen, a hydrogen attached to carbon can also participate in hydrogen bonding when the carbon atom is bound to electronegative atoms, as is the case in chloroform, CHCl3. An example of a hydrogen donor is the hydrogen from the hydroxyl group of ethanol. In a hydrogen bond, the electronegative atom not covalently attached to the hydrogen is named proton acceptor, because of the small size of hydrogen relative to other atoms and molecules, the resulting charge, though only partial, represents a large charge density. A hydrogen bond results when this positive charge density attracts a lone pair of electrons on another heteroatom. The hydrogen bond is described as an electrostatic dipole-dipole interaction. These covalent features are more substantial when acceptors bind hydrogens from more electronegative donors, the partially covalent nature of a hydrogen bond raises the following questions, To which molecule or atom does the hydrogen nucleus belong. And Which should be labeled donor and which acceptor, liquids that display hydrogen bonding are called associated liquids. Hydrogen bonds can vary in strength from weak to extremely strong. For example, the central interresidue N−H···N hydrogen bond between guanine and cytosine is much stronger in comparison to the N−H···N bond between the adenine-thymine pair, the length of hydrogen bonds depends on bond strength, temperature, and pressure. The bond strength itself is dependent on temperature, pressure, bond angle, the typical length of a hydrogen bond in water is 197 pm. The ideal bond angle depends on the nature of the hydrogen bond donor, moore and Winmill used the hydrogen bond to account for the fact that trimethylammonium hydroxide is a weaker base than tetramethylammonium hydroxide
11.
Electromagnetic force
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Electromagnetism is a branch of physics involving the study of the electromagnetic force, a type of physical interaction that occurs between electrically charged particles. The electromagnetic force usually exhibits electromagnetic fields such as fields, magnetic fields. The other three fundamental interactions are the interaction, the weak interaction, and gravitation. The word electromagnetism is a form of two Greek terms, ἤλεκτρον, ēlektron, amber, and μαγνῆτις λίθος magnētis lithos, which means magnesian stone. The electromagnetic force plays a role in determining the internal properties of most objects encountered in daily life. Ordinary matter takes its form as a result of forces between individual atoms and molecules in matter, and is a manifestation of the electromagnetic force. Electrons are bound by the force to atomic nuclei, and their orbital shapes. The electromagnetic force governs the processes involved in chemistry, which arise from interactions between the electrons of neighboring atoms, there are numerous mathematical descriptions of the electromagnetic field. In classical electrodynamics, electric fields are described as electric potential, although electromagnetism is considered one of the four fundamental forces, at high energy the weak force and electromagnetic force are unified as a single electroweak force. In the history of the universe, during the epoch the unified force broke into the two separate forces as the universe cooled. Originally, electricity and magnetism were considered to be two separate forces, Magnetic poles attract or repel one another in a manner similar to positive and negative charges and always exist as pairs, every north pole is yoked to a south pole. An electric current inside a wire creates a corresponding magnetic field outside the wire. Its direction depends on the direction of the current in the wire. A current is induced in a loop of wire when it is moved toward or away from a field, or a magnet is moved towards or away from it. While preparing for a lecture on 21 April 1820, Hans Christian Ørsted made a surprising observation. As he was setting up his materials, he noticed a compass needle deflected away from north when the electric current from the battery he was using was switched on. At the time of discovery, Ørsted did not suggest any explanation of the phenomenon. However, three later he began more intensive investigations
12.
Electron
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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
13.
Proton
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A proton is a subatomic particle, symbol p or p+, with a positive electric charge of +1e elementary charge and mass slightly less than that of a neutron. Protons and neutrons, each with masses of one atomic mass unit, are collectively referred to as nucleons. One or more protons are present in the nucleus of every atom, the number of protons in the nucleus is the defining property of an element, and is referred to as the atomic number. Since each element has a number of protons, each element has its own unique atomic number. The word proton is Greek for first, and this name was given to the nucleus by Ernest Rutherford in 1920. In previous years, Rutherford had discovered that the nucleus could be extracted from the nuclei of nitrogen by atomic collisions. Protons were therefore a candidate to be a particle, and hence a building block of nitrogen. In the modern Standard Model of particle physics, protons are hadrons, and like neutrons, although protons were originally considered fundamental or elementary particles, they are now known to be composed of three valence quarks, two up quarks and one down quark. The rest masses of quarks contribute only about 1% of a protons mass, the remainder of a protons mass is due to quantum chromodynamics binding energy, which includes the kinetic energy of the quarks and the energy of the gluon fields that bind the quarks together. At sufficiently low temperatures, free protons will bind to electrons, however, the character of such bound protons does not change, and they remain protons. A fast proton moving through matter will slow by interactions with electrons and nuclei, the result is a protonated atom, which is a chemical compound of hydrogen. In vacuum, when electrons are present, a sufficiently slow proton may pick up a single free electron, becoming a neutral hydrogen atom. Such free hydrogen atoms tend to react chemically with other types of atoms at sufficiently low energies. When free hydrogen atoms react with other, they form neutral hydrogen molecules. Protons are spin-½ fermions and are composed of three quarks, making them baryons. Protons have an exponentially decaying positive charge distribution with a mean square radius of about 0.8 fm. Protons and neutrons are both nucleons, which may be together by the nuclear force to form atomic nuclei. The nucleus of the most common isotope of the atom is a lone proton
14.
Atomic nucleus
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After the discovery of the neutron in 1932, models for a nucleus composed of protons and neutrons were quickly developed by Dmitri Ivanenko and Werner Heisenberg. Almost all of the mass of an atom is located in the nucleus, protons and neutrons are bound together to form a nucleus by the nuclear force. The diameter of the nucleus is in the range of 6985175000000000000♠1.75 fm for hydrogen to about 6986150000000000000♠15 fm for the heaviest atoms and these dimensions are much smaller than the diameter of the atom itself, by a factor of about 23,000 to about 145,000. The branch of physics concerned with the study and understanding of the nucleus, including its composition. The nucleus was discovered in 1911, as a result of Ernest Rutherfords efforts to test Thomsons plum pudding model of the atom, the electron had already been discovered earlier by J. J. Knowing that atoms are electrically neutral, Thomson postulated that there must be a charge as well. In his plum pudding model, Thomson suggested that an atom consisted of negative electrons randomly scattered within a sphere of positive charge, to his surprise, many of the particles were deflected at very large angles. This justified the idea of an atom with a dense center of positive charge. The term nucleus is from the Latin word nucleus, a diminutive of nux, in 1844, Michael Faraday used the term to refer to the central point of an atom. The modern atomic meaning was proposed by Ernest Rutherford in 1912, the adoption of the term nucleus to atomic theory, however, was not immediate. In 1916, for example, Gilbert N, the nuclear strong force extends far enough from each baryon so as to bind the neutrons and protons together against the repulsive electrical force between the positively charged protons. The nuclear strong force has a short range, and essentially drops to zero just beyond the edge of the nucleus. The collective action of the charged nucleus is to hold the electrically negative charged electrons in their orbits about the nucleus. The collection of negatively charged electrons orbiting the nucleus display an affinity for certain configurations, which chemical element an atom represents is determined by the number of protons in the nucleus, the neutral atom will have an equal number of electrons orbiting that nucleus. Individual chemical elements can create more stable electron configurations by combining to share their electrons and it is that sharing of electrons to create stable electronic orbits about the nucleus that appears to us as the chemistry of our macro world. Protons define the entire charge of a nucleus, and hence its chemical identity, neutrons are electrically neutral, but contribute to the mass of a nucleus to nearly the same extent as the protons. Neutrons explain the phenomenon of isotopes – varieties of the chemical element which differ only in their atomic mass. They are sometimes viewed as two different quantum states of the particle, the nucleon
15.
Matter wave
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Matter waves are a central part of the theory of quantum mechanics, being an example of wave–particle duality. All matter can exhibit wave-like behavior, for example, a beam of electrons can be diffracted just like a beam of light or a water wave. The concept that matter behaves like a wave is referred to as the de Broglie hypothesis due to having been proposed by Louis de Broglie in 1924. Matter waves are referred to as de Broglie waves, the de Broglie wavelength is the wavelength, λ, associated with a massive particle and is related to its momentum, p, through the Planck constant, h, λ = h p. The wave-like behavior of matter is crucial to the theory of atomic structure. In 1900, this division was exposed to doubt, when, investigating the theory of black body thermal radiation and it was thoroughly challenged in 1905. Extending Plancks investigation in several ways, including its connection with the photoelectric effect, light quanta are now called photons. In the modern convention, frequency is symbolized by f as is done in the rest of this article, einstein’s postulate was confirmed experimentally by Robert Millikan and Arthur Compton over the next two decades. De Broglie, in his 1924 PhD thesis, proposed that just as light has both wave-like and particle-like properties, electrons also have wave-like properties, the relationship is now known to hold for all types of matter, all matter exhibits properties of both particles and waves. In 1926, Erwin Schrödinger published an equation describing how a matter wave should evolve—the matter wave analogue of Maxwell’s equations—and used it to derive the energy spectrum of hydrogen, furthermore, neutral atoms and even molecules have been shown to be wave-like. In 1927 at Bell Labs, Clinton Davisson and Lester Germer fired slow-moving electrons at a nickel target. The angular dependence of the electron intensity was measured, and was determined to have the same diffraction pattern as those predicted by Bragg for x-rays. At the same time George Paget Thomson at the University of Aberdeen was independently firing electrons at very thin metal foils to demonstrate the same effect, before the acceptance of the de Broglie hypothesis, diffraction was a property that was thought to be exhibited only by waves. Therefore, the presence of any diffraction effects by matter demonstrated the wave-like nature of matter, when the de Broglie wavelength was inserted into the Bragg condition, the observed diffraction pattern was predicted, thereby experimentally confirming the de Broglie hypothesis for electrons. This was a result in the development of quantum mechanics. Just as the photoelectric effect demonstrated the particle nature of light, the Davisson–Germer experiment showed the wave-nature of matter, advances in laser cooling have allowed cooling of neutral atoms down to nanokelvin temperatures. At these temperatures, the thermal de Broglie wavelengths come into the micrometre range and this effect has been used to demonstrate atomic holography, and it may allow the construction of an atom probe imaging system with nanometer resolution. The description of phenomena is based on the wave properties of neutral atoms
16.
Crystal
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A crystal or crystalline solid is a solid material whose constituents are arranged in a highly ordered microscopic structure, forming a crystal lattice that extends in all directions. In addition, macroscopic single crystals are usually identifiable by their geometrical shape, the scientific study of crystals and crystal formation is known as crystallography. The process of crystal formation via mechanisms of crystal growth is called crystallization or solidification, the word crystal derives from the Ancient Greek word κρύσταλλος, meaning both ice and rock crystal, from κρύος, icy cold, frost. Examples of large crystals include snowflakes, diamonds, and table salt, most inorganic solids are not crystals but polycrystals, i. e. many microscopic crystals fused together into a single solid. Examples of polycrystals include most metals, rocks, ceramics, a third category of solids is amorphous solids, where the atoms have no periodic structure whatsoever. Examples of amorphous solids include glass, wax, and many plastics, Crystals are often used in pseudoscientific practices such as crystal therapy, and, along with gemstones, are sometimes associated with spellwork in Wiccan beliefs and related religious movements. The scientific definition of a crystal is based on the arrangement of atoms inside it. A crystal is a solid where the form a periodic arrangement. For example, when liquid water starts freezing, the change begins with small ice crystals that grow until they fuse. Most macroscopic inorganic solids are polycrystalline, including almost all metals, ceramics, ice, rocks, solids that are neither crystalline nor polycrystalline, such as glass, are called amorphous solids, also called glassy, vitreous, or noncrystalline. These have no periodic order, even microscopically, there are distinct differences between crystalline solids and amorphous solids, most notably, the process of forming a glass does not release the latent heat of fusion, but forming a crystal does. A crystal structure is characterized by its cell, a small imaginary box containing one or more atoms in a specific spatial arrangement. The unit cells are stacked in three-dimensional space to form the crystal, the symmetry of a crystal is constrained by the requirement that the unit cells stack perfectly with no gaps. There are 219 possible crystal symmetries, called space groups. These are grouped into 7 crystal systems, such as cubic crystal system or hexagonal crystal system, Crystals are commonly recognized by their shape, consisting of flat faces with sharp angles. Euhedral crystals are those with obvious, well-formed flat faces, anhedral crystals do not, usually because the crystal is one grain in a polycrystalline solid. The flat faces of a crystal are oriented in a specific way relative to the underlying atomic arrangement of the crystal. This occurs because some surface orientations are more stable than others, as a crystal grows, new atoms attach easily to the rougher and less stable parts of the surface, but less easily to the flat, stable surfaces
17.
Metal
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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
18.
Lewis structure
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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
19.
Carbon
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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
20.
Hydrogen
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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
21.
Oxygen
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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
22.
Quantum mechanics
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Quantum mechanics, including quantum field theory, is a branch of physics which is the fundamental theory of nature at small scales and low energies of atoms and subatomic particles. Classical physics, the physics existing before quantum mechanics, derives from quantum mechanics as an approximation valid only at large scales, early quantum theory was profoundly reconceived in the mid-1920s. The reconceived theory is formulated in various specially developed mathematical formalisms, in one of them, a mathematical function, the wave function, provides information about the probability amplitude of position, momentum, and other physical properties of a particle. In 1803, Thomas Young, an English polymath, performed the famous experiment that he later described in a paper titled On the nature of light. This experiment played a role in the general acceptance of the wave theory of light. In 1838, Michael Faraday discovered cathode rays, Plancks hypothesis that energy is radiated and absorbed in discrete quanta precisely matched the observed patterns of black-body radiation. In 1896, Wilhelm Wien empirically determined a distribution law of black-body radiation, ludwig Boltzmann independently arrived at this result by considerations of Maxwells equations. However, it was only at high frequencies and underestimated the radiance at low frequencies. Later, Planck corrected this model using Boltzmanns statistical interpretation of thermodynamics and proposed what is now called Plancks law, following Max Plancks solution in 1900 to the black-body radiation problem, Albert Einstein offered a quantum-based theory to explain the photoelectric effect. Among the first to study quantum phenomena in nature were Arthur Compton, C. V. Raman, robert Andrews Millikan studied the photoelectric effect experimentally, and Albert Einstein developed a theory for it. In 1913, Peter Debye extended Niels Bohrs theory of structure, introducing elliptical orbits. This phase is known as old quantum theory, according to Planck, each energy element is proportional to its frequency, E = h ν, where h is Plancks constant. Planck cautiously insisted that this was simply an aspect of the processes of absorption and emission of radiation and had nothing to do with the reality of the radiation itself. In fact, he considered his quantum hypothesis a mathematical trick to get the right rather than a sizable discovery. He won the 1921 Nobel Prize in Physics for this work, lower energy/frequency means increased time and vice versa, photons of differing frequencies all deliver the same amount of action, but do so in varying time intervals. High frequency waves are damaging to human tissue because they deliver their action packets concentrated in time, the Copenhagen interpretation of Niels Bohr became widely accepted. In the mid-1920s, developments in mechanics led to its becoming the standard formulation for atomic physics. In the summer of 1925, Bohr and Heisenberg published results that closed the old quantum theory, out of deference to their particle-like behavior in certain processes and measurements, light quanta came to be called photons
23.
Octet rule
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The rule is especially applicable to carbon, nitrogen, oxygen, and the halogens, but also to metals such as sodium or magnesium. The valence electrons can be counted using a Lewis electron dot diagram as shown at the right for carbon dioxide, the electrons shared by the two atoms in a covalent bond are counted twice, once for each atom. In carbon dioxide each oxygen shares four electrons with the central carbon, all four of these electrons are counted in both the carbon octet and the oxygen octet. Ionic bonding is common between pairs of atoms, where one of the pair is a metal of low electronegativity, a chlorine atom has seven electrons in its outer electron shell, the first and second shells being filled with two and eight electrons respectively. The first electron affinity of chlorine is +328.8 kJ per mole of chlorine atoms, adding a second electron to chlorine requires energy, energy that cannot be recovered by formation of a chemical bond. The result is that chlorine will very often form a compound in which it has eight electrons in its outer shell, a sodium atom has a single electron in its outermost electron shell, the first and second shells again being full with two and eight electrons respectively. To remove this outer electron requires only the first ionization energy, which is +495.8 kJ per mole of sodium atoms, a small amount of energy. By contrast, the electron resides in the deeper second electron shell. Thus sodium will, in most cases, form a compound in which it has lost an electron and have a full outer shell of eight electrons. The energy required to transfer an electron from an atom to a chlorine atom is small. This energy is offset by the lattice energy of sodium chloride. This completes the explanation of the rule in this case. In 1893, Alfred Werner showed that the number of atoms or groups associated with an atom is often 4 or 6, other coordination numbers up to a maximum of 8 were known. Abegg noted that the difference between the positive and negative valences of an element under his model is frequently eight. In 1919 Irving Langmuir refined these concepts further and renamed them the cubical octet atom, the octet theory evolved into what is now known as the octet rule. The quantum theory of the atom explains the eight electrons as a shell with an s2p6 electron configuration. A closed-shell configuration is one in which low-lying energy levels are full, for example, the neon atom ground state has a full n =2 shell and an empty n =3 shell. According to the rule, the atoms immediately before and after neon in the periodic table, tend to attain a similar configuration by gaining, losing
24.
VSEPR theory
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Valence shell electron pair repulsion theory is a model used in chemistry to predict the geometry of individual molecules from the number of electron pairs surrounding their central atoms. It is also named the Gillespie-Nyholm theory after its two main developers, the acronym VSEPR is pronounced either ves-pur or vuh-seh-per. Gillespie has emphasized that the repulsion due to the Pauli exclusion principle is more important in determining molecular geometry than the electrostatic repulsion. VSEPR theory is based on electron density rather than mathematical wave functions and hence unrelated to orbital hybridisation. While it is qualitative, VSEPR has a quantitative basis in quantum chemical topology methods such as the electron localization function. In 1957, Ronald Gillespie and Ronald Sydney Nyholm of University College London refined this concept into a detailed theory. In VSEPR theory, a bond or triple bond are treated as a single bonding group. The sum of the number of atoms bonded to a central atom, the number of electron pairs, therefore, determines the overall geometry that they will adopt. For example, when there are two electron pairs surrounding the atom, their mutual repulsion is minimal when they lie at opposite poles of the sphere. Therefore, the atom is predicted to adopt a linear geometry. If there are 3 electron pairs surrounding the atom, their repulsion is minimized by placing them at the vertices of an equilateral triangle centered on the atom. Therefore, the geometry is trigonal. Likewise, for 4 electron pairs, the arrangement is tetrahedral. The overall geometry is further refined by distinguishing between bonding and nonbonding electron pairs, the bonding electron pair shared in a sigma bond with an adjacent atom lies further from the central atom than a nonbonding pair of that atom, which is held close to its positively charged nucleus. VSEPR theory therefore views repulsion by the pair to be greater than the repulsion by a bonding pair. As such, when a molecule has 2 interactions with different degrees of repulsion, for instance, when 5 valence electron pairs surround a central atom, they adopt a trigonal bipyramidal molecular geometry with two collinear axial positions and three equatorial positions. The difference between pairs and bonding pairs may also be used to rationalize deviations from idealized geometries. For example, the H2O molecule has four pairs in its valence shell
25.
Valence bond theory
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In chemistry, valence bond theory is one of two basic theories, along with molecular orbital theory, that were developed to use the methods of quantum mechanics to explain chemical bonding. It focuses on how the orbitals of the dissociated atoms combine to give individual chemical bonds when a molecule is formed. In contrast, molecular orbital theory has orbitals that cover the whole molecule, in 1916, G. N. Lewis proposed that a chemical bond forms by the interaction of two shared bonding electrons, with the representation of molecules as Lewis structures. In 1927 the Heitler–London theory was formulated which for the first time enabled the calculation of bonding properties of the hydrogen molecule H2 based on quantum mechanical considerations. Specifically, Walter Heitler determined how to use Schrödingers wave equation to show how two hydrogen atom wavefunctions join together, with plus, minus, and exchange terms, to form a covalent bond. He then called up his associate Fritz London and they worked out the details of the theory over the course of the night. Later, Linus Pauling used the pair bonding ideas of Lewis together with Heitler–London theory to two other key concepts in VB theory, resonance and orbital hybridization. Resonance theory was criticized as imperfect by Soviet chemists during the 1950s, according to this theory a covalent bond is formed between the two atoms by the overlap of half filled valence atomic orbitals of each atom containing one unpaired electron. A valence bond structure is similar to a Lewis structure, but where a single Lewis structure cannot be written, each of these VB structures represents a specific Lewis structure. This combination of valence bond structures is the point of resonance theory. Valence bond theory considers that the atomic orbitals of the participating atoms form a chemical bond. Because of the overlapping, it is most probable that electrons should be in the bond region, valence bond theory views bonds as weakly coupled orbitals. Valence bond theory is typically easier to employ in ground state molecules, the inner-shell orbitals and electrons remain essentially unchanged during the formation of bonds. The overlapping atomic orbitals can differ, the two types of overlapping orbitals are sigma and pi. Sigma bonds occur when the orbitals of two shared electrons overlap head-to-head, pi bonds occur when two orbitals overlap when they are parallel. For example, a bond between two electrons is a sigma bond, because two spheres are always coaxial. In terms of order, single bonds have one sigma bond, double bonds consist of one sigma bond and one pi bond. However, the orbitals for bonding may be hybrids
26.
Orbital hybridisation
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In chemistry, hybridisation is the concept of mixing atomic orbitals into new hybrid orbitals suitable for the pairing of electrons to form chemical bonds in valence bond theory. Hybrid orbitals are very useful in the explanation of molecular geometry, although sometimes taught together with the valence shell electron-pair repulsion theory, valence bond and hybridisation are in fact not related to the VSEPR model. Chemist Linus Pauling first developed the theory in 1931 in order to explain the structure of simple molecules such as methane using atomic orbitals. In reality however, methane has four bonds of equivalent strength separated by the bond angle of 109. 5°. It gives a simple orbital picture equivalent to Lewis structures, hybridisation theory finds its use mainly in organic chemistry. Orbitals are a representation of the behaviour of electrons within molecules. In heavier atoms, such as carbon, nitrogen, and oxygen, hybrid orbitals are assumed to be mixtures of atomic orbitals, superimposed on each other in various proportions. For example, in methane, the C hybrid orbital which forms each carbon–hydrogen bond consists of 25% s character and 75% p character and is thus described as sp3 hybridised. Quantum mechanics describes this hybrid as an sp3 wavefunction of the form N, the ratio of coefficients is √3 in this example. Since the electron density associated with an orbital is proportional to the square of the wavefunction, the p character or the weight of the p component is N2λ2 = 3/4. The amount of p character or s character, which is decided mainly by orbital hybridisation, molecules with multiple bonds or multiple lone pairs can have orbitals represented in terms of sigma and pi symmetry or equivalent orbitals. The sigma and pi representation of Erich Hückel is the common one compared to the equivalent orbital representation of Linus Pauling. The two have mathematically equivalent total many-electron wave functions, and are related by a transformation of the set of occupied molecular orbitals. Hybridisation describes the bonding atoms from a point of view. For a tetrahedrally coordinated carbon, the carbon should have 4 orbitals with the symmetry to bond to the 4 hydrogen atoms. The carbon atom can also bond to four hydrogen atoms by an excitation of an electron from the doubly occupied 2s orbital to the empty 2p orbital, producing four singly occupied orbitals. The energy released by formation of two additional bonds more than compensates for the energy required, energetically favouring the formation of four C-H bonds. Quantum mechanically, the lowest energy is obtained if the four bonds are equivalent, a set of four equivalent orbitals can be obtained that are linear combinations of the valence-shell s and p wave functions, which are the four sp3 hybrids
27.
Resonance (chemistry)
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In chemistry, resonance or mesomerism is a way of describing delocalized electrons within certain molecules or polyatomic ions where the bonding cannot be expressed by one single Lewis structure. A molecule or ion with such delocalized electrons is represented by several contributing structures, each contributing structure can be represented by a Lewis structure, with only an integer number of covalent bonds between each pair of atoms within the structure. Several Lewis structures are used collectively to describe the molecular structure. Electron delocalization lowers the energy of the substance and thus makes it more stable than any of the contributing structures. The difference between the energy of the actual structure and that of the contributing structure with the lowest potential energy is called the resonance energy or delocalization energy. An isomer is a molecule with the chemical formula but with different arrangements of atoms in space. Resonance contributors of a molecule, on the contrary, can differ by the arrangements of electrons. Therefore, the resonance hybrid cannot be represented by a combination of isomers, benzene undergoes substitution reactions, rather than addition reactions as typical for alkenes. He proposed that the bond in benzene is intermediate of a single and double bond. The mechanism of resonance was introduced into quantum mechanics by Werner Heisenberg in 1926 in a discussion of the states of the helium atom. He compared the structure of the atom with the classical system of resonating coupled harmonic oscillators. Linus Pauling used this mechanism to explain the partial valence of molecules in 1928, the alternative term mesomerism popular in German and French publications with the same meaning was introduced by C. K. Ingold in 1938, but did not catch on in the English literature. The current concept of effect has taken on a related. The double headed arrow was introduced by the German chemist Fritz Arndt who preferred the German phrase zwischenstufe or intermediate stage, the real structure is an intermediate of these structures represented by a resonance hybrid. The contributing structures are not isomers and they differ only in the position of electrons, not in the position of nuclei. Each Lewis formula must have the number of valence electrons. Bonds that have different bond orders in different contributing structures do not have typical bond lengths, the real structure has a lower total potential energy than each of the contributing structures would have. This means that it is more stable than each separate contributing structure would be and it is a common misconception that resonance structures are actual transient states of the molecule, with the molecule oscillating between them or existing as an equilibrium between them
28.
Molecular orbital theory
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The spatial and energetic properties of electrons within atoms are fixed by quantum mechanics to form orbitals that contain these electrons. While atomic orbitals contain electrons ascribed to an atom, molecular orbitals. These approximations are now made by applying the density functional theory or Hartree–Fock models to the Schrödinger equation, one may determine cij coefficients numerically by substituting this equation into the Schrödingers equation and applying the variational principle. The variational principle is a technique used in quantum mechanics to build up the coefficients of each atomic orbital basis. A larger coefficient means that the basis is composed more of that particular contributing atomic orbital—hence. This method of quantifying orbital contribution as Linear Combinations of Atomic Orbitals is used in computational chemistry, an additional unitary transformation can be applied on the system to accelerate the convergence in some computational schemes. Molecular orbital theory was seen as a competitor to valence bond theory in the 1930s, Molecular orbital theory was developed, in the years after valence bond theory had been established, primarily through the efforts of Friedrich Hund, Robert Mulliken, John C. MO theory was originally called the Hund-Mulliken theory, according to German physicist and physical chemist Erich Hückel, the first quantitative use of molecular orbital theory was the 1929 paper of Lennard-Jones. This paper notably predicted a triplet ground state for the molecule which explained its paramagnetism before valence bond theory. The word orbital was introduced by Mulliken in 1932, by 1933, the molecular orbital theory had been accepted as a valid and useful theory. This method provided an explanation of the stability of molecules with six pi-electrons such as benzene, the first accurate calculation of a molecular orbital wavefunction was that made by Charles Coulson in 1938 on the hydrogen molecule. This rigorous approach is known as the Hartree–Fock method for molecules although it had its origins in calculations on atoms, in calculations on molecules, the molecular orbitals are expanded in terms of an atomic orbital basis set, leading to the Roothaan equations. This led to the development of ab initio quantum chemistry methods. In parallel, molecular orbital theory was applied in a more approximate manner using some empirically derived parameters in methods now known as quantum chemistry methods. The success of Molecular Orbital Theory also spawned ligand field theory, Molecular orbital theory uses a linear combination of atomic orbitals to represent molecular orbitals resulting from bonds between atoms. These are often divided into bonding orbitals, anti-bonding orbitals, an anti-bonding orbital concentrates electron density behind each nucleus, and so tends to pull each of the two nuclei away from the other and actually weaken the bond between the two nuclei. Molecular orbitals are further divided according to the types of atomic orbitals they are formed from, chemical substances will form bonding interactions if their orbitals become lower in energy when they interact with each other. Different bonding orbitals are distinguished that differ by electron configuration and by energy levels, the molecular orbitals of a molecule can be illustrated in molecular orbital diagrams
29.
Linear combination of atomic orbitals
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A linear combination of atomic orbitals or LCAO is a quantum superposition of atomic orbitals and a technique for calculating molecular orbitals in quantum chemistry. In quantum mechanics, electron configurations of atoms are described as wavefunctions, in mathematical sense, these wave functions are the basis set of functions, the basis functions, which describe the electrons of a given atom. In chemical reactions, orbital wavefunctions are modified, i. e. the electron cloud shape is changed, an initial assumption is that the number of molecular orbitals is equal to the number of atomic orbitals included in the linear expansion. In a sense, n atomic orbitals combine to form n molecular orbitals, which can be numbered i =1 to n, the coefficients are the weights of the contributions of the n atomic orbitals to the molecular orbital. The Hartree–Fock procedure is used to obtain the coefficients of the expansion, in either case the basis functions are usually also referred to as atomic orbitals. By minimizing the energy of the system, an appropriate set of coefficients of the linear combinations is determined. This quantitative approach is now known as the Hartree–Fock method, the graphs that are plotted to make this discussion clearer are called correlation diagrams. The required atomic orbital energies can come from calculations or directly from experiment via Koopmans theorem and this is done by using the symmetry of the molecules and orbitals involved in bonding, and thus is sometimes called Symmetry Adapted Linear Combination. The first step in process is assigning a point group to the molecule. A common example is water, which is of C2v symmetry, then a reducible representation of the bonding is determined demonstrated below for water, Each operation in the point group is performed upon the molecule. The number of bonds that are unmoved is the character of that operation and this reducible representation is decomposed into the sum of irreducible representations. These irreducible representations correspond to the symmetry of the orbitals involved, MO diagrams provide simple qualitative LCAO treatment. Quantitative theories are the Hückel method, the extended Hückel method, quantum chemistry computer programs Hartree–Fock method Basis set Tight binding Holstein–Herring method LCAO @ chemistry. umeche. maine. edu Link
30.
Ligand field theory
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Ligand field theory describes the bonding, orbital arrangement, and other characteristics of coordination complexes. It represents an application of molecular orbital theory to transition metal complexes, a transition metal ion has nine valence atomic orbitals - consisting of five nd, three p, and one s orbitals. These orbitals are of appropriate energy to form bonding interaction with ligands, the LFT analysis is highly dependent on the geometry of the complex, but most explanations begin by describing octahedral complexes, where six ligands coordinate to the metal. Other complexes can be described by reference to field theory. Griffith and Orgel championed ligand field theory as an accurate description of such complexes. In their paper, they propose that the cause of color differences in transition metal complexes in solution is the incomplete d orbital subshells. That is, the d orbitals of transition metals participate in bonding. In an octahedral complex, the orbitals created by coordination can be seen as resulting from the donation of two electrons by each of six σ-donor ligands to the d-orbitals on the metal. In octahedral complexes, ligands approach along the x-, y- and z-axes, so their σ-symmetry orbitals form bonding and anti-bonding combinations with the dz2, the dxy, dxz and dyz orbitals remain non-bonding orbitals. The irreducible representations that these span are a1g, t1u and eg, the six σ-bonding molecular orbitals result from the combinations of ligand SALCs with metal orbitals of the same symmetry. π bonding in octahedral complexes occurs in two ways, via any ligand p-orbitals that are not being used in σ bonding, and via any π or π* molecular orbitals present on the ligand. In the usual analysis, the p-orbitals of the metal are used for σ bonding, so the π interactions take place with the appropriate metal d-orbitals, i. e. dxy, dxz and these are the orbitals that are non-bonding when only σ bonding takes place. One important π bonding in coordination complexes is metal-to-ligand π bonding and it occurs when the LUMOs of the ligand are anti-bonding π* orbitals. These orbitals are close in energy to the dxy, dxz and dyz orbitals, the ligands end up with electrons in their π* molecular orbital, so the corresponding π bond within the ligand weakens. The other form of coordination π bonding is ligand-to-metal bonding and this situation arises when the π-symmetry p or π orbitals on the ligands are filled. They combine with the dxy, dxz and dyz orbitals on the metal and it is filled with electrons from the metal d-orbitals, however, becoming the HOMO of the complex. For that reason, ΔO decreases when ligand-to-metal bonding occurs, the greater stabilization that results from metal-to-ligand bonding is caused by the donation of negative charge away from the metal ion, towards the ligands. This allows the metal to accept the σ bonds more easily, the combination of ligand-to-metal σ-bonding and metal-to-ligand π-bonding is a synergic effect, as each enhances the other
31.
Electrostatics
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Electrostatics is a branch of physics that deals with the phenomena and properties of stationary or slow-moving electric charges. Since classical physics, it has known that some materials such as amber attract lightweight particles after rubbing. The Greek word for amber, ήλεκτρον, or electron, was the source of the word electricity, Electrostatic phenomena arise from the forces that electric charges exert on each other. Such forces are described by Coulombs law, Electrostatics involves the buildup of charge on the surface of objects due to contact with other surfaces. This is because the charges that transfer are trapped there for a long enough for their effects to be observed. We begin with the magnitude of the force between two point charges q and Q. It is convenient to one of these charges, q, as a test charge. As we develop the theory, more source charges will be added.854187817 ×10 −12 C2 N −1 m −2, the SI units of ε0 are equivalently A2s4 kg−1m−3 or C2N−1m−2 or F m−1. Coulombs constant is, k e ≈14 π ε0 ≈8.987551787 ×109 N m 2 C −2. A single proton has a charge of e, and the electron has a charge of −e and these physical constants are currently defined so that ε0 and k0 are exactly defined, and e is a measured quantity. Electric field lines are useful for visualizing the electric field, field lines begin on positive charge and terminate on negative charge. Electric field lines are parallel to the direction of the field. The electric field, E →, is a field that can be defined everywhere. It is convenient to place a hypothetical test charge at a point, by Coulombs Law, this test charge will experience a force that can be used to define the electric field as follow F → = q E →. For a single point charge at the origin, the magnitude of electric field is E = k e Q / R2. The fact that the force can be calculated by summing all the contributions due to individual source particles is an example of the superposition principle. If the charge is distributed over a surface or along a line, the Divergence Theorem allows Gausss Law to be written in differential form, ∇ → ⋅ E → = ρ ε0. Where ∇ → ⋅ is the divergence operator, the definition of electrostatic potential, combined with the differential form of Gausss law, provides a relationship between the potential Φ and the charge density ρ, ∇2 ϕ = − ρ ε0
32.
Valence electron
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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
33.
Condensed matter physics
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Condensed matter physics is a branch of physics that deals with the physical properties of condensed phases of matter, where particles adhere to each other. Condensed matter physicists seek to understand the behavior of these phases by using physical laws, in particular, they include the laws of quantum mechanics, electromagnetism and statistical mechanics. The field overlaps with chemistry, materials science, and nanotechnology, the theoretical physics of condensed matter shares important concepts and methods with that of particle physics and nuclear physics. A variety of topics in physics such as crystallography, metallurgy, elasticity, magnetism, etc. were treated as distinct areas until the 1940s, when they were grouped together as solid state physics. Around the 1960s, the study of properties of liquids was added to this list, forming the basis for the new. The Bell Telephone Laboratories was one of the first institutes to conduct a program in condensed matter physics. References to condensed state can be traced to earlier sources, as a matter of fact, it would be more correct to unify them under the title of condensed bodies. One of the first studies of condensed states of matter was by English chemist Humphry Davy, Davy observed that of the forty chemical elements known at the time, twenty-six had metallic properties such as lustre, ductility and high electrical and thermal conductivity. This indicated that the atoms in John Daltons atomic theory were not indivisible as Dalton claimed, Davy further claimed that elements that were then believed to be gases, such as nitrogen and hydrogen could be liquefied under the right conditions and would then behave as metals. In 1823, Michael Faraday, then an assistant in Davys lab, successfully liquefied chlorine and went on to all known gaseous elements, except for nitrogen, hydrogen. By 1908, James Dewar and Heike Kamerlingh Onnes were successfully able to hydrogen and then newly discovered helium. Paul Drude in 1900 proposed the first theoretical model for an electron moving through a metallic solid. Drudes model described properties of metals in terms of a gas of free electrons, the phenomenon completely surprised the best theoretical physicists of the time, and it remained unexplained for several decades. Drudes classical model was augmented by Wolfgang Pauli, Arnold Sommerfeld, Felix Bloch, Pauli realized that the free electrons in metal must obey the Fermi–Dirac statistics. Using this idea, he developed the theory of paramagnetism in 1926, shortly after, Sommerfeld incorporated the Fermi–Dirac statistics into the free electron model and made it better able to explain the heat capacity. Two years later, Bloch used quantum mechanics to describe the motion of an electron in a periodic lattice. Magnetism as a property of matter has been known in China since 4000 BC, Pierre Curie studied the dependence of magnetization on temperature and discovered the Curie point phase transition in ferromagnetic materials. In 1906, Pierre Weiss introduced the concept of magnetic domains to explain the properties of ferromagnets
34.
Chemical polarity
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In chemistry, polarity is a separation of electric charge leading to a molecule or its chemical groups having an electric dipole or multipole moment. Polar molecules must contain polar bonds due to a difference in electronegativity between the bonded atoms, a polar molecule with two or more polar bonds must have an asymmetric geometry so that the bond dipoles do not cancel each other. Polar molecules interact through dipole–dipole intermolecular forces and hydrogen bonds, Polarity underlies a number of physical properties including surface tension, solubility, and melting and boiling points. Not all atoms attract electrons with the same force, the amount of pull an atom exerts on its electrons is called its electronegativity. Atoms with high electronegativities – such as fluorine, oxygen and nitrogen – exert a pull on electrons than atoms with lower electronegativities. In a bond, this leads to sharing of electrons between the atoms, as electrons will be drawn closer to the atom with the higher electronegativity. Because electrons have a charge, the unequal sharing of electrons within a bond leads to the formation of an electric dipole. Because the amount of charge separated in such dipoles is usually smaller than a charge, they are called partial charges, denoted as δ+. These symbols were introduced by Christopher Kelk Ingold and Edith Hilda Ingold in 1926, the bond dipole moment is calculated by multiplying the amount of charge separated and the distance between the charges. These dipoles within molecules can interact with dipoles in other molecules, Bonds can fall between one of two extremes – being completely nonpolar or completely polar. A completely nonpolar bond occurs when the electronegativities are identical and therefore possess a difference of zero, a completely polar bond is more correctly called an ionic bond, and occurs when the difference between electronegativities is large enough that one atom actually takes an electron from the other. The terms polar and nonpolar are usually applied to covalent bonds, to determine the polarity of a covalent bond using numerical means, the difference between the electronegativity of the atoms is used. Bond polarity is typically divided into three groups that are based on the difference in electronegativity between the two bonded atoms. He estimated that a difference of 1.7 corresponds to 50% ionic character, see also dipole § Molecular dipoles. While the molecules can be described as covalent, nonpolar covalent, or ionic. However, the properties are typical of such molecules. A molecule is composed of one or more chemical bonds between molecular orbitals of different atoms, a polar molecule has a net dipole as a result of the opposing charges from polar bonds arranged asymmetrically. Water is an example of a polar molecule since it has a positive charge on one side
35.
Nylon
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Nylon is a generic designation for a family of synthetic polymers, more specifically aliphatic or semi-aromatic polyamides. They can be melt-processed into fibers, films or shapes, the first example of nylon was produced on February 28,1935, by Wallace Carothers at DuPonts research facility at the DuPont Experimental Station. Nylon polymers have significant commercial applications in fibers, in shapes. Nylon is made of repeating units linked by bonds and is a type of polyamide and is frequently referred to as such. Nylon was the first commercially successful synthetic thermoplastic polymer, commercially, nylon polymer is made by reacting monomers which are either lactams, acid/amines or stoichiometric mixtures of diamines and diacids. Mixtures of these can be polymerized together to make copolymers, Nylon polymers can be mixed with a wide variety of additives to achieve many different property variations. Nylon was invented accidentally by Julian W. Hill, a chemist for DuPont under Wallace Carotherss supervision and it was only given this name at the 1939 New York Worlds Fair. The patents were owned by DuPont, Nylon was intended to be a synthetic replacement for silk and substituted for it in many different products after silk became scarce during World War II. It replaced silk in military applications such as parachutes and flak vests, after initial commercialization of nylon as a fiber, applications in the form of shapes and films were also developed. The main market for nylon shapes now is in auto components, in 1940, John W. Eckelberry of DuPont stated that the letters nyl were arbitrary and the on was copied from the suffixes of other fibers such as cotton and rayon. A later publication by DuPont explained that the name was intended to be No-Run. Since the products were not really run-proof, the vowels were swapped to produce nuron, for clarity in pronunciation, the i was changed to y. Most nylons are made from the reaction of an acid with a diamine or a lactam or amino acid with itself. In the first case, the structure is so-called ABAB similar to polyesters and polyurethanes, in the second case, the repeating unit corresponds to the single monomer. It is difficult to get the proportions exactly correct, and deviations can lead to termination at molecular weights less than a desirable 10,000 daltons. To overcome this problem, a crystalline, solid nylon salt can be formed at room temperature, using an exact 1,1 ratio of the acid, the salt is crystallized to purify it and obtain the desired precise stoichiometry. Heated to 285 °C, the salt reacts to form nylon polymer with the production of water, Wallace Carothers at DuPont patented nylon 66, but overlooked the possibility to use lactams. That synthetic route was developed by Paul Schlack at IG Farben, leading to nylon 6, the peptide bond within the caprolactam is broken with the exposed active groups on each side being incorporated into two new bonds as the monomer becomes part of the polymer backbone
36.
Diamond
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Diamond is a metastable allotrope of carbon, where the carbon atoms are arranged in a variation of the face-centered cubic crystal structure called a diamond lattice. Diamond is less stable than graphite, but the rate from diamond to graphite is negligible at standard conditions. Diamond is renowned as a material with superlative physical qualities, most of which originate from the covalent bonding between its atoms. In particular, diamond has the highest hardness and thermal conductivity of any bulk material and those properties determine the major industrial application of diamond in cutting and polishing tools and the scientific applications in diamond knives and diamond anvil cells. Because of its extremely rigid lattice, it can be contaminated by very few types of impurities, such as boron, small amounts of defects or impurities color diamond blue, yellow, brown, green, purple, pink, orange or red. Diamond also has relatively high optical dispersion, most natural diamonds are formed at high temperature and pressure at depths of 140 to 190 kilometers in the Earths mantle. Carbon-containing minerals provide the source, and the growth occurs over periods from 1 billion to 3.3 billion years. Diamonds are brought close to the Earths surface through deep volcanic eruptions by magma, Diamonds can also be produced synthetically in a HPHT method which approximately simulates the conditions in the Earths mantle. An alternative, and completely different growth technique is chemical vapor deposition, several non-diamond materials, which include cubic zirconia and silicon carbide and are often called diamond simulants, resemble diamond in appearance and many properties. Special gemological techniques have developed to distinguish natural diamonds, synthetic diamonds. The word is from the ancient Greek ἀδάμας – adámas unbreakable, the name diamond is derived from the ancient Greek αδάμας, proper, unalterable, unbreakable, untamed, from ἀ-, un- + δαμάω, I overpower, I tame. Diamonds have been known in India for at least 3,000 years, Diamonds have been treasured as gemstones since their use as religious icons in ancient India. Their usage in engraving tools also dates to early human history, later in 1797, the English chemist Smithson Tennant repeated and expanded that experiment. By demonstrating that burning diamond and graphite releases the same amount of gas, the most familiar uses of diamonds today are as gemstones used for adornment, a use which dates back into antiquity, and as industrial abrasives for cutting hard materials. The dispersion of light into spectral colors is the primary gemological characteristic of gem diamonds. In the 20th century, experts in gemology developed methods of grading diamonds, four characteristics, known informally as the four Cs, are now commonly used as the basic descriptors of diamonds, these are carat, cut, color, and clarity. A large, flawless diamond is known as a paragon and these conditions are met in two places on Earth, in the lithospheric mantle below relatively stable continental plates, and at the site of a meteorite strike. The conditions for diamond formation to happen in the mantle occur at considerable depth corresponding to the requirements of temperature and pressure
37.
Quartz
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Quartz is the second most abundant mineral in Earths continental crust, behind feldspar. There are many different varieties of quartz, several of which are semi-precious gemstones, since antiquity, varieties of quartz have been the most commonly used minerals in the making of jewelry and hardstone carvings, especially in Eurasia. The word quartz is derived from the German word Quarz and its Middle High German ancestor twarc, the Ancient Greeks referred to quartz as κρύσταλλος derived from the Ancient Greek κρύος meaning icy cold, because some philosophers apparently believed the mineral to be a form of supercooled ice. Today, the rock crystal is sometimes used as an alternative name for the purest form of quartz. Quartz belongs to the crystal system. The ideal crystal shape is a six-sided prism terminating with six-sided pyramids at each end, well-formed crystals typically form in a bed that has unconstrained growth into a void, usually the crystals are attached at the other end to a matrix and only one termination pyramid is present. However, doubly terminated crystals do occur where they develop freely without attachment, a quartz geode is such a situation where the void is approximately spherical in shape, lined with a bed of crystals pointing inward. α-quartz crystallizes in the crystal system, space group P3121 and P3221 respectively. β-quartz belongs to the system, space group P6222 and P6422. These space groups are truly chiral, both α-quartz and β-quartz are examples of chiral crystal structures composed of achiral building blocks. The transformation between α- and β-quartz only involves a comparatively minor rotation of the tetrahedra with respect to one another, although many of the varietal names historically arose from the color of the mineral, current scientific naming schemes refer primarily to the microstructure of the mineral. Color is an identifier for the cryptocrystalline minerals, although it is a primary identifier for the macrocrystalline varieties. Pure quartz, traditionally called rock crystal or clear quartz, is colorless and transparent or translucent, common colored varieties include citrine, rose quartz, amethyst, smoky quartz, milky quartz, and others. The most important distinction between types of quartz is that of macrocrystalline and the microcrystalline or cryptocrystalline varieties, the cryptocrystalline varieties are either translucent or mostly opaque, while the transparent varieties tend to be macrocrystalline. Chalcedony is a form of silica consisting of fine intergrowths of both quartz, and its monoclinic polymorph moganite. Other opaque gemstone varieties of quartz, or mixed rocks including quartz, often including contrasting bands or patterns of color, are agate, carnelian or sard, onyx, heliotrope, amethyst is a form of quartz that ranges from a bright to dark or dull purple color. The worlds largest deposits of amethysts can be found in Brazil, Mexico, Uruguay, Russia, France, Namibia, sometimes amethyst and citrine are found growing in the same crystal. It is then referred to as ametrine, an amethyst is formed when there is iron in the area where it was formed
38.
Silicate minerals
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Silicate minerals are rock-forming minerals made up of silicate groups. They are the largest and most important class of rock-forming minerals and they are classified based on the structure of their silicate groups, which contain different ratios of silicon and oxygen. Nesosilicates, or orthosilicates, have the orthosilicate ion, which constitute isolated 4− tetrahedra that are connected only by interstitial cations and these exist as 3-member 6− and 6-member 12− rings, where T stands for a tetrahedrally coordinated cation. Inosilicates, or chain silicates, have interlocking chains of silicate tetrahedra with either SiO3,1,3 ratio, for single chains or Si4O11,4,11 ratio, for double chains. Nickel–Strunz classification,09. D Pyroxene group Enstatite – orthoferrosilite series Enstatite – MgSiO3 Ferrosilite – FeSiO3 Pigeonite – Ca0.251, all phyllosilicate minerals are hydrated, with either water or hydroxyl groups attached. Serpentine subgroup Antigorite – Mg3Si2O54 Chrysotile – Mg3Si2O54 Lizardite – Mg3Si2O54 Clay minerals group Halloysite – Al2Si2O54 Kaolinite – Al2Si2O54 Illite – 24O10 Montmorillonite –0 and this group comprises nearly 75% of the crust of the Earth. Tectosilicates, with the exception of the group, are aluminosilicates. Nickel–Strunz classification,09. F and 09. G,04. A, an introduction to the rock-forming minerals. Wise, W. S. Zussman, J. Rock-forming minerals, P.982 pp. Hurlbut, Cornelius S. Danas Manual of Mineralogy. Mindat. org, Dana classification Webmineral, Danas New Silicate Classification Media related to Silicates at Wikimedia Commons
39.
Atomic orbital
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In quantum mechanics, an atomic orbital is a mathematical function that describes the wave-like behavior of either one electron or a pair of electrons in an atom. This function can be used to calculate the probability of finding any electron of an atom in any specific region around the atoms nucleus. The term, atomic orbital, may refer to the physical region or space where the electron can be calculated to be present. Each such orbital can be occupied by a maximum of two electrons, each with its own quantum number s. The simple names s orbital, p orbital, d orbital and these names, together with the value of n, are used to describe the electron configurations of atoms. They are derived from the description by early spectroscopists of certain series of alkali metal spectroscopic lines as sharp, principal, diffuse, Orbitals for ℓ >3 continue alphabetically, omitting j because some languages do not distinguish between the letters i and j. Atomic orbitals are the building blocks of the atomic orbital model. In this model the electron cloud of an atom may be seen as being built up in an electron configuration that is a product of simpler hydrogen-like atomic orbitals. The lowest possible energy an electron can take is therefore analogous to the frequency of a wave on a string. Higher energy states are similar to harmonics of the fundamental frequency. The electrons are never in a point location, although the probability of interacting with the electron at a single point can be found from the wave function of the electron. Particle-like properties, There is always a number of electrons orbiting the nucleus. Electrons jump between orbitals in a particle-like fashion, for example, if a single photon strikes the electrons, only a single electron changes states in response to the photon. The electrons retain particle-like properties such as, each state has the same electrical charge as the electron particle. Each wave state has a single discrete spin and this can depend upon its superposition. Thus, despite the popular analogy to planets revolving around the Sun, in addition, atomic orbitals do not closely resemble a planets elliptical path in ordinary atoms. A more accurate analogy might be that of a large and often oddly shaped atmosphere, atomic orbitals exactly describe the shape of this atmosphere only when a single electron is present in an atom. This is due to the uncertainty principle, atomic orbitals may be defined more precisely in formal quantum mechanical language
40.
Delocalized electron
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In chemistry, delocalized electrons are electrons in a molecule, ion or solid metal that are not associated with a single atom or a covalent bond. The term is general and can have different meanings in different fields. In organic chemistry, this refers to resonance in conjugated systems, in solid-state physics, this refers to free electrons that facilitate electrical conduction. In quantum chemistry, this refers to molecular orbitals that extend over several adjacent atoms, in the simple aromatic ring of benzene the delocalization of six π electrons over the C6 ring is often graphically indicated by a circle. The fact that the six C-C bonds are equidistant is one indication of this delocalization, in valence bond theory, delocalization in benzene is represented by resonance structures. Delocalized electrons also exist in the structure of solid metals, metallic structure consists of aligned positive ions in a sea of delocalized electrons. This means that the electrons are free to move throughout the structure, in diamond all four outer electrons of each carbon atom are localized between the atoms in covalent bonding. The movement of electrons is restricted and diamond does not conduct an electric current, in graphite, each carbon atom uses only 3 of its 4 outer energy level electrons in covalently bonding to three other carbon atoms in a plane. Each carbon atom contributes one electron to a system of electrons that is also a part of the chemical bonding. The delocalized electrons are free to move throughout the plane, for this reason, graphite conducts electricity along the planes of carbon atoms, but does not conduct in a direction at right angles to the plane. Standard ab initio quantum chemistry methods lead to delocalized orbitals that, in general, localized orbitals may then be found as linear combinations of the delocalized orbitals, given by an appropriate unitary transformation. In the methane molecule for example, ab initio calculations show bonding character in four molecular orbitals, there are two orbital levels, a bonding molecular orbital formed from the 2s orbital on carbon and triply degenerate bonding molecular orbitals from each of the 2p orbitals on carbon. The localized sp3 orbitals corresponding to each individual bond in valence bond theory can be obtained from a combination of the four molecular orbitals
41.
Ultimate tensile strength
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In other words, tensile strength resists tension, whereas compressive strength resists compression. Ultimate tensile strength is measured by the stress that a material can withstand while being stretched or pulled before breaking. In the study of strength of materials, tensile strength, compressive strength, some materials break very sharply, without plastic deformation, in what is called a brittle failure. Others, which are more ductile, including most metals, experience some plastic deformation, the UTS is usually found by performing a tensile test and recording the engineering stress versus strain. The highest point of the curve is the UTS. It is a property, therefore its value does not depend on the size of the test specimen. However, it is dependent on other factors, such as the preparation of the specimen, the presence or otherwise of surface defects, Tensile strengths are rarely used in the design of ductile members, but they are important in brittle members. They are tabulated for common materials such as alloys, composite materials, ceramics, plastics, Tensile strength can be defined for liquids as well as solids under certain conditions. Tensile strength is defined as a stress, which is measured as force per unit area, for some non-homogeneous materials it can be reported just as a force or as a force per unit width. In the International System of Units, the unit is the pascal, or, equivalently to pascals, Many materials can display linear elastic behavior, defined by a linear stress–strain relationship, as shown in the left figure up to point 3. Beyond this elastic region, for materials, such as steel. A plastically deformed specimen does not completely return to its original size, for many applications, plastic deformation is unacceptable, and is used as the design limitation. The reversal point is the stress on the engineering stress–strain curve. The UTS is not used in the design of static members because design practices dictate the use of the yield stress. It is, however, used for quality control, because of the ease of testing and it is also used to roughly determine material types for unknown samples. The UTS is a common engineering parameter to design members made of material because such materials have no yield point. Typically, the testing involves taking a sample with a fixed cross-sectional area. When testing some metals, indentation hardness correlates linearly with tensile strength and this important relation permits economically important nondestructive testing of bulk metal deliveries with lightweight, even portable equipment, such as hand-held Rockwell hardness testers
42.
Lustre (mineralogy)
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Lustre or luster is the way light interacts with the surface of a crystal, rock, or mineral. The word traces its origins back to the latin lux, meaning light, a range of terms are used to describe lustre, such as earthy, metallic, greasy, and silky. Similarly, the term refers to a glassy lustre. A list of terms is given below. Lustre varies over a continuum, and so there are no rigid boundaries between the different types of lustre. The terms are frequently combined to describe types of lustre. Some minerals exhibit unusual optical phenomena, such as asterism or chatoyancy, a list of such phenomena is given below. Adamantine minerals possess a superlative lustre, which is most notably seen in diamond, such minerals are transparent or translucent, and have a high refractive index. Minerals with an adamantine lustre are uncommon, with examples being cerussite. Minerals with a degree of lustre are referred to as subadamantine, with some examples being garnet. Dull minerals exhibit little to no lustre, due to coarse granulations which scatter light in all directions, a distinction is sometimes drawn between dull minerals and earthy minerals, with the latter being coarser, and having even less lustre. Greasy minerals resemble fat or grease, a greasy lustre often occurs in minerals containing a great abundance of microscopic inclusions, with examples including opal and cordierite. Many minerals with a greasy lustre also feel greasy to the touch, metallic minerals have the lustre of polished metal, and with ideal surfaces will work as a reflective surface. Examples include galena, pyrite and magnetite, pearly minerals consist of thin transparent co-planar sheets. Light reflecting from these layers give them a lustre reminiscent of pearls, such minerals possess perfect cleavage, with examples including muscovite and stilbite. Resinous minerals have the appearance of resin, chewing gum or plastic, a principal example is amber, which is a form of fossilized resin. Silky minerals have an arrangement of extremely fine fibres, giving them a lustre reminiscent of silk. Examples include asbestos, ulexite and the satin spar variety of gypsum, a fibrous lustre is similar, but has a coarser texture
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History of chemistry
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The history of alchemy represents a time span from ancient history to the present. By 1000 BC, civilizations used technologies that would form the basis of the various branches of chemistry. The protoscience of chemistry, alchemy, was unsuccessful in explaining the nature of matter, however, by performing experiments and recording the results, alchemists set the stage for modern chemistry. The distinction began to emerge when a clear differentiation was made between chemistry and alchemy by Robert Boyle in his work The Sceptical Chymist, while both alchemy and chemistry are concerned with matter and its transformations, chemists are seen as applying scientific method to their work. The history of chemistry is intertwined with the history of thermodynamics, the earliest recorded metal employed by humans seems to be gold which can be found free or native. Small amounts of gold have been found in Spanish caves used during the late Paleolithic period. Silver, copper, tin and meteoric iron can also be found native, egyptian weapons made from meteoric iron in about 3000 BC were highly prized as Daggers from Heaven. Arguably the first chemical reaction used in a manner was fire. However, for fire was seen simply as a mystical force that could transform one substance into another while producing heat. Fire affected many aspects of early societies and these ranged from the simplest facets of everyday life, such as cooking and habitat lighting, to more advanced technologies, such as pottery, bricks, and melting of metals to make tools. It was fire that led to the discovery of glass and the purification of metals which in turn gave way to the rise of metallurgy. During the early stages of metallurgy, methods of purification of metals were sought, certain metals can be recovered from their ores by simply heating the rocks in a fire, notably tin, lead and copper, a process known as smelting. The first evidence of this extractive metallurgy dates from the 5th and 6th millennium BC, to date, the earliest copper smelting is found at the Belovode site, these examples include a copper axe from 5500 BC belonging to the Vinča culture. Other signs of early metals are found from the third millennium BC in places like Palmela, Los Millares, however, as often happens with the study of prehistoric times, the ultimate beginnings cannot be clearly defined and new discoveries are continuous and ongoing. These first metals were single ones or as found, by combining copper and tin, a superior metal could be made, an alloy called bronze, a major technological shift which began the Bronze Age about 3500 BC. These naturally occurring ores typically included arsenic as a common impurity, copper/tin ores are rare, as reflected in the fact that there were no tin bronzes in western Asia before 3000 BC. After the Bronze Age, the history of metallurgy was marked by armies seeking better weaponry, countries in Eurasia prospered when they made the superior alloys, which, in turn, made better armor and better weapons. This often determined the outcomes of battles, significant progress in metallurgy and alchemy was made in ancient India
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History of molecular theory
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In chemistry, the history of molecular theory traces the origins of the concept or idea of the existence of strong chemical bonds between two or more atoms. The modern concept of molecules can be traced back towards pre-scientific and Greek philosophers such as Leucippus who argued that all the universe is composed of atoms, circa 450 BC Empedocles imagined fundamental elements and forces of attraction and repulsion allowing the elements to interact. Prior to this, Heraclitus had claimed that fire or change was fundamental to our existence, a fifth element, the incorruptible quintessence aether, was considered to be the fundamental building block of the heavenly bodies. The viewpoint of Leucippus and Empedocles, along with the aether, was accepted by Aristotle and passed to medieval and it was Democritus that was the main proponent of this view. Moreover, connections were explained by material links in which atoms were supplied with attachments, some with hooks and eyes others with balls. With the rise of scholasticism and the decline of the Roman Empire, the 17th century, however, saw a resurgence in the atomic theory primarily through the works of Gassendi, and Newton. Among other scientists of that time Gassendi deeply studied ancient history and he reasoned that to account for the size and shape of atoms moving in a void could account for the properties of matter. Heat was due to small, round atoms, cold, to atoms with sharp points, which accounted for the pricking sensation of severe cold. Newton, though he acknowledged the various atom attachment theories in vogue at the time, i. e. ”A molecule, according to this view, consisted of corpuscles united through a geometric locking of points and pores. An early precursor to the idea of bonded combinations of atoms, was the theory of combination via chemical affinity, geoffroys name is best known in connection with his tables of affinities, which he presented to the French Academy in 1718 and 1720. These were lists, prepared by collating observations on the actions of substances one upon another and these tables retained their vogue for the rest of the century, until displaced by the profounder conceptions introduced by CL Berthollet. In 1738, Swiss physicist and mathematician Daniel Bernoulli published Hydrodynamica, in 1789, William Higgins published views on what he called combinations of ultimate particles, which foreshadowed the concept of valency bonds. Interestingly, Dalton incorrectly imagined that atoms “hooked” together to form molecules. e, note that this quote is not a literal translation. Avogadro uses the name molecule for both atoms and molecules, specifically, he uses the name elementary molecule when referring to atoms and to complicate the matter also speaks of compound molecules and composite molecules. Hence, relative molecular masses could now be calculated from the masses of gas samples, Avogadro developed this hypothesis in order to reconcile Joseph Louis Gay-Lussacs 1808 law on volumes and combining gases with Daltons 1803 atomic theory. Dalton, by contrast, did not consider this possibility, curiously, Avogadro considers only molecules containing even numbers of atoms, he does not say why odd numbers are left out. He proposed that atoms were tetravalent, and could bond to themselves to form the carbon skeletons of organic molecules. In 1856, Scottish chemist Archibald Couper began research on the bromination of benzene at the laboratory of Charles Wurtz in Paris, one month after Kekulés second paper appeared, Coupers independent and largely identical theory of molecular structure was published
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Chemical affinity
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In chemical physics and physical chemistry, chemical affinity is the electronic property by which dissimilar chemical species are capable of forming chemical compounds. Chemical affinity can also refer to the tendency of an atom or compound to combine by chemical reaction with atoms or compounds of unlike composition, the idea of affinity is extremely old. Many attempts have been made at identifying its origins, the majority of such attempts, however, except in a general manner, end in futility since affinities lie at the basis of all magic, thereby pre-dating science. Physical chemistry, however, was one of the first branches of science to study, the name affinitas was first used in the sense of chemical relation by German philosopher Albertus Magnus near the year 1250. Later, those as Robert Boyle, John Mayow, Johann Glauber, Isaac Newton, the term affinity has been used figuratively since c.1600 in discussions of structural relationships in chemistry, philology, etc. and reference to natural attraction is from 1616. Chemical affinity, historically, has referred to the force that causes chemical reactions, as well as, more generally, and earlier, the ″tendency to combine″ of any pair of substances. The broad definition, used throughout history, is that chemical affinity is that whereby substances enter into or resist decomposition. Antoine Lavoisier, in his famed 1789 Traité Élémentaire de Chimie, refers to Bergman’s work, according to Prigogine, the term was introduced and developed by Théophile de Donder. Goethe used the concept in his novel Elective Affinities, the affinity concept was very closely linked to the visual representation of substances on a table. The first-ever affinity table, which was based on displacement reactions, was published in 1718 by the French chemist Étienne François Geoffroy, crucially, the table was the central graphic tool used to teach chemistry to students and its visual arrangement was often combined with other kinds diagrams. Joseph Black, for example, used the table in combination with chiastic, Affinity tables were used throughout Europe until the early 19th century when they were displaced by affinity concepts introduced by Claude Berthollet. In chemical physics and physical chemistry, chemical affinity is the property by which dissimilar chemical species are capable of forming chemical compounds. Chemical affinity can also refer to the tendency of an atom or compound to combine by chemical reaction with atoms or compounds of unlike composition, in modern terms, we relate affinity to the phenomenon whereby certain atoms or molecules have the tendency to aggregate or bond. In this antiquated context, chemical affinity is sometimes synonymous with the term magnetic attraction. Many writings, up until about 1925, also refer to a law of chemical affinity, ilya Prigogine summarized the concept of affinity, saying, All chemical reactions drive the system to a state of equilibrium in which the affinities of the reactions vanish. The present IUPAC definition is that affinity A is the partial derivative of Gibbs free energy G with respect to extent of reaction ξ at constant pressure and temperature. It follows that affinity is positive for spontaneous reactions, in 1923, the Belgian mathematician and physicist Théophile de Donder derived a relation between affinity and the Gibbs free energy of a chemical reaction. This definition is useful for quantifying the factors responsible both for the state of systems, and for changes of state of non-equilibrium systems
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Isaac Newton
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His book Philosophiæ Naturalis Principia Mathematica, first published in 1687, laid the foundations of classical mechanics. Newton also made contributions to optics, and he shares credit with Gottfried Wilhelm Leibniz for developing the infinitesimal calculus. Newtons Principia formulated the laws of motion and universal gravitation that dominated scientists view of the universe for the next three centuries. Newtons work on light was collected in his influential book Opticks. He also formulated a law of cooling, made the first theoretical calculation of the speed of sound. Newton was a fellow of Trinity College and the second Lucasian Professor of Mathematics at the University of Cambridge, politically and personally tied to the Whig party, Newton served two brief terms as Member of Parliament for the University of Cambridge, in 1689–90 and 1701–02. He was knighted by Queen Anne in 1705 and he spent the last three decades of his life in London, serving as Warden and Master of the Royal Mint and his father, also named Isaac Newton, had died three months before. Born prematurely, he was a child, his mother Hannah Ayscough reportedly said that he could have fit inside a quart mug. When Newton was three, his mother remarried and went to live with her new husband, the Reverend Barnabas Smith, leaving her son in the care of his maternal grandmother, Newtons mother had three children from her second marriage. From the age of twelve until he was seventeen, Newton was educated at The Kings School, Grantham which taught Latin and Greek. He was removed from school, and by October 1659, he was to be found at Woolsthorpe-by-Colsterworth, Henry Stokes, master at the Kings School, persuaded his mother to send him back to school so that he might complete his education. Motivated partly by a desire for revenge against a bully, he became the top-ranked student. In June 1661, he was admitted to Trinity College, Cambridge and he started as a subsizar—paying his way by performing valets duties—until he was awarded a scholarship in 1664, which guaranteed him four more years until he would get his M. A. He set down in his notebook a series of Quaestiones about mechanical philosophy as he found it, in 1665, he discovered the generalised binomial theorem and began to develop a mathematical theory that later became calculus. Soon after Newton had obtained his B. A. degree in August 1665, in April 1667, he returned to Cambridge and in October was elected as a fellow of Trinity. Fellows were required to become ordained priests, although this was not enforced in the restoration years, however, by 1675 the issue could not be avoided and by then his unconventional views stood in the way. Nevertheless, Newton managed to avoid it by means of a special permission from Charles II. A and he was elected a Fellow of the Royal Society in 1672. Newtons work has been said to distinctly advance every branch of mathematics then studied and his work on the subject usually referred to as fluxions or calculus, seen in a manuscript of October 1666, is now published among Newtons mathematical papers
. Bibcode:2005PNAS..10211985S. doi:10.1073/pnas.0505778102.
