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]