In chemistry and manufacturing, electrolysis is a technique that uses a direct electric current to drive an otherwise non-spontaneous chemical reaction. Electrolysis is commercially important as a stage in the separation of elements from occurring sources such as ores using an electrolytic cell; the voltage, needed for electrolysis to occur is called the decomposition potential. The word "electrolysis" was introduced by Michael Faraday in the 19th century, on the suggestion of the Rev. William Whewell, using the Greek words ἤλεκτρον "amber", which since the 17th century was associated with electric phenomena, λύσις meaning "dissolution". Electrolysis, as a tool to study chemical reactions and obtain pure elements, precedes the coinage of the term and formal description by Faraday. 1785 – Martinus van Marum's electrostatic generator was used to reduce tin and antimony from their salts using electrolysis. 1800 – William Nicholson and Anthony Carlisle, decomposed water into hydrogen and oxygen.
1808 – Potassium, barium and magnesium were discovered by Sir Humphry Davy using electrolysis. 1821 – Lithium was discovered by the English chemist William Thomas Brande, who obtained it by electrolysis of lithium oxide. 1833 – Michael Faraday develops his two laws of electrolysis, provides a mathematical explanation of his laws. 1875 – Paul Émile Lecoq de Boisbaudran discovered gallium using electrolysis. 1886 – Fluorine was discovered by Henri Moissan using electrolysis. 1886 – Hall–Héroult process developed for making aluminium 1890 – Castner–Kellner process developed for making sodium hydroxide Electrolysis is the passing of a direct electric current through an ionic substance, either molten or dissolved in a suitable solvent, producing chemical reactions at the electrodes and a decomposition of the materials. The main components required to achieve electrolysis are: An electrolyte: a substance an ion-conducting polymer that contains free ions, which carry electric current in the electrolyte.
If the ions are not mobile, as in most solid salts electrolysis cannot occur. A direct current electrical supply: provides the energy necessary to create or discharge the ions in the electrolyte. Electric current is carried by electrons in the external circuit. Two electrodes: electrical conductors that provide the physical interface between the electrolyte and the electrical circuit that provides the energy. Electrodes of metal and semiconductor material are used. Choice of suitable electrode depends on chemical reactivity between the electrode and electrolyte and manufacturing cost; the key process of electrolysis is the interchange of atoms and ions by the removal or addition of electrons from the external circuit. The desired products of electrolysis are in a different physical state from the electrolyte and can be removed by some physical processes. For example, in the electrolysis of brine to produce hydrogen and chlorine, the products are gaseous; these are collected. 2 NaCl + 2 H2O → 2 NaOH + H2 + Cl2A liquid containing electrolyte is produced by: Solvation or reaction of an ionic compound with a solvent to produce mobile ions An ionic compound is melted by heatingAn electrical potential is applied across a pair of electrodes immersed in the electrolyte.
Each electrode attracts ions. Positively charged ions move towards the electron-providing cathode. Negatively charged ions move towards the electron-extracting anode. In this process electrons are either released. Neutral atoms gain or lose electrons and become charged ions that pass into the electrolyte; the formation of uncharged atoms from ions is called discharging. When an ion gains or loses enough electrons to become uncharged atoms, the newly formed atoms separate from the electrolyte. Positive metal ions like Cu2+deposit onto the cathode in a layer; the terms for this are electroplating and electrorefining. When an ion gains or loses electrons without becoming neutral, its electronic charge is altered in the process. In chemistry, the loss of electrons is called oxidation. Oxidation of ions or neutral molecules occurs at the anode. For example, it is possible to oxidize ferrous ions to ferric ions at the anode: Fe2+ → Fe3+ + e−Reduction of ions or neutral molecules occurs at the cathode, it is possible to reduce ferricyanide ions to ferrocyanide ions at the cathode: Fe3-6 + e− → Fe4-6Neutral molecules can react at either of the electrodes.
For example: p-Benzoquinone can be reduced to hydroquinone at the cathode: + 2 e− + 2 H+ → In the last example, H+ ions take part in the reaction, are provided by an acid in the solution, or by the solvent itself. Electrolysis reactions involving H+ ions are common in acidic solutions. In aqueous alkaline solutions, reactions involving OH− are common. Sometimes the solvents themselves are reduced at the electrodes, it is possible to have electrolysis involving gases. Such as when using a Gas diffusion electrode; the amount of electrical energy that must be added equals the change in Gibbs free energy of the reaction plus the losses in the system. The losses can be arbitrarily close to zero, so the maximum thermodynamic efficiency equals the enthalpy change divided by the free energy change of the reaction. In most cases, the electric input is larger than the enthalpy change of the reaction, so some energy is released in the form of heat. In some cases, for instance, in
Timeline of hydrogen technologies
This is a timeline of the history of hydrogen technology. 1625 – First description of hydrogen by Johann Baptista van Helmont. First to use the word "gas". 1650 – Turquet de Mayerne obtained a gas or "inflammable air" by the action of dilute sulphuric acid on iron. 1662 – Boyle's law 1670 – Robert Boyle produced hydrogen by reacting metals with acid. 1672 – "New Experiments touching the Relation between Flame and Air" by Robert Boyle. 1679 – Denis Papin – safety valve 1700 – Nicolas Lemery showed that the gas produced in the sulfuric acid/iron reaction was explosive in air 1755 – Joseph Black confirmed that different gases exist. / Latent heat 1766 – Henry Cavendish published in "On Factitious Airs" a description of "dephlogisticated air" by reacting zinc metal with hydrochloric acid and isolated a gas 7 to 11 times lighter than air. 1774 – Joseph Priestley isolated and categorized oxygen. 1780 – Felice Fontana discovers the water-gas shift reaction 1783 – Antoine Lavoisier gave hydrogen its name 1783 – Jacques Charles made the first flight with his hydrogen balloon "La Charlière".
1783 – Antoine Lavoisier and Pierre Laplace measured the heat of combustion of hydrogen using an ice calorimeter. 1784 -- Jean-Pierre Blanchard, attempted a dirigible hydrogen balloon. 1784 – The invention of the Lavoisier Meusnier iron-steam process, generating hydrogen by passing water vapor over a bed of red-hot iron at 600 °C. 1785 – Jean-François Pilâtre de Rozier built the hybrid Rozière balloon. 1787 – Charles's law 1789 – Jan Rudolph Deiman and Adriaan Paets van Troostwijk using an electrostatic machine and a Leyden jar for the first electrolysis of water. 1800 – William Nicholson and Anthony Carlisle decomposed water into hydrogen and oxygen by electrolysis with a voltaic pile. 1800 – Johann Wilhelm Ritter duplicated the experiment with a rearranged set of electrodes to collect the two gases separately. 1801 – Humphry Davy discovers the concept of the Fuel Cell. 1806 – François Isaac de Rivaz built the de Rivaz engine, the first internal combustion engine powered by a mixture of hydrogen and oxygen.
1809 – Thomas Forster observed with a theodolite the drift of small free pilot balloons filled with "inflammable gas" 1809 – Gay-Lussac's law 1811 – Amedeo Avogadro – Avogadro's law a gas law 1819 – Edward Daniel Clarke invented the hydrogen gas blowpipe. 1820 – W. Cecil wrote a letter "On the application of hydrogen gas to produce a moving power in machinery" 1823 – Goldsworthy Gurney demonstrated limelight. 1823 – Döbereiner's Lamp a lighter invented by Johann Wolfgang Döbereiner. 1823 – Goldsworthy Gurney devised an oxy-hydrogen blowpipe. 1824 – Michael Faraday invented the rubber balloon. 1826 – Thomas Drummond built the Drummond Light. 1826 – Samuel Brown tested his internal combustion engine by using it to propel a vehicle up Shooter's Hill 1834 – Michael Faraday published Faraday's laws of electrolysis. 1834 – Benoît Paul Émile Clapeyron – Ideal gas law 1836 – John Frederic Daniell invented a primary cell in which hydrogen was eliminated in the generation of the electricity. 1839 – Christian Friedrich Schönbein published the principle of the fuel cell in the "Philosophical Magazine".
1839 – William Robert Grove developed the Grove cell. 1842 – William Robert Grove developed the first fuel cell 1849 – Eugène Bourdon – Bourdon gauge 1863 – Etienne Lenoir made a test drive from Paris to Joinville-le-Pont with the 1-cylinder, 2-stroke Hippomobile. 1866 – August Wilhelm von Hofmann invents the Hofmann voltameter for the electrolysis of water. 1873 – Thaddeus S. C. Lowe – Water gas, the process used the water gas shift reaction. 1874 – Jules Verne – The Mysterious Island, "water will one day be employed as fuel, that hydrogen and oxygen of which it is constituted will be used" 1884 – Charles Renard and Arthur Constantin Krebs launch the airship La France. 1885 – Zygmunt Florenty Wróblewski published hydrogen's critical temperature as 33 K. 1893 – Friedrich Wilhelm Ostwald experimentally determined the interconnected roles of the various components of the fuel cell. 1895 – Hydrolysis 1896 – Jackson D. D. and Ellms J. W. hydrogen production by microalgae 1896 – Leon Teisserenc de Bort carries out experiments with high flying instrumental weather balloons.
1897 – Paul Sabatier facilitated the use of hydrogenation with the discovery of the Sabatier reaction. 1898 – James Dewar liquefied hydrogen by using regenerative cooling and his invention, the vacuum flask at the Royal Institution of Great Britain in London. 1899 – James Dewar collected solid hydrogen for the first time. 1900 – Count Ferdinand von Zeppelin launched the first hydrogen-filled Zeppelin LZ1 airship. 1901 – Wilhelm Normann introduced the hydrogenation of fats. 1903 – Konstantin Eduardovich Tsiolkovskii published "The Exploration of Cosmic Space by Means of Reaction Devices" 1907 – Lane hydrogen producer 1909 – Count Ferdinand Adolf August von Zeppelin made the first long distance flight with the Zeppelin LZ5. 1909 – Linde–Frank–Caro process 1910 – The first Zeppelin passenger flight with the Zeppelin LZ7. 1910 – Fritz Haber patented the Haber process. 1912 – The first scheduled international Zeppelin passenger flights with the Zeppelin LZ13. 1913 – Niels Bohr explains the Rydberg formula for the spectrum of hydrogen by imposing a quantization condition on classical orbits of the electron in hydrogen 1919 – The first Atlantic crossing by airship with the Beardmore HMA R34.
1920 – Hydrocracking, a plant for the commercial
Potassium is a chemical element with symbol K and atomic number 19. It was first isolated from the ashes of plants, from which its name derives. In the periodic table, potassium is one of the alkali metals. All of the alkali metals have a single valence electron in the outer electron shell, removed to create an ion with a positive charge – a cation, which combines with anions to form salts. Potassium in nature occurs only in ionic salts. Elemental potassium is a soft silvery-white alkali metal that oxidizes in air and reacts vigorously with water, generating sufficient heat to ignite hydrogen emitted in the reaction, burning with a lilac-colored flame, it is found dissolved in sea water, is part of many minerals. Potassium is chemically similar to sodium, the previous element in group 1 of the periodic table, they have a similar first ionization energy, which allows for each atom to give up its sole outer electron. That they are different elements that combine with the same anions to make similar salts was suspected in 1702, was proven in 1807 using electrolysis.
Occurring potassium is composed of three isotopes, of which 40K is radioactive. Traces of 40K are found in all potassium, it is the most common radioisotope in the human body. Potassium ions are vital for the functioning of all living cells; the transfer of potassium ions across nerve cell membranes is necessary for normal nerve transmission. Fresh fruits and vegetables are good dietary sources of potassium; the body responds to the influx of dietary potassium, which raises serum potassium levels, with a shift of potassium from outside to inside cells and an increase in potassium excretion by the kidneys. Most industrial applications of potassium exploit the high solubility in water of potassium compounds, such as potassium soaps. Heavy crop production depletes the soil of potassium, this can be remedied with agricultural fertilizers containing potassium, accounting for 95% of global potassium chemical production; the English name for the element potassium comes from the word "potash", which refers to an early method of extracting various potassium salts: placing in a pot the ash of burnt wood or tree leaves, adding water and evaporating the solution.
When Humphry Davy first isolated the pure element using electrolysis in 1807, he named it potassium, which he derived from the word potash. The symbol "K" stems from kali, itself from the root word alkali, which in turn comes from Arabic: القَلْيَه al-qalyah "plant ashes". In 1797, the German chemist Martin Klaproth discovered "potash" in the minerals leucite and lepidolite, realized that "potash" was not a product of plant growth but contained a new element, which he proposed to call kali. In 1807, Humphry Davy produced the element via electrolysis: in 1809, Ludwig Wilhelm Gilbert proposed the name Kalium for Davy's "potassium". In 1814, the Swedish chemist Berzelius advocated the name kalium for potassium, with the chemical symbol "K"; the English and French speaking countries adopted Davy and Gay-Lussac/Thénard's name Potassium, while the Germanic countries adopted Gilbert/Klaproth's name Kalium. The "Gold Book" of the International Union of Physical and Applied Chemistry has designated the official chemical symbol as K.
Potassium is the second least dense metal after lithium. It is a soft solid with a low melting point, can be cut with a knife. Freshly cut potassium is silvery in appearance, but it begins to tarnish toward gray on exposure to air. In a flame test and its compounds emit a lilac color with a peak emission wavelength of 766.5 nanometers. Neutral potassium atoms have 19 electrons, one more than the stable configuration of the noble gas argon; because of this and its low first ionization energy of 418.8 kJ/mol, the potassium atom is much more to lose the last electron and acquire a positive charge than to gain one and acquire a negative charge. This process requires so little energy that potassium is oxidized by atmospheric oxygen. In contrast, the second ionization energy is high, because removal of two electrons breaks the stable noble gas electronic configuration. Potassium therefore does not form compounds with the oxidation state of higher. Potassium is an active metal that reacts violently with oxygen in water and air.
With oxygen it forms potassium peroxide, with water potassium forms potassium hydroxide. The reaction of potassium with water is dangerous because of its violent exothermic character and the production of hydrogen gas. Hydrogen reacts again with atmospheric oxygen, producing water, which reacts with the remaining potassium; this reaction requires only traces of water. Because of the sensitivity of potassium to water and air, reactions with other elements are possible only in an inert atmosphere such as argon gas using air-free techniques. Potassium does not react with most hydrocarbons such as mineral kerosene, it dissolves in liquid ammonia, up to 480 g per 1000 g of ammonia at 0 °C. Depending on the concentration, the ammonia solutions are blue to yellow, their electrical conductivity is similar to that of liquid metals. In a pure solution, potassium reacts with ammonia to form KNH2, but this reaction is accelerated by minute amounts of transition metal s
Calcium is a chemical element with symbol Ca and atomic number 20. As an alkaline earth metal, calcium is a reactive metal that forms a dark oxide-nitride layer when exposed to air, its physical and chemical properties are most similar to its heavier homologues strontium and barium. It is the fifth most abundant element in Earth's crust and the third most abundant metal, after iron and aluminium; the most common calcium compound on Earth is calcium carbonate, found in limestone and the fossilised remnants of early sea life. The name derives from Latin calx "lime", obtained from heating limestone; some calcium compounds were known to the ancients, though their chemistry was unknown until the seventeenth century. Pure calcium was isolated in 1808 via electrolysis of its oxide by Humphry Davy, who named the element. Calcium compounds are used in many industries: in foods and pharmaceuticals for calcium supplementation, in the paper industry as bleaches, as components in cement and electrical insulators, in the manufacture of soaps.
On the other hand, the metal in pure form has few applications due to its high reactivity. Calcium is the fifth-most abundant element in the human body; as electrolytes, calcium ions play a vital role in the physiological and biochemical processes of organisms and cells: in signal transduction pathways where they act as a second messenger. Calcium ions outside cells are important for maintaining the potential difference across excitable cell membranes as well as proper bone formation. Calcium is a ductile silvery metal whose properties are similar to the heavier elements in its group, strontium and radium. A calcium atom has twenty electrons, arranged in the electron configuration 4s2. Like the other elements placed in group 2 of the periodic table, calcium has two valence electrons in the outermost s-orbital, which are easily lost in chemical reactions to form a dipositive ion with the stable electron configuration of a noble gas, in this case argon. Hence, calcium is always divalent in its compounds, which are ionic.
Hypothetical univalent salts of calcium would be stable with respect to their elements, but not to disproportionation to the divalent salts and calcium metal, because the enthalpy of formation of MX2 is much higher than those of the hypothetical MX. This occurs because of the much greater lattice energy afforded by the more charged Ca2+ cation compared to the hypothetical Ca+ cation. Calcium, strontium and radium are always considered to be alkaline earth metals. Beryllium and magnesium are different from the other members of the group in their physical and chemical behaviour: they behave more like aluminium and zinc and have some of the weaker metallic character of the post-transition metals, why the traditional definition of the term "alkaline earth metal" excludes them; this classification is obsolete in English-language sources, but is still used in other countries such as Japan. As a result, comparisons with strontium and barium are more germane to calcium chemistry than comparisons with magnesium.
Calcium metal melts at 842 °C and boils at 1494 °C. It crystallises in the face-centered cubic arrangement like strontium, its density of 1.55 g/cm3 is the lowest in its group. Calcium can be cut with a knife with effort. While calcium is a poorer conductor of electricity than copper or aluminium by volume, it is a better conductor by mass than both due to its low density. While calcium is infeasible as a conductor for most terrestrial applications as it reacts with atmospheric oxygen, its use as such in space has been considered; the chemistry of calcium is that of a typical heavy alkaline earth metal. For example, calcium spontaneously reacts with water more than magnesium and less than strontium to produce calcium hydroxide and hydrogen gas, it reacts with the oxygen and nitrogen in the air to form a mixture of calcium oxide and calcium nitride. When finely divided, it spontaneously burns in air to produce the nitride. In bulk, calcium is less reactive: it forms a hydration coating in moist air, but below 30% relative humidity it may be stored indefinitely at room temperature.
Besides the simple oxide CaO, the peroxide CaO2 can be made by direct oxidation of calcium metal under a high pressure of oxygen, there is some evidence for a yellow superoxide Ca2. Calcium hydroxide, Ca2, is a strong base, though it is not as strong as the hydroxides of strontium, barium or the alkali metals. All four dihalides of calcium are known. Calcium carbonate and calcium sulfate are abundant minerals. Like strontium and barium, as well as the alkali metals and the divalent lanthanides europium and ytterbium, calcium metal dissolves directly in liquid ammonia to give a dark blue solution. Due to the large size of the Ca2+ ion, high coordination numbers are common, up to 24 in some intermetallic compounds such as CaZn13. Calcium is complexed by oxygen chelates such as EDTA and polyphosphates, which are useful in an
Stoichiometry is the calculation of reactants and products in chemical reactions. Stoichiometry is founded on the law of conservation of mass where the total mass of the reactants equals the total mass of the products, leading to the insight that the relations among quantities of reactants and products form a ratio of positive integers; this means that if the amounts of the separate reactants are known the amount of the product can be calculated. Conversely, if one reactant has a known quantity and the quantity of the products can be empirically determined the amount of the other reactants can be calculated; this is illustrated in the image here, where the balanced equation is: CH4 + 2 O2 → CO2 + 2 H2O. Here, one molecule of methane reacts with two molecules of oxygen gas to yield one molecule of carbon dioxide and two molecules of water; this particular chemical equation is an example of complete combustion. Stoichiometry measures these quantitative relationships, is used to determine the amount of products and reactants that are produced or needed in a given reaction.
Describing the quantitative relationships among substances as they participate in chemical reactions is known as reaction stoichiometry. In the example above, reaction stoichiometry measures the relationship between the methane and oxygen as they react to form carbon dioxide and water; because of the well known relationship of moles to atomic weights, the ratios that are arrived at by stoichiometry can be used to determine quantities by weight in a reaction described by a balanced equation. This is called composition stoichiometry. Gas stoichiometry deals with reactions involving gases, where the gases are at a known temperature and volume and can be assumed to be ideal gases. For gases, the volume ratio is ideally the same by the ideal gas law, but the mass ratio of a single reaction has to be calculated from the molecular masses of the reactants and products. In practice, due to the existence of isotopes, molar masses are used instead when calculating the mass ratio; the term stoichiometry was first used by Jeremias Benjamin Richter in 1792 when the first volume of Richter's Stoichiometry or the Art of Measuring the Chemical Elements was published.
The term is derived from the Ancient Greek words στοιχεῖον stoicheion "element" and μέτρον metron "measure". In patristic Greek, the word Stoichiometria was used by Nicephorus to refer to the number of line counts of the canonical New Testament and some of the Apocrypha. A stoichiometric amount or stoichiometric ratio of a reagent is the optimum amount or ratio where, assuming that the reaction proceeds to completion: All of the reagent is consumed There is no deficiency of the reagent There is no excess of the reagent. Stoichiometry rests upon the basic laws that help to understand it better, i.e. law of conservation of mass, the law of definite proportions, the law of multiple proportions and the law of reciprocal proportions. In general, chemical reactions combine in definite ratios of chemicals. Since chemical reactions can neither create nor destroy matter, nor transmute one element into another, the amount of each element must be the same throughout the overall reaction. For example, the number of atoms of a given element X on the reactant side must equal the number of atoms of that element on the product side, whether or not all of those atoms are involved in a reaction.
Chemical reactions, as macroscopic unit operations, consist of a large number of elementary reactions, where a single molecule reacts with another molecule. As the reacting molecules consist of a definite set of atoms in an integer ratio, the ratio between reactants in a complete reaction is in integer ratio. A reaction may consume more than one molecule, the stoichiometric number counts this number, defined as positive for products and negative for reactants. Different elements have a different atomic mass, as collections of single atoms, molecules have a definite molar mass, measured with the unit mole. By definition, carbon-12 has a molar mass of 12 g/mol. Thus, to calculate the stoichiometry by mass, the number of molecules required for each reactant is expressed in moles and multiplied by the molar mass of each to give the mass of each reactant per mole of reaction; the mass ratios can be calculated by dividing each by the total in the whole reaction. Elements in their natural state are mixtures of isotopes of differing mass, thus atomic masses and thus molar masses are not integers.
For instance, instead of an exact 14:3 proportion, 17.04 kg of ammonia consists of 14.01 kg of nitrogen and 3 × 1.01 kg of hydrogen, because natural nitrogen includes a small amount of nitrogen-15, natural hydrogen includes hydrogen-2. A stoichiometric reactant is a reactant, consumed in a reaction, as opposed to a catalytic reactant, not consumed in the overall reaction because it reacts in one step and is regenerated in another step. Stoichiometry is not only used to balance chemical equations but used in conversions, i.e. converting from grams to moles using molar mass as the conversion factor, or from grams to milliliters using density. For example, to find the amount of NaCl in 2.00 g, one would do the following: 2.00 g NaCl 58.44 g NaCl mol − 1 = 0.034 mol In the above example, when written out in fraction form, the units of grams form a multiplicative identity, equivalent to one, wit
Copper extraction refers to the methods used to obtaining copper from its ores. The conversion of copper consists of a series of electrochemical processes. Methods have evolved and vary with country depending on the ore source, local environmental regulations, other factors; as in all mining operations, the ore must be beneficiated. The processing techniques depend on the nature of the ore. If the ore is sulfide copper minerals, the ore is crushed and ground to liberate the valuable minerals from the waste minerals, it is concentrated using mineral flotation. The concentrate is then sold to distant smelters, although some large mines have smelters located nearby; such colocation of mines and smelters was more typical in the 19th and early 20th centuries, when smaller smelters could be economic. The sulfide concentrates are smelted in such furnaces as the Outokumpo or Inco flash furnace or the ISASMELT furnace to produce matte, which must be converted and refined to produce anode copper; the final refining process is electrolysis.
For economic and environmental reasons, many of the byproducts of extraction are reclaimed. Sulfur dioxide gas, for example, is captured and turned into sulfuric acid — which can be used in the extraction process or sold for such purposes as fertiliser manufacture. Oxidised copper ores can be treated by hydrometallurgical extraction; the earliest evidence of cold-hammering of native copper comes from the excavation at Çaÿonü Tepesi in eastern Anatolia, which dates between 7200 to 6600 BCE. Among the various items considered to be votive or amulets there was one that looked like a fishhook and one like an awl. Another find, at Shanidar Cave in Mergasur, contained copper beads, dates to 8,700 BCE; the world's oldest known copper mine, as opposed to usage of surface deposits, is at Timna Valley, since the fourth millennium BC, with smelting and surface deposit usage since the sixth to fifth millennium. Pločnik archaeological site in southeastern Europe contains the oldest securely dated evidence of copper making at high temperature, from 5,000 BCE.
The find in June 2010 extends for additional 500 years the earlier record of copper smelting from Rudna Glava, dated to 5th millennium BCE. Copper smelting technology gave rise to the Copper Age, aka Chalcolithic Age, the Bronze Age; the Bronze Age would not have been possible without humans developing smelting technology. Most copper ores contain only a small percentage of copper metal bound up within valuable ore minerals, with the remainder of the ore being unwanted rock or gangue minerals silicate minerals or oxide minerals for which there is no value. In some cases, tailings have been retreated to recover lost value as the technology for recovering copper has improved; the average grade of copper ores in the 21st century is below 0.6% copper, with a proportion of economic ore minerals being less than 2% of the total volume of the ore rock. A key objective in the metallurgical treatment of any ore is the separation of ore minerals from gangue minerals within the rock; the first stage of any process within a metallurgical treatment circuit is accurate grinding or comminution, where the rock is crushed to produce small particles consisting of individual mineral phases.
These particles are separated to remove gangue, thereafter followed by a process of physical liberation of the ore minerals from the rock. The process of liberation of copper ores depends upon whether they are sulfide ores. Subsequent steps depend on the nature of the ore containing the copper. For oxide ores, a hydrometallurgical liberation process is undertaken, which uses the soluble nature of the ore minerals to the advantage of the metallurgical treatment plant. For sulfide ores, both secondary and primary, froth flotation is used to physically separate ore from gangue. For special native copper bearing ore bodies or sections of ore bodies rich in supergene native copper, this mineral can be recovered by a simple gravity circuit; the modern froth flotation process was independently invented the early 1900s in Australia by C. V Potter and around the same time by G. D. Delprat. All primary sulfide ores of copper sulfides, most concentrates of secondary copper sulfides, are subjected to smelting.
Some vat leach or pressure leach processes exist to solubilise chalcocite concentrates and produce copper cathode from the resulting leachate solution, but this is a minor part of the market. Carbonate concentrates are a minor product produced from copper cementation plants as the end-stage of a heap-leach operation; such carbonate concentrates can be treated by a solvent extraction and electrowinning plant or smelted. The copper ore is crushed and ground to a size such that an acceptably high degree of liberation has occurred between the copper sulfide ore minerals and the gangue minerals; the ore is wet, suspended in a slurry, mixed with xanthates or other reagents, which render the sulfide particles hydrophobic. Typical reagents include potassium ethylxanthate and sodium ethylxanthate, but dithiophosphates and dithiocarbamates are used; the treated ore is introduced to a water-filled aeration tank containing surfactant such as methylisobutyl carbinol. Air is forced through the slurry and the air bubbles attach to the hydrophobic copper sulfide particles, which are conducted to the surface, where they form a froth and are skimmed off.
These skimmings are subjected to a cleaner-scavenger cell to remove excess silicates and to remove other sulfide minerals that can deleteriousl
Oxygen is the chemical element with the symbol O and atomic number 8. It is a member of the chalcogen group on the periodic table, a reactive nonmetal, an oxidizing agent that forms oxides with most elements as well as with other compounds. By mass, oxygen is the third-most abundant element in the universe, after helium. At standard temperature and pressure, two atoms of the element bind to form dioxygen, a colorless and odorless diatomic gas with the formula O2. Diatomic oxygen gas constitutes 20.8% of the Earth's atmosphere. As compounds including oxides, the element makes up half of the Earth's crust. Dioxygen is used in cellular respiration and many major classes of organic molecules in living organisms contain oxygen, such as proteins, nucleic acids and fats, as do the major constituent inorganic compounds of animal shells and bone. Most of the mass of living organisms is oxygen as a component of water, the major constituent of lifeforms. Oxygen is continuously replenished in Earth's atmosphere by photosynthesis, which uses the energy of sunlight to produce oxygen from water and carbon dioxide.
Oxygen is too chemically reactive to remain a free element in air without being continuously replenished by the photosynthetic action of living organisms. Another form of oxygen, ozone absorbs ultraviolet UVB radiation and the high-altitude ozone layer helps protect the biosphere from ultraviolet radiation. However, ozone present at the surface is a byproduct of thus a pollutant. Oxygen was isolated by Michael Sendivogius before 1604, but it is believed that the element was discovered independently by Carl Wilhelm Scheele, in Uppsala, in 1773 or earlier, Joseph Priestley in Wiltshire, in 1774. Priority is given for Priestley because his work was published first. Priestley, called oxygen "dephlogisticated air", did not recognize it as a chemical element; the name oxygen was coined in 1777 by Antoine Lavoisier, who first recognized oxygen as a chemical element and characterized the role it plays in combustion. Common uses of oxygen include production of steel and textiles, brazing and cutting of steels and other metals, rocket propellant, oxygen therapy, life support systems in aircraft, submarines and diving.
One of the first known experiments on the relationship between combustion and air was conducted by the 2nd century BCE Greek writer on mechanics, Philo of Byzantium. In his work Pneumatica, Philo observed that inverting a vessel over a burning candle and surrounding the vessel's neck with water resulted in some water rising into the neck. Philo incorrectly surmised that parts of the air in the vessel were converted into the classical element fire and thus were able to escape through pores in the glass. Many centuries Leonardo da Vinci built on Philo's work by observing that a portion of air is consumed during combustion and respiration. In the late 17th century, Robert Boyle proved. English chemist John Mayow refined this work by showing that fire requires only a part of air that he called spiritus nitroaereus. In one experiment, he found that placing either a mouse or a lit candle in a closed container over water caused the water to rise and replace one-fourteenth of the air's volume before extinguishing the subjects.
From this he surmised that nitroaereus is consumed in both combustion. Mayow observed that antimony increased in weight when heated, inferred that the nitroaereus must have combined with it, he thought that the lungs separate nitroaereus from air and pass it into the blood and that animal heat and muscle movement result from the reaction of nitroaereus with certain substances in the body. Accounts of these and other experiments and ideas were published in 1668 in his work Tractatus duo in the tract "De respiratione". Robert Hooke, Ole Borch, Mikhail Lomonosov, Pierre Bayen all produced oxygen in experiments in the 17th and the 18th century but none of them recognized it as a chemical element; this may have been in part due to the prevalence of the philosophy of combustion and corrosion called the phlogiston theory, the favored explanation of those processes. Established in 1667 by the German alchemist J. J. Becher, modified by the chemist Georg Ernst Stahl by 1731, phlogiston theory stated that all combustible materials were made of two parts.
One part, called phlogiston, was given off when the substance containing it was burned, while the dephlogisticated part was thought to be its true form, or calx. Combustible materials that leave little residue, such as wood or coal, were thought to be made of phlogiston. Air did not play a role in phlogiston theory, nor were any initial quantitative experiments conducted to test the idea. Polish alchemist and physician Michael Sendivogius in his work De Lapide Philosophorum Tractatus duodecim e naturae fonte et manuali experientia depromti described a substance contained in air, referring to it as'cibus vitae', this substance is identical with oxygen. Sendivogius, during his experiments performed between 1598 and 1604, properly recognized that the substance is equivalent to the gaseous byproduct released by the thermal decomposition of potassium nitrate. In Bugaj’s view, the isolation of oxygen and the proper association of the substance to that part of air, required for life, lends sufficient weight to the discovery of oxygen by Sendivogius.