1.8 Ions and salts

Main article: Ion
An ion is a charged species, an atom or a molecule, that has lost or gained one or more electrons. Positively charged cations (e.g. sodium cation Na+) and negatively charged anions (e.g. chloride Cl−) can form a crystalline lattice of neutral salts (e.g. sodium chloride NaCl). Examples of polyatomic ions that do not split up during acid-base reactions are hydroxide (OH−) and phosphate (PO43−).
Ions in the gaseous phase are often known as plasma.

An ion is an atom or molecule with more or less electrons than usual, giving it a positive or negative electric charge. Because an ion "wants" to become neutral by acquiring or losing additional electrons, it has interesting chemical features. Ions usually come in the form of solutions. There are three types -- anions, which are negatively charged, cations, which are positively charged, and radicals, or radical ions, which are highly reactive due to having a large amount of free electrons.

An ion solution is typically created when certain molecules, such as salt, are dissolved in water. The interactions between the solute and solvent molecules disassociate weakly-bonded molecules, like salt, into two or more ions. When salt is put into water, it breaks into sodium and chloride ions, both of which are essential to human life. If a person is deprived of salt for too long, they will eventually die. Other life-essential ions (called electrolytes) are potassium (K+), calcium (Ca2+), magnesium (Mg2+), hydrogen phosphate (HPO42−), and hydrogen carbonate (HCO3−).

Other ion solutions besides dissolved salts are acids and bases. For instance, hydrochloric acid is formed when hydrogen chloride is dissolved in water. The hydrogen chloride breaks into two ions, H+ and Cl−. The H+ reacts with water to produce H3O+, hydronium ion, while the chlorine remains as a chloride ion. Because the ionization process is complete, hydrochloric acid is called a strong acid. The strength of acids is measured by how many H+ ions they have -- hydrochloric acid many, so it is very acidic. In fact, hydrochloric acid can dissolve through glass and all metals except iridium, though its corrosion speed is slow.

Another category of ion solutions are bases. Instead of being measured by the concentration of hydrogen ions, a base is measured by its concentration of hydroxide ions (OH-). One example of a base is potassium hydroxide, which is the chemical precursor of most soft and liquid soaps. When dissolved in water, potassium hydroxide produces a strong alkali solution. An alkali is a type of base. Potassium hydroxide has a number of interesting chemical properties -- one of them is its hygroscopic (water-attracting) nature, which makes it a powerful desiccant. When it reacts with water, the reaction is exothermic, meaning it generates significant heat.

The last major group of ions are radical ions, or just radicals. Examples of radicals include superoxide (O2-), which the immune system uses to kill invading microorganisms, and nitric oxide, which the body uses as an important signaling molecule. Like acids and bases, radicals are highly reactive and have a number of applications in chemistry. One unfortunate aspect of radicals is that trace amounts in the body tend to break down tissues over time, contributing to aging.

Salt:
In chemistry, salts are ionic compounds that can result from the neutralization reaction of an acid and a base. Salts are ionic compounds composed of cations (positively charged ions) and anions (negative ions) so that the product is electrically neutral (without a net charge). These component ions can be inorganic such as chloride (Cl−), as well as organic such as acetate (CH3COO−) and monatomic ions such as fluoride (F−), as well as polyatomic ions such as sulfate (SO42−).

There are several varieties of salts. Salts that hydrolyze to produce hydroxide ions when dissolved in water are basic salts and salts that hydrolyze to produce hydronium ions in water are acid salts. Neutral salts are those that are neither acid nor basic salts. Zwitterions contain an anionic center and a cationic center in the same molecule but are not considered to be salts. Examples include amino acids, many metabolites, peptides and proteins.

Molten salts and solutions containing dissolved salts (e.g. sodium chloride in water) are called electrolytes, as they are able to conduct electricity. As observed in the cytoplasm of cells, in blood, urine, plant saps and mineral waters, mixtures of many different ions in solution usually do not form defined salts after evaporation of the water. Therefore, their salt content is given for the respective ions.
Different salts can elicit all five basic tastes, e.g., salty (sodium chloride), sweet (lead diacetate, which will cause lead poisoning if ingested), sour (potassium bitartrate), bitter (magnesium sulfate), and umami or savory (monosodium glutamate).
Acid salts is a somewhat obscure term for a class of salts formed by the partial neutralization of diprotic or polyprotic acids. Because the parent acid is only partially neutralized, one or more replaceable protons remain. Typical acid salts have one or more alkali metal ions as well as one or more protons. Well known examples are sodium bicarbonate (NaHCO3), sodium hydrosulfide (NaHS), sodium bisulfate (NaHSO4), monosodium phosphate (NaH2PO4), and disodium phosphate (Na2HPO4). Often acid salts are used as buffers.[1]

For example, the acid salt sodium bisulfate is the main species formed upon the half neutralization of sulfuric acid with sodium hydroxide:

H2SO4 + NaOH → NaHSO4 + H2O
Acid salts compounds can act either as an acid or a base: addition of a suitably strong acid will restore protons, and addition of a suitably strong base will remove protons. The pH of a solution of an acid salt will depend on the relevant equilibrium constants and the amounts of any additional base or acid. A comparison between the Kb and Ka will indicate this: if Kb > Ka, the solution will be basic, whereas if Kb < Ka, the solution will be acidic.

Use in food
Main article: baking powder
Some acid salts are used in baking. They are found in baking powders and are typically divided into low-temperature (or single-acting) and high-temperature (or double-acting) acid salts. Common low-temperature acid salts react at room temperature to produce a leavening effect. They include cream of tartar, calcium phosphate, and citrates. High-temperature acid salts produce a leavening effect during baking and are usually aluminium salts such as calcium aluminium phosphate. Some acid salts may also be found in non-dairy coffee creamers.

1.6.1 History of mole

The first table of atomic weights was published by John Dalton (1766–1844) in 1805, based on a system in which the atomic weight of hydrogen was defined as 1. These atomic weights were based on the stoichiometric proportions of chemical reactions and compounds, a fact which greatly aided their acceptance: it was not necessary for a chemist to subscribe to atomic theory (an unproven hypothesis at the time) to make practical use of the tables. This would lead to some confusion between atomic weights (promoted by proponents of atomic theory) and equivalent weights (promoted by its opponents and which sometimes differed from atomic weights by an integer factor), which would last throughout much of the nineteenth century.

Jöns Jacob Berzelius (1779–1848) was instrumental in the determination of atomic weights to ever increasing accuracy. He was also the first chemist to use oxygen as the standard to which other weights were referred. Oxygen is a useful standard, as, unlike hydrogen, it forms compounds with most other elements, especially metals. However he chose to fix the atomic weight of oxygen as 100, an innovation which did not catch on.

Charles Frédéric Gerhardt (1816–56), Henri Victor Regnault (1810–78) and Stanislao Cannizzaro (1826–1910) expanded on Berzelius' work, resolving many of the problems of unknown stoichiometry of compounds, and the use of atomic weights attracted a large consensus by the time of the Karlsruhe Congress (1860). The convention had reverted to defining the atomic weight of hydrogen as 1, although at the level of precision of measurements at that time—relative uncertainties of around 1%—this was numerically equivalent to the later standard of oxygen = 16. However the chemical convenience of having oxygen as the primary atomic weight standard became ever more evident with advances in analytical chemistry and the need for ever more accurate atomic weight determinations.

Scale basis Scale basis
relative to 12C = 12 Relative deviation
from the 12C = 12 scale
Atomic weight of hydrogen = 1 1.007 94(7) −0.788%
Atomic weight of oxygen = 16 15.9994(3) +37.5 ppm
Relative atomic mass of 16O = 16 15.994 914 6221(15) +318 ppm

Other units called "mole"
Chemical engineers use the concept extensively, but the unit is rather small for industrial use. For convenience in avoiding conversions, some American engineers adopted the pound-mole (noted lb-mol or lbmol), which is defined as the number of entities in 12 lb of 12C. One lb-mol is equal to 453.59237 mol.[11] In the metric system, chemical engineers once used the kilogram-mole (noted kg-mol), which is defined as the number of entities in 12 kg of 12C, and often referred to the mole as the gram-mole (noted g-mol), when dealing with laboratory data.[11] However modern chemical engineering practice is to use the kilomole (kmol), which is identical to the kilogram-mole, but whose name and symbol adopt the SI convention for standard multiples of metric units.

Proposed future definition
Kilogram
As with other SI base units, there have been proposals to redefine the kilogram in such a way as to define some currently measured physical constants to fixed values. One proposed definition of the kilogram is:

The kilogram is the mass of exactly (6.0221415×1023⁄0.012) unbound carbon-12 atoms at rest and in their ground state.
This would have the effect of defining the Avogadro constant to be precisely 6.0221415×1023 elementary entities per mole.

Holiday
October the 23rd (10/23), "Mole Day", is an informal holiday in honor of the unit among chemists in North America. The date is derived from the Avogadro constant, which is approximately 6.02×1023.

1.6 Mole

Main article: Mole (unit)
A mole is the amount of a substance that contains as many elementary entities (atoms, molecules or ions) as there are atoms in 0.012 kilogram (or 12 grams) of carbon-12, where the carbon-12 atoms are unbound, at rest and in their ground state.[38] This number is known as the Avogadro constant, and is determined empirically. The currently accepted value is 6.02214179(30) × 1023 mol−1 (2007 CODATA). The best way to understand the meaning of the term "mole" is to compare it to terms such as dozen. Just as one dozen is equal to 12, one mole is equal to 6.02214179(30) × 1023. The term is used because it is much easier to say, for example, 1 mole of carbon atoms, than it is to say 6.02214179(30) × 1023 carbon atoms. Likewise, we can describe the number of entities as a multiple or fraction of 1 mole, e.g. 2 mole or 0.5 moles. Mole is an absolute number (having no units) and can describe any type of elementary object, although the mole's use is usually limited to measurement of subatomic, atomic, and molecular structures.
The number of moles of a substance in one liter of a solution is known as its molarity. Molarity is the common unit used to express the concentration of a solution in physical chemistry

The mole (symbol mol) is the SI base unit[1] of amount of substance, one of a few units used to measure this physical quantity. The name "mole" is an 1897 translation[2][3] of the German Mol, coined by the chemist Wilhelm Ostwald in 1893,[4] although the related concept of equivalent mass had been in use at least a century earlier. The name is assumed to be derived from the German word Molekül (molecule).

The mole is commonly used in titration, a laboratory method in chemistry to determine the concentration of some substance in a solution. In this context, millimoles per litre (mmol/L), micromoles/litre (µmol/L), or nanomoles/L (nmol/L) are often used.

For pure substances (that is, not being an admixture of different substances) the mole is defined as the amount of substance that contains as many "elementary entities" (e.g. atoms, molecules, ions, electrons) as there are atoms in 12 gram of carbon-12 (12C). Thus, by definition, one mole of pure 12C has a mass of exactly 12 g. The number of atoms or molecules contained in one mole of a pure substance is known as the Avogadro constant (or Avogadro's number). By convention it has dimension mol−1, and its experimentally determined value is approximately 6.022142×1023 mol−1. So a mole of any pure substance has mass in grams exactly equal to that substance's molecular or atomic mass; e.g., 1 mol of calcium-40 is equal to 40 g approximately (because Ca-40 has a mass of 39.9625906 amu on the C-12 scale). In other words, the numerical value of a substance's molecular or atomic mass in atomic mass units is the same as that of its molar mass (the mass of one mole of that substance) in grams. (Although the SI base unit for mass is the kilogram, for practical and historical reasons grams are commonly used in this context, especially in chemistry.)

The most common method of determining the amount, expressed in moles, of pure substance the value of whose molar mass is known, is to measure its mass in grams and then to divide by its molar mass (expressed in g/mol). Molar masses may be easily calculated from tabulated values of atomic weights and the molar mass constant (which has a convenient defined value of 1 g/mol). Other methods include the use of the molar volume or the measurement of electric charge.

The current definition of the mole was approved during the 1960s.[1][8] Earlier definitions had been based on the atomic mass of hydrogen (about one gram of hydrogen-1 gas, excluding its heavy isotopes), the atomic weight of oxygen, and the relative atomic mass of oxygen-16; the four different definitions were equivalent to within 1%.

The names gram-atom (abbreviated gat.) and gram-molecule have also been used in the same sense as "mole"[1]. However, modern conventions define the gram-atom and the mole differently. While the elementary entity defining a mole will vary depending on the substance, the elementary entity for the gram-atom is always the atom. For example, 1 mole of He is equivalent to 1 gram-atom of He, but 1 mole of MgB2 is equivalent to 3 gram-atoms of MgB2.

1.5 Molecule

Main article: Molecule
A molecule is the smallest indivisible portion, besides an atom, of a pure chemical substance that has its unique set of chemical properties, that is, its potential to undergo a certain set of chemical reactions with other substances. Molecules can exist as electrically neutral units unlike ions. Molecules are typically a set of atoms bound together by covalent bonds, such that the structure is electrically neutral and all valence electrons are paired with other electrons either in bonds or in lone pairs.
One of the main characteristic of a molecule is its geometry often called its structure. While the structure of diatomic, triatomic or tetra atomic molecules may be trivial, (linear, angular pyramidal etc.) the structure of polyatomic molecules, that are constituted of more than six atoms (of several elements) can be crucial for its chemical nature.

A molecule is defined as an electrically neutral group of at least two atoms in a definite arrangement held together by very strong (covalent) chemical bonds. Molecules are distinguished from polyatomic ions in this strict sense. In organic chemistry and biochemistry, the term molecule is used less strictly and also is applied to charged organic molecules and biomolecules.
In the kinetic theory of gases, the term molecule is often used for any gaseous particle regardless of its composition. According to this definition noble gas atoms are considered molecules despite the fact that they are composed of a single non-bonded atom.

A molecule may consist of atoms of a single chemical element, as with oxygen (O2), or of different elements, as with water (H2O). Atoms and complexes connected by non-covalent bonds such as hydrogen bonds or ionic bonds are generally not considered single molecules.

Molecules as components of matter are common in organic substances (and therefore biochemistry). They also make up most of the oceans and atmosphere. A large number of familiar solid substances, however, including most of the minerals that make up the crust, mantle, and core of the Earth itself, contain many chemical bonds, but are not made of identifiable molecules. No typical molecule can be defined for ionic crystals (salts) and covalent crystals (network solids), although these are often composed of repeating unit cells that extend either in a plane (such as in graphene) or three-dimensionally (such as in diamond or sodium chloride). The theme of repeated unit-cellular-structure also holds for most condensed phases with metallic bonding. In glasses (solids that exist in a vitreous disordered state), atoms may also be held together by chemical bonds without any definable molecule, but also without any of the regularity of repeating units that characterises crystals.

History and etymology
Main article: History of the molecule
According to Merriam-Webster and the Online Etymology Dictionary, the word "molecule" derives from the Latin "moles" or small unit of mass.

Molecule (1794) – "extremely minute particle," from Fr. molécule (1678), from Mod.L. molecula, dim. of L. moles "mass, barrier". A vague meaning at first; the vogue for the word (used until late 18th century only in Latin form) can be traced to the philosophy of Descartes.
Although the existence of molecules has been accepted by many chemists since the early 19th century as a result of Dalton's laws of Definite and Multiple Proportions (1803–1808) and Avogadro's law (1811), there was some resistance among positivists and physicists such as Mach, Boltzmann, Maxwell, and Gibbs, who saw molecules merely as convenient mathematical constructs. The work of Perrin on Brownian motion (1911) is considered to be the final proof of the existence of molecules.

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.This definition often breaks down since many substances in ordinary experience, such as rocks, salts, and metals, are composed of large networks of chemically bonded atoms or ions, but are not made of discrete molecules.

1.4 Substance

Main article: Chemical substance

Substance may refer to:
Chemical substance, in chemistry, any material with a definite chemical composition
Physical substance, any quantity of material with any density, which can be measured or perceived without changing its identity.
Substance theory, in philosophy, that which "stands under" or supports the existence of properties; a particular, concrete object; or what exists only by itself (causa sui)
Substantial (rapper)
Substance (Joy Division album)
Substance (New Order album)
"Substance", a song by Haste the Day from That They May Know You and Burning Bridges
Metal Gear Solid 2: Substance, a tactical espionage video game

A chemical substance is a kind of matter with a definite composition and set of properties. Strictly speaking, a mixture of compounds, elements or compounds and elements is not a chemical substance, but it may be called a chemical. Most of the substances we encounter in our daily life are some kind of mixture; for example: air, alloys, biomass, etc.
Nomenclature of substances is a critical part of the language of chemistry. Generally it refers to a system for naming chemical compounds. Earlier in the history of chemistry substances were given name by their discoverer, which often led to some confusion and difficulty. However, today the IUPAC system of chemical nomenclature allows chemists to specify by name specific compounds amongst the vast variety of possible chemicals. The standard nomenclature of chemical substances is set by the International Union of Pure and Applied Chemistry (IUPAC). There are well-defined systems in place for naming chemical species. Organic compounds are named according to the organic nomenclature system. Inorganic compounds are named according to the inorganic nomenclature system.In addition the Chemical Abstracts Service has devised a method to index chemical substance. In this scheme each chemical substance is identifiable by a number known as CAS registry number.

1.3 Compound

Main article: Chemical compound
Wider definitions:
There are a few exceptions to the definition above. Certain crystalline compounds are called "non-stoichiometric" because they vary in composition due to either the presence of foreign elements trapped within the crystal structure or a deficit or excess of the constituent elements. Some compounds regarded as chemically identical may have varying amounts of heavy or light isotopes of the constituent elements, which will make the ratio of elements by mass vary slightly. A compound therefore may not be completely homogeneous, but for most chemical purposes it can be regarded as such.

A compound is a substance with a particular ratio of atoms of particular chemical elements which determines its composition, and a particular organization which determines chemical properties. For example, water is a compound containing hydrogen and oxygen in the ratio of two to one, with the oxygen atom between the two hydrogen atoms, and an angle of 104.5° between them. Compounds are formed and interconverted by chemical reactions.

A chemical compound is a pure chemical substance consisting of two or more different chemical elements that can be separated into simpler substances by chemical reactions. Chemical compounds have a unique and defined chemical structure; they consist of a fixed ratio of atoms that are held together in a defined spatial arrangement by chemical bonds. Chemical compounds can be molecular compounds held together by covalent bonds, salts held together by ionic bonds, intermetallic compounds held together by metallic bonds, or complexes held together by coordinate covalent bonds. Pure chemical elements are not considered chemical compounds, even if they consist of molecules which contain only multiple atoms of a single element (such as H2, S8, etc.) .

Elements form compounds to become more stable. They become stable when they have the maximum number of possible electrons in their outermost energy level, which is normally two or eight valence electrons. This is the reason that noble gases do not frequently react: they already possess eight valence electrons (the exception being helium, which requires only two valence electrons to achieve stability).

1.2 Element

Main article: Chemical element
The concept of chemical element is related to that of chemical substance. A chemical element is characterized by a particular number of protons in the nuclei of its atoms. This number is known as the atomic number of the element. For example, all atoms with 6 protons in their nuclei are atoms of the chemical element carbon, and all atoms with 92 protons in their nuclei are atoms of the element uranium. 94 different chemical elements or types of atoms based on the number of protons exist naturally. A further 18 have been recognised by IUPAC as existing artificially only. Although all the nuclei of all atoms belonging to one element will have the same number of protons, they may not necessarily have the same number of neutrons, such atoms are termed isotopes. In fact several isotopes of an element may exist.

The most convenient presentation of the chemical elements is in the periodic table of the chemical elements, which groups elements by atomic number. Due to its ingenious arrangement, groups, or columns, and periods, or rows, of elements in the table either share several chemical properties, or follow a certain trend in characteristics such as atomic radius, electronegativity, etc. Lists of the elements by name, by symbol, and by atomic number are also available.

A chemical element is a pure chemical substance consisting of one type of atom distinguished by its atomic number, which is the number of protons in its nucleus.[1] Common examples of elements are iron, copper, silver, gold, hydrogen, carbon, nitrogen, and oxygen. In total, 118 elements have been observed as of March 2010, of which 94 occur naturally on Earth. 80 elements have stable isotopes, namely all elements with atomic numbers 1 to 82, except elements 43 and 61 (technetium and promethium). Elements with atomic numbers 83 or higher (bismuth and above) are inherently unstable, and undergo radioactive decay. The elements from atomic number 83 to 94 have no stable nuclei, but are nevertheless found in nature, either surviving as remnants of the primordial stellar nucleosynthesis that produced the elements in the solar system, or else produced as short-lived daughter-isotopes through the natural decay of uranium and thorium.[2]

All chemical matter consists of these elements. New elements of higher atomic number are discovered from time to time, as products of artificial nuclear reactions.
History
Mendeleev's 1869 periodic tableAncient philosophy posited a set of classical elements to explain patterns in nature. Elements originally referred to earth, water, air and fire rather than the chemical elements of modern science.
The term 'elements' (stoicheia) was first used by the Greek philosopher Plato in about 360 BCE, in his dialogue Timaeus, which includes a discussion of the composition of inorganic and organic bodies and is a speculative treatise on chemistry. Plato believed the elements introduced a century earlier by Empedocles were composed of small polyhedral forms: tetrahedron (fire), octahedron (air), icosahedron (water), and cube (earth).[3][4]
Aristotle, c. 350 BCE, also used the term stoicheia and added a fifth element called aether, which formed the heavens. Aristotle defined an element as:
Element – one of those bodies into which other bodies can decompose, and that itself is not capable of being divided into other.[5]
Building on the theory, Arab/Persian chemist and alchemist, Jābir ibn Hayyān (Geber c. 790), postulated that metals were formed out of two elements: sulfur, ‘the stone that burns’, which characterized the principle of combustibility, and mercury, which contained the idealized principle of metallic properties.[6] Shortly thereafter, this evolved into the Arabic concept of the three principles: sulfur giving flammability or combustion, mercury giving volatility and stability, and in the 10th century, Persian physician and alchemist Muhammad ibn Zakarīya Rāzi (Rhazes) hints at salt giving solidity.

In 1524, Swiss chemist Paracelsus adopted Aristotle’s four element theory, but reasoned that they appeared in bodies as three principles. Paracelsus saw these principles as fundamental, and justified them by recourse to the description of how wood burns in fire. Mercury included the cohesive principle, so that when it left in smoke the wood fell apart. Smoke represented the volatility (the mercury principle), the heat-giving flames represented flammability (sulfur), and the remnant ash represented solidity (salt).[6]

In 1669, German physician and chemist Johann Becher published his Physica Subterranea. In modification on the ideas of Paracelsus, he argued that the constituents of bodies are air, water, and three types of earth: terra fluida, the mercurial element, which contributes fluidity and volatility; terra lapida, the solidifying element, which produces fusibility or the binding quality; and terra pinguis, the fatty element, which gives material substance its oily and combustible qualities.[7] These three earths correspond with Geber’s three principles. A piece of wood, for example, according to Becher, is composed of ash and terra pinguis; when the wood is burnt, the terra pinguis is released, leaving the ash. In other words, in combustion the fatty earth burns away.

In 1661, Robert Boyle showed that there were more than just four classical elements as the ancients had assumed.[8] The first modern list of chemical elements was given in Antoine Lavoisier's 1789 Elements of Chemistry, which contained thirty-three elements, including light and caloric.[9] By 1818, Jöns Jakob Berzelius had determined atomic weights for forty-five of the forty-nine accepted elements. Dmitri Mendeleev had sixty-six elements in his periodic table of 1869.

From Boyle until the early 20th century, an element was defined as a pure substance that cannot be decomposed into any simpler substance.[8] Put another way, a chemical element cannot be transformed into other chemical elements by chemical processes. In 1913, Henry Moseley discovered that the physical basis of the atomic number of the atom was its nuclear charge, which eventually led to the current definition. The current definition also avoids some ambiguities due to isotopes and allotropes.

By 1919, there were seventy-two known elements.[10] In 1955, element 101 was discovered and named mendelevium in honor of Mendeleev, the first to arrange the elements in a periodic manner. In October 2006, the synthesis of element 118 was reported; the synthesis of element 117 was reported in April 2010.[11]

1.1.11 Origin and current state

Origin Of The Atom Theory:
The theory of atoms, and the name atom originated from a Greek Philosopher named Democritus who lived around 400 BC. He proposed that all life and matter is made up of tiny particles which he called atoms. The word atom comes from the Greek word atomos meaning invisible.



Atoms form about 4% of the total energy density of the observable universe, with an average density of about 0.25 atoms/m3. Within a galaxy such as the Milky Way, atoms have a much higher concentration, with the density of matter in the interstellar medium (ISM) ranging from 105 to 109 atoms/m3. The Sun is believed to be inside the Local Bubble, a region of highly ionized gas, so the density in the solar neighborhood is only about 103 atoms/m3. Stars form from dense clouds in the ISM, and the evolutionary processes of stars result in the steady enrichment of the ISM with elements more massive than hydrogen and helium. Up to 95% of the Milky Way's atoms are concentrated inside stars and the total mass of atoms forms about 10% of the mass of the galaxy. (The remainder of the mass is an unknown dark matter.)

1.1.10 Identification

Scanning tunneling microscope image showing the individual atoms making up this gold surface. Reconstruction causes the surface atoms to deviate from the bulk crystal structure and arrange in columns several atoms wide with pits between them.
The scanning tunneling microscope is a device for viewing surfaces at the atomic level. It uses the quantum tunneling phenomenon, which allows particles to pass through a barrier that would normally be insurmountable. Electrons tunnel through the vacuum between two planar metal electrodes, on each of which is an adsorbed atom, providing a tunneling-current density that can be measured. Scanning one atom (taken as the tip) as it moves past the other (the sample) permits plotting of tip displacement versus lateral separation for a constant current. The calculation shows the extent to which scanning-tunneling-microscope images of an individual atom are visible. It confirms that for low bias, the microscope images the space-averaged dimensions of the electron orbitals across closely packed energy levels—the Fermi level local density of states.
An atom can be ionized by removing one of its electrons. The electric charge causes the trajectory of an atom to bend when it passes through a magnetic field. The radius by which the trajectory of a moving ion is turned by the magnetic field is determined by the mass of the atom. The mass spectrometer uses this principle to measure the mass-to-charge ratio of ions. If a sample contains multiple isotopes, the mass spectrometer can determine the proportion of each isotope in the sample by measuring the intensity of the different beams of ions. Techniques to vaporize atoms include inductively coupled plasma atomic emission spectroscopy and inductively coupled plasma mass spectrometry, both of which use a plasma to vaporize samples for analysis.[111]
A more area-selective method is electron energy loss spectroscopy, which measures the energy loss of an electron beam within a transmission electron microscope when it interacts with a portion of a sample. The atom-probe tomograph has sub-nanometer resolution in 3-D and can chemically identify individual atoms using time-of-flight mass spectrometry.[112]
Spectra of excited states can be used to analyze the atomic composition of distant stars. Specific light wavelengths contained in the observed light from stars can be separated out and related to the quantized transitions in free gas atoms. These colors can be replicated using a gas-discharge lamp containing the same element.[113] Helium was discovered in this way in the spectrum of the Sun 23 years before it was found on Earth.[114]

1.1.9 States

Main articles: State of matter and Phase (matter)









Snapshots illustrating the formation of a Bose–Einstein condensate.
Quantities of atoms are found in different states of matter that depend on the physical conditions, such as temperature and pressure. By varying the conditions, materials can transition between solids, liquids, gases and plasmas. Within a state, a material can also exist in different phases. An example of this is solid carbon, which can exist as graphite or diamond.
At temperatures close to absolute zero, atoms can form a Bose–Einstein condensate, at which point quantum mechanical effects, which are normally only observed at the atomic scale, become apparent on a macroscopic scale. This super-cooled collection of atoms then behaves as a single super atom, which may allow fundamental checks of quantum mechanical behavior.

1.1.8 Energy levels

Main articles: Energy level and Atomic spectral line
When an electron is bound to an atom, it has a potential energy that is inversely proportional to its distance from the nucleus. This is measured by the amount of energy needed to unbind the electron from the atom, and is usually given in units of electronvolts (eV). In the quantum mechanical model, a bound electron can only occupy a set of states centered on the nucleus, and each state corresponds to a specific energy level. The lowest energy state of a bound electron is called the ground state, while an electron at a higher energy level is in an excited state.
For an electron to transition between two different states, it must absorb or emit a photon at an energy matching the difference in the potential energy of those levels. The energy of an emitted photon is proportional to its frequency, so these specific energy levels appear as distinct bands in the electromagnetic spectrum.Each element has a characteristic spectrum that can depend on the nuclear charge, subshells filled by electrons, the electromagnetic interactions between the electrons and other factors.
An example of absorption lines in a spectrum.
When a continuous spectrum of energy is passed through a gas or plasma, some of the photons are absorbed by atoms, causing electrons to change their energy level. Those excited electrons that remain bound to their atom spontaneously emit this energy as a photon, traveling in a random direction, and so drop back to lower energy levels. Thus the atoms behave like a filter that forms a series of dark absorption bands in the energy output. (An observer viewing the atoms from a view that doesn't include the continuous spectrum in the background, instead sees a series of emission lines from the photons emitted by the atoms.) Spectroscopic measurements of the strength and width of spectral lines allow the composition and physical properties of a substance to be determined.
Close examination of the spectral lines reveals that some display a fine structure splitting. This occurs because of spin-orbit coupling, which is an interaction between the spin and motion of the outermost electron. When an atom is in an external magnetic field, spectral lines become split into three or more components; a phenomenon called the Zeeman effect. This is caused by the interaction of the magnetic field with the magnetic moment of the atom and its electrons. Some atoms can have multiple electron configurations with the same energy level, which thus appear as a single spectral line. The interaction of the magnetic field with the atom shifts these electron configurations to slightly different energy levels, resulting in multiple spectral lines.[97] The presence of an external electric field can cause a comparable splitting and shifting of spectral lines by modifying the electron energy levels, a phenomenon called the Stark effect.[98]
If a bound electron is in an excited state, an interacting photon with the proper energy can cause stimulated emission of a photon with a matching energy level. For this to occur, the electron must drop to a lower energy state that has an energy difference matching the energy of the interacting photon. The emitted photon and the interacting photon then move off in parallel and with matching phases. That is, the wave patterns of the two photons are synchronized. This physical property is used to make lasers, which can emit a coherent beam of light energy in a narrow frequency band.

1.1.7 Magnetic moment

Main articles: Electron magnetic dipole moment and Nuclear magnetic moment
Elementary particles possess an intrinsic quantum mechanical property known as spin. This is analogous to the angular momentum of an object that is spinning around its center of mass, although strictly speaking these particles are believed to be point-like and cannot be said to be rotating. Spin is measured in units of the reduced Planck constant (ħ), with electrons, protons and neutrons all having spin ½ ħ, or "spin-½". In an atom, electrons in motion around the nucleus possess orbital angular momentum in addition to their spin, while the nucleus itself possesses angular momentum due to its nuclear spin.
The magnetic field produced by an atom—its magnetic moment—is determined by these various forms of angular momentum, just as a rotating charged object classically produces a magnetic field. However, the most dominant contribution comes from spin. Due to the nature of electrons to obey the Pauli exclusion principle, in which no two electrons may be found in the same quantum state, bound electrons pair up with each other, with one member of each pair in a spin up state and the other in the opposite, spin down state. Thus these spins cancel each other out, reducing the total magnetic dipole moment to zero in some atoms with even number of electrons.
In ferromagnetic elements such as iron, an odd number of electrons leads to an unpaired electron and a net overall magnetic moment. The orbitals of neighboring atoms overlap and a lower energy state is achieved when the spins of unpaired electrons are aligned with each other, a process known as an exchange interaction. When the magnetic moments of ferromagnetic atoms are lined up, the material can produce a measurable macroscopic field. Paramagnetic materials have atoms with magnetic moments that line up in random directions when no magnetic field is present, but the magnetic moments of the individual atoms line up in the presence of a field.
The nucleus of an atom can also have a net spin. Normally these nuclei are aligned in random directions because of thermal equilibrium. However, for certain elements (such as xenon-129) it is possible to polarize a significant proportion of the nuclear spin states so that they are aligned in the same direction—a condition called hyperpolarization. This has important applications in magnetic resonance imaging.

Magnetic Moment:
The magnetic moment of a magnet is a measure of its tendency to align with a magnetic field. Both the magnetic moment and magnetic field may be considered to be vectors having a magnitude and direction. The direction of the magnetic moment points from the south to north pole of a magnet. The magnetic field produced by a magnet is proportional to its magnetic moment as well. For example, a loop of electric current, a bar magnet, an electron, a molecule, and a planet all have magnetic moments. More precisely, the term magnetic moment normally refers to a system's magnetic dipole moment, which produces the first term in the multipole expansion of a general magnetic field. The dipole component of an object's magnetic field is symmetric about the direction of its magnetic dipole moment, and decreases as the inverse cube of the distance from the object.

1.1.6 Radioactive decay

Main article: Radioactive decay













This diagram shows the half-life (T½) of various isotopes with Z protons and N neutrons.
Every element has one or more isotopes that have unstable nuclei that are subject to radioactive decay, causing the nucleus to emit particles or electromagnetic radiation. Radioactivity can occur when the radius of a nucleus is large compared with the radius of the strong force, which only acts over distances on the order of 1 fm.
The most common forms of radioactive decay are:
• Alpha decay is caused when the nucleus emits an alpha particle, which is a helium nucleus consisting of two protons and two neutrons. The result of the emission is a new element with a lower atomic number.
• Beta decay is regulated by the weak force, and results from a transformation of a neutron into a proton, or a proton into a neutron. The first is accompanied by the emission of an electron and an antineutrino, while the second causes the emission of a positron and a neutrino. The electron or positron emissions are called beta particles. Beta decay either increases or decreases the atomic number of the nucleus by one.
• Gamma decay results from a change in the energy level of the nucleus to a lower state, resulting in the emission of electromagnetic radiation. This can occur following the emission of an alpha or a beta particle from radioactive decay.
Other more rare types of radioactive decay include ejection of neutrons or protons or clusters of nucleons from a nucleus, or more than one beta particle, or result (through internal conversion) in production of high-speed electrons that are not beta rays, and high-energy photons that are not gamma rays.
Each radioactive isotope has a characteristic decay time period—the half-life—that is determined by the amount of time needed for half of a sample to decay. This is an exponential decay process that steadily decreases the proportion of the remaining isotope by 50% every half life. Hence after two half-lives have passed only 25% of the isotope is present, and so forth.

1.1.5 Mass + Shape & Size

Mass
Main article: Atomic mass
Because the large majority of an atom's mass comes from the protons and neutrons, the total number of these particles in an atom is called the mass number. The mass of an atom at rest is often expressed using the unified atomic mass unit (u), which is also called a Dalton (Da). This unit is defined as a twelfth of the mass of a free neutral atom of carbon-12, which is approximately 1.66 × 10−27 kg. Hydrogen-1, the lightest isotope of hydrogen and the atom with the lowest mass, has an atomic weight of 1.007825 u. An atom has a mass approximately equal to the mass number times the atomic mass unit. The heaviest stable atom is lead-208,[68] with a mass of 207.9766521 u.
As even the most massive atoms are far too light to work with directly, chemists instead use the unit of moles. The mole is defined such that one mole of any element always has the same number of atoms (about 6.022 × 1023). This number was chosen so that if an element has an atomic mass of 1 u, a mole of atoms of that element has a mass close to 0.001 kg, or 1 gram. Because of the definition of the unified atomic mass unit, carbon has an atomic mass of exactly 12 u, and so a mole of carbon atoms weighs exactly 0.012 kg.

Shape and size
Main article: Atomic radius
Atoms lack a well-defined outer boundary, so their dimensions are usually described in terms of an atomic radius. This is a measure of the distance out to which the electron cloud extends from the nucleus. However, this assumes the atom to exhibit a spherical shape, which is only obeyed for atoms in vacuum or free space. Atomic radii may be derived from the distances between two nuclei when the two atoms are joined in a chemical bond. The radius varies with the location of an atom on the atomic chart, the type of chemical bond, the number of neighboring atoms (coordination number) and a quantum mechanical property known as spin. On the periodic table of the elements, atom size tends to increase when moving down columns, but decrease when moving across rows (left to right). Consequently, the smallest atom is helium with a radius of 32 pm, while one of the largest is caesium at 225 pm.
When subjected to external fields, like an electrical field, the shape of an atom may deviate from that a sphere. The deformation depends on the field magnitude and the orbital type of outer shell electrons, as shown by group-theoretical considerations. Aspherical deviations might be elicited for instance in crystals, where large crystal-electrical fields may occur at low-symmetry lattice sites.Significant ellipsoidal deformations have recently been shown to occur for sulfur ions in pyrite-type compounds.
Atomic dimensions are thousands of times smaller than the wavelengths of light (400–700 nm) so they can not be viewed using an optical microscope. However, individual atoms can be observed using a scanning tunneling microscope. To visualize the minuteness of the atom, consider that a typical human hair is about 1 million carbon atoms in width. A single drop of water contains about 2 sextillion (2 × 1021) atoms of oxygen, and twice the number of hydrogen atoms. A single carat diamond with a mass of 2 × 10-4 kg contains about 10 sextillion (1022) atoms of carbon.[note 2] If an apple were magnified to the size of the Earth, then the atoms in the apple would be approximately the size of the original apple.

1.1.4 Nuclear properties:

Main articles:
Isotope, Stable isotope, List of elements by nuclear stability, and List of elements by stability of isotopes ;
By definition, any two atoms with an identical number of protons in their nuclei belong to the same chemical element. Atoms with equal numbers of protons but a different number of neutrons are different isotopes of the same element. For example, all hydrogen atoms admit exactly one proton, but isotopes exist with no neutrons hydrogen-1, one neutron (deuterium), two neutrons (tritium) and more than two neutrons. The hydrogen-1 is by far the most common form, and is sometimes called protium. The known elements form a set of atomic numbers from hydrogen with a single proton up to the 118-proton element ununoctium. All known isotopes of elements with atomic numbers greater than 82 are radioactive.
About 339 nuclides occur naturally on Earth, of which 256 (about 76%) have not been observed to decay, and are referred to as "stable isotopes". For 80 of the chemical elements, there is at least one stable isotope. Elements 43, 61, and all elements numbered 83 or higher have no stable isotopes. As a rule, there is, for each element, only a handful of stable isotopes, the average being 3.1 stable isotopes per element among those that have stable isotopes. Twenty-seven elements have only a single stable isotope, while the largest number of stable isotopes observed for any element is ten, for the element tin.
Stability of isotopes is affected by the ratio of protons to neutrons, and also by the presence of certain "magic numbers" of neutrons or protons that represent closed and filled quantum shells. These quantum shells correspond to a set of energy levels within the shell model of the nucleus; filled shells, such as the filled shell of 50 protons for tin, confers unusual stability on the nuclide. Of the 256 known stable nuclides, only four have both an odd number of protons and odd number of neutrons: hydrogen-2 (deuterium), lithium-6, boron-10 and nitrogen-14. Also, only four naturally occurring, radioactive odd-odd nuclides have a half-life over a billion years: potassium-40, vanadium-50, lanthanum-138 and tantalum-180m. Most odd-odd nuclei are highly unstable with respect to beta decay, because the decay products are even-even, and are therefore more strongly bound, due to nuclear pairing effects.

1.1.3 Properties of an atom

1. Nuclear properties
2. Mass
3. Shape and size
4. Radioactive decay
5. Magnetic moment
6. Energy levels
7. Valence and bonding behavior
8. States

1.1.2 Concepts

Subatomic particles :
Though the word atom originally denoted a particle that cannot be cut into smaller particles, in modern scientific usage the atom is composed of various subatomic particles. The constituent particles of an atom are the electron, the proton and the neutron. However, the hydrogen-1 atom has no neutrons and a positive hydrogen ion has no electrons.
The electron is by far the least massive of these particles at 9.11 × 10−31 kg, with a negative electrical charge and a size that is too small to be measured using available techniques. Protons have a positive charge and a mass 1,836 times that of the electron, at 1.6726 × 10−27 kg, although this can be reduced by changes to the energy binding the proton into an atom. Neutrons have no electrical charge and have a free mass of 1,839 times the mass of electrons,[47] or 1.6929 × 10−27 kg. Neutrons and protons have comparable dimensions—on the order of 2.5 × 10−15 m—although the 'surface' of these particles is not sharply defined.[48]
In the Standard Model of physics, both protons and neutrons are composed of elementary particles called quarks. The quark belongs to the fermion group of particles, and is one of the two basic constituents of matter—the other being the lepton, of which the electron is an example. There are six types of quarks, each having a fractional electric charge of either +2/3 or −1/3. Protons are composed of two up quarks and one down quark, while a neutron consists of one up quark and two down quarks. This distinction accounts for the difference in mass and charge between the two particles. The quarks are held together by the strong nuclear force, which is mediated by gluons. The gluon is a member of the family of gauge bosons, which are elementary particles that mediate physical forces.
Nucleus :













All the bound protons and neutrons in an atom make up a tiny atomic nucleus, and are collectively called nucleons. The radius of a nucleus is approximately equal to fm, where A is the total number of nucleons. This is much smaller than the radius of the atom, which is on the order of 105 fm. The nucleons are bound together by a short-ranged attractive potential called the residual strong force. At distances smaller than 2.5 fm this force is much more powerful than the electrostatic force that causes positively charged protons to repel each other.Atoms of the same element have the same number of protons, called the atomic number. Within a single element, the number of neutrons may vary, determining the isotope of that element. The total number of protons and neutrons determine the nuclide. The number of neutrons relative to the protons determines the stability of the nucleus, with the binding energy needed for a nucleon to escape the nucleus, for various isotopes.
certain isotopes undergoing radioactive decay.
The neutron and the proton are different types of fermions. The Pauli exclusion principle is a quantum mechanical effect that prohibits identical fermions, such as multiple protons, from occupying the same quantum physical state at the same time. Thus every proton in the nucleus must occupy a different state, with its own energy level, and the same rule applies to all of the neutrons. This prohibition does not apply to a proton and neutron occupying the same quantum state.
For atoms with low atomic numbers, a nucleus that has a different number of protons than neutrons can potentially drop to a lower energy state through a radioactive decay that causes the number of protons and neutrons to more closely match. As a result, atoms with roughly matching numbers of protons and neutrons are more stable against decay. However, with increasing atomic number, the mutual repulsion of the protons requires an increasing proportion of neutrons to maintain the stability of the nucleus, which modifies this trend. Thus, there are no stable nuclei with equal proton and neutron numbers above atomic number Z = 20 (calcium); and as Z increases toward the heaviest nuclei, the ratio of neutrons per proton required for stability increases to about 1.5.










Illustration of a nuclear fusion process that forms a deuterium nucleus, consisting of a proton and a neutron, from two protons. A positron (e+)—an antimatter electron—is emitted along with an electron neutrino.
The number of protons and neutrons in the atomic nucleus can be modified, although this can require very high energies because of the strong force. Nuclear fusion occurs when multiple atomic particles join to form a heavier nucleus, such as through the energetic collision of two nuclei. For example, at the core of the Sun protons require energies of 3–10 keV to overcome their mutual repulsion—the coulomb barrier—and fuse together into a single nucleus.[55] Nuclear fission is the opposite process, causing a nucleus to split into two smaller nuclei—usually through radioactive decay. The nucleus can also be modified through bombardment by high energy subatomic particles or photons. If this modifies the number of protons in a nucleus, the atom changes to a different chemical element.
If the mass of the nucleus following a fusion reaction is less than the sum of the masses of the separate particles, then the difference between these two values can be emitted as a type of usable energy (such as a gamma ray, or the kinetic energy of a beta particle), as described by Albert Einstein's mass–energy equivalence formula, E = mc2, where m is the mass loss and c is the speed of light. This deficit is part of the binding energy of the new nucleus, and it is the non-recoverable loss of the energy that causes the fused particles to remain together in a state that requires this energy to separate.
The fusion of two nuclei that create larger nuclei with lower atomic numbers than iron and nickel—a total nucleon number of about 60—is usually an exothermic process that releases more energy than is required to bring them together. It is this energy-releasing process that makes nuclear fusion in stars a self-sustaining reaction. For heavier nuclei, the binding energy per nucleon in the nucleus begins to decrease. That means fusion processes producing nuclei that have atomic numbers higher than about 26, and atomic masses higher than about 60, is an endothermic process. These more massive nuclei can not undergo an energy-producing fusion reaction that can sustain the hydrostatic equilibrium of a star.
Electron cloud :







A potential well, showing, according to classical mechanics, the minimum energy V(x) needed to reach each position x. Classically, a particle with energy E is constrained to a range of positions between x1 and x2.
The electrons in an atom are attracted to the protons in the nucleus by the electromagnetic force. This force binds the electrons inside an electrostatic potential well surrounding the smaller nucleus, which means that an external source of energy is needed for the electron to escape. The closer an electron is to the nucleus, the greater the attractive force. Hence electrons bound near the center of the potential well require more energy to escape than those at greater separations.
Electrons, like other particles, have properties of both a particle and a wave. The electron cloud is a region inside the potential well where each electron forms a type of three-dimensional standing wave—a wave form that does not move relative to the nucleus. This behavior is defined by an atomic orbital, a mathematical function that characterises the probability that an electron appears to be at a particular location when its position is measured.Only a discrete (or quantized) set of these orbitals exist around the nucleus, as other possible wave patterns rapidly decay into a more stable form. Orbitals can have one or more ring or node structures, and they differ from each other in size, shape and orientation.
Wave functions of the first five atomic orbitals. The three 2p orbitals each display a single angular node that has an orientation and a minimum at the center.
Each atomic orbital corresponds to a particular energy level of the electron. The electron can change its state to a higher energy level by absorbing a photon with sufficient energy to boost it into the new quantum state. Likewise, through spontaneous emission, an electron in a higher energy state can drop to a lower energy state while radiating the excess energy as a photon. These characteristic energy values, defined by the differences in the energies of the quantum states, are responsible for atomic spectral lines.
The amount of energy needed to remove or add an electron—the electron binding energy—is far less than the binding energy of nucleons. For example, it requires only 13.6 eV to strip a ground-state electron from a hydrogen atom, compared to 2.23 million eV for splitting a deuterium nucleus. Atoms are electrically neutral if they have an equal number of protons and electrons. Atoms that have either a deficit or a surplus of electrons are called ions. Electrons that are farthest from the nucleus may be transferred to other nearby atoms or shared between atoms. By this mechanism, atoms are able to bond into molecules and other types of chemical compounds like ionic and covalent network crystals.

History of an atom

Subcomponents and quantum theory:
The physicist J. J. Thomson, through his work on cathode rays in 1897, discovered the electron, and concluded that they were a component of every atom. Thus he overturned the belief that atoms are the indivisible, ultimate particles of matter.Thomson postulated that the low mass, negatively charged electrons were distributed throughout the atom, possibly rotating in rings, with their charge balanced by the presence of a uniform sea of positive charge. This later became known as the plum pudding model.
In 1909, Hans Geiger and Ernest Marsden, under the direction of physicist Ernest Rutherford, bombarded a sheet of gold foil with alpha rays—by then known to be positively charged helium atoms—and discovered that a small percentage of these particles were deflected through much larger angles than was predicted using Thomson's proposal. Rutherford interpreted the gold foil experiment as suggesting that the positive charge of a heavy gold atom and most of its mass was concentrated in a nucleus at the center of the atom—the Rutherford model.
While experimenting with the products of radioactive decay, in 1913 radiochemist Frederick Soddy discovered that there appeared to be more than one type of atom at each position on the periodic table. The term isotope was coined by Margaret Todd as a suitable name for different atoms that belong to the same element. J.J. Thomson created a technique for separating atom types through his work on ionized gases, which subsequently led to the discovery of stable isotopes.
A Bohr model of the hydrogen atom, showing an electron jumping between fixed orbits and emitting a photon of energy with a specific frequency.
Meanwhile, in 1913, physicist Niels Bohr suggested that the electrons were confined into clearly defined, quantized orbits, and could jump between these, but could not freely spiral inward or outward in intermediate states. An electron must absorb or emit specific amounts of energy to transition between these fixed orbits. When the light from a heated material was passed through a prism, it produced a multi-colored spectrum. The appearance of fixed lines in this spectrum was successfully explained by these orbital transitions.
Chemical bonds between atoms were now explained, by Gilbert Newton Lewis in 1916, as the interactions between their constituent electrons.As the chemical properties of the elements were known to largely repeat themselves according to the periodic law,in 1919 the American chemist Irving Langmuir suggested that this could be explained if the electrons in an atom were connected or clustered in some manner. Groups of electrons were thought to occupy a set of electron shells about the nucleus.The Stern–Gerlach experiment of 1922 provided further evidence of the quantum nature of the atom. When a beam of silver atoms was passed through a specially shaped magnetic field, the beam was split based on the direction of an atom's angular momentum, or spin. As this direction is random, the beam could be expected to spread into a line. Instead, the beam was split into two parts, depending on whether the atomic spin was oriented up or down.
In 1924, Louis de Broglie proposed that all particles behave to an extent like waves. In 1926, Erwin Schrödinger used this idea to develop a mathematical model of the atom that described the electrons as three-dimensional waveforms rather than point particles. A consequence of using waveforms to describe particles is that it is mathematically impossible to obtain precise values for both the position and momentum of a particle at the same time; this became known as the uncertainty principle, formulated by Werner Heisenberg in 1926. In this concept, for a given accuracy in measuring a position one could only obtain a range of probable values for momentum, and vice versa. This model was able to explain observations of atomic behavior that previous models could not, such as certain structural and spectral patterns of atoms larger than hydrogen. Thus, the planetary model of the atom was discarded in favor of one that described atomic orbital zones around the nucleus where a given electron is most likely to be observed.
The development of the mass spectrometer allowed the exact mass of atoms to be measured. The device uses a magnet to bend the trajectory of a beam of ions, and the amount of deflection is determined by the ratio of an atom's mass to its charge. The chemist Francis William Aston used this instrument to show that isotopes had different masses. The atomic mass of these isotopes varied by integer amounts, called the whole number rule. The explanation for these different isotopes awaited the discovery of the neutron, a neutral-charged particle with a mass similar to the proton, by the physicist James Chadwick in 1932. Isotopes were then explained as elements with the same number of protons, but different numbers of neutrons within the nucleus.

Fission, high energy physics and condensed matter:
In 1938, the German chemist Otto Hahn, a student of Rutherford, directed neutrons onto uranium atoms expecting to get transuranium elements. Instead, his chemical experiments showed barium as a product. A year later, Lise Meitner and her nephew Otto Frisch verified that Hahn's result were the first experimental nuclear fission. In 1944, Hahn received the Nobel prize in chemistry. Despite Hahn's efforts, the contributions of Meitner and Frisch were not recognized.
In the 1950s, the development of improved particle accelerators and particle detectors allowed scientists to study the impacts of atoms moving at high energies. Neutrons and protons were found to be hadrons, or composites of smaller particles called quarks. Standard models of nuclear physics were developed that successfully explained the properties of the nucleus in terms of these sub-atomic particles and the forces that govern their interactions.
Fission and fusion are different types of nuclear reactions in which energy is released from the high-powered bonds between particles in the atomic nucleus. The atomic nucleus is most stable when binding energies between particles are strongest. This occurs with iron and nickel. For lighter atomic nuclei, energy can be extracted by combining these nuclei together, a process known as nuclear fusion. For nuclei heavier than those of iron or nickel, energy can be extracted by splitting them apart in a process called nuclear fission.

In primatology, a fission-fusion society is one in which the social group, e.g. bonobo collectives of 100-strong, sleep in one locality together, but forage in small groups going off in different directions during the day. This form of social organization occurs in several other species of primates, though usually less organised and less social than bonobos (e.g. chimpanzees, hamadryas, gelada baboons, spider monkeys, and humans), most carnivores including the spotted hyena, African lion, and cetaceans such as bottlenose dolphins, ungulates such as deer, and fish such as guppies. These societies change frequently in their size and composition, making up a permanent social group called the 'parent group.' Permanent social networks consist of all individual members of a faunal community and often varies to track changes in their environment and based on individual animal dynamics.

In a fission-fusion society, the main parent group can fracture (fission) into smaller stable subgroups or individuals to adapt to environmental or social circumstances. For example, a number of males may break off from the main group in order to hunt or forage for food during the day, but at night they may return to join (fusion) the primary group to share food and partake in other activities.

Overlapping of so-called 'parent groups' territorially is also frequent, resulting in more interaction and mingling of community members, further altering the make-up of the parent group. This results in instances where, say, a female chimpanzee may generally belong to one parent group, but encounters a male who belongs to a neighboring community. If they copulate, the female may stay with the male for several days and come into contact with his parent group, temporarily 'fusing' into the male's community. In some cases, animals may leave one parent group in favor of associating themselves with another, usually for reproductively motivated reasons