Lead is a chemical element with symbol Pb and atomic number 82. It is a heavy metal, denser than most common materials. Lead is soft and malleable, has a low melting point; when freshly cut, lead is silvery with a hint of blue. Lead has the highest atomic number of any stable element and three of its isotopes each include a major decay chain of heavier elements. Lead is a unreactive post-transition metal, its weak metallic character is illustrated by its amphoteric nature. Compounds of lead are found in the +2 oxidation state rather than the +4 state common with lighter members of the carbon group. Exceptions are limited to organolead compounds. Like the lighter members of the group, lead tends to bond with itself. Lead is extracted from its ores. Galena, a principal ore of lead bears silver, interest in which helped initiate widespread extraction and use of lead in ancient Rome. Lead production declined after the fall of Rome and did not reach comparable levels until the Industrial Revolution. In 2014, the annual global production of lead was about ten million tonnes, over half of, from recycling.
Lead's high density, low melting point and relative inertness to oxidation make it useful. These properties, combined with its relative abundance and low cost, resulted in its extensive use in construction, batteries and shot, solders, fusible alloys, white paints, leaded gasoline, radiation shielding. In the late 19th century, lead's toxicity was recognized, its use has since been phased out of many applications. However, many countries still allow the sale of products that expose humans to lead, including some types of paints and bullets. Lead is a toxin that accumulates in soft tissues and bones, it acts as a neurotoxin damaging the nervous system and interfering with the function of biological enzymes, causing neurological disorders, such as brain damage and behavioral problems. A lead atom has 82 electrons, arranged in an electron configuration of 4f145d106s26p2; the sum of lead's first and second ionization energies—the total energy required to remove the two 6p electrons—is close to that of tin, lead's upper neighbor in the carbon group.
This is unusual. The similarity of ionization energies is caused by the lanthanide contraction—the decrease in element radii from lanthanum to lutetium, the small radii of the elements from hafnium onwards; this is due to poor shielding of the nucleus by the lanthanide 4f electrons. The sum of the first four ionization energies of lead exceeds that of tin, contrary to what periodic trends would predict. Relativistic effects, which become significant in heavier atoms, contribute to this behavior. One such effect is the inert pair effect: the 6s electrons of lead become reluctant to participate in bonding, making the distance between nearest atoms in crystalline lead unusually long. Lead's lighter carbon group congeners form stable or metastable allotropes with the tetrahedrally coordinated and covalently bonded diamond cubic structure; the energy levels of their outer s- and p-orbitals are close enough to allow mixing into four hybrid sp3 orbitals. In lead, the inert pair effect increases the separation between its s- and p-orbitals, the gap cannot be overcome by the energy that would be released by extra bonds following hybridization.
Rather than having a diamond cubic structure, lead forms metallic bonds in which only the p-electrons are delocalized and shared between the Pb2+ ions. Lead has a face-centered cubic structure like the sized divalent metals calcium and strontium. Pure lead has a silvery appearance with a hint of blue, it tarnishes on contact with moist air and takes on a dull appearance, the hue of which depends on the prevailing conditions. Characteristic properties of lead include high density, malleability and high resistance to corrosion due to passivation. Lead's close-packed face-centered cubic structure and high atomic weight result in a density of 11.34 g/cm3, greater than that of common metals such as iron and zinc. This density is the origin of the idiom to go over like a lead balloon; some rarer metals are denser: tungsten and gold are both at 19.3 g/cm3, osmium—the densest metal known—has a density of 22.59 g/cm3 twice that of lead. Lead is a soft metal with a Mohs hardness of 1.5. It is somewhat ductile.
The bulk modulus of lead—a measure of its ease of compressibility—is 45.8 GPa. In comparison, that of aluminium is 75.2 GPa. Lead's tensile strength, at 12–17 MPa, is low; the melting point of lead—at 327.5 °C —is low compared to most metals. Its boiling point of 1749 °C is the lowest among the carbon group elements; the electrical resistivity of lead at 20 °C is 192 nanoohm-meters an order of magnitude higher than those of other industrial metals. Lead is a superconductor at temperatures lower than 7.19 K.
A battery is a device consisting of one or more electrochemical cells with external connections provided to power electrical devices such as flashlights and electric cars. When a battery is supplying electric power, its positive terminal is the cathode and its negative terminal is the anode; the terminal marked negative is the source of electrons that will flow through an external electric circuit to the positive terminal. When a battery is connected to an external electric load, a redox reaction converts high-energy reactants to lower-energy products, the free-energy difference is delivered to the external circuit as electrical energy; the term "battery" referred to a device composed of multiple cells, however the usage has evolved to include devices composed of a single cell. Primary batteries are discarded. Common examples are the alkaline battery used for flashlights and a multitude of portable electronic devices. Secondary batteries can be discharged and recharged multiple times using an applied electric current.
Examples include the lead-acid batteries used in vehicles and lithium-ion batteries used for portable electronics such as laptops and smartphones. Batteries come in many shapes and sizes, from miniature cells used to power hearing aids and wristwatches to small, thin cells used in smartphones, to large lead acid batteries or lithium-ion batteries in vehicles, at the largest extreme, huge battery banks the size of rooms that provide standby or emergency power for telephone exchanges and computer data centers. According to a 2005 estimate, the worldwide battery industry generates US$48 billion in sales each year, with 6% annual growth. Batteries have much lower specific energy than common fuels such as gasoline. In automobiles, this is somewhat offset by the higher efficiency of electric motors in converting chemical energy to mechanical work, compared to combustion engines; the usage of "battery" to describe a group of electrical devices dates to Benjamin Franklin, who in 1748 described multiple Leyden jars by analogy to a battery of cannon.
Italian physicist Alessandro Volta built and described the first electrochemical battery, the voltaic pile, in 1800. This was a stack of copper and zinc plates, separated by brine-soaked paper disks, that could produce a steady current for a considerable length of time. Volta did not understand, he thought that his cells were an inexhaustible source of energy, that the associated corrosion effects at the electrodes were a mere nuisance, rather than an unavoidable consequence of their operation, as Michael Faraday showed in 1834. Although early batteries were of great value for experimental purposes, in practice their voltages fluctuated and they could not provide a large current for a sustained period; the Daniell cell, invented in 1836 by British chemist John Frederic Daniell, was the first practical source of electricity, becoming an industry standard and seeing widespread adoption as a power source for electrical telegraph networks. It consisted of a copper pot filled with a copper sulfate solution, in, immersed an unglazed earthenware container filled with sulfuric acid and a zinc electrode.
These wet cells used liquid electrolytes, which were prone to leakage and spillage if not handled correctly. Many used glass jars to hold their components, which made them fragile and dangerous; these characteristics made. Near the end of the nineteenth century, the invention of dry cell batteries, which replaced the liquid electrolyte with a paste, made portable electrical devices practical. Batteries convert chemical energy directly to electrical energy. In many cases, the electrical energy released is the difference in the cohesive or bond energies of the metals, oxides, or molecules undergoing the electrochemical reaction. For instance, energy can be stored in Zn or Li, which are high-energy metals because they are not stabilized by d-electron bonding, unlike transition metals. Batteries are designed such that the energetically favorable redox reaction can occur only if electrons move through the external part of the circuit. A battery consists of some number of voltaic cells; each cell consists of two half-cells connected in series by a conductive electrolyte containing metal cations.
One half-cell includes electrolyte and the negative electrode, the electrode to which anions migrate. Cations are reduced at the cathode; some cells use different electrolytes for each half-cell. Each half-cell has an electromotive force relative to a standard; the net emf of the cell is the difference between the emfs of its half-cells. Thus, if the electrodes have emfs E 1 and E 2 the net emf is E 2 − E 1.
The electrolysis of brine is an industrial process for the electrolysis of sodium chloride solutions. It is the technology used to produce chlorine and sodium hydroxide, which are commodity chemicals required by industry. 35 million tons of chlorine were prepared by this process in 1987. Industrial scale production began in 1892; the process is conducted on a brine, in which case NaOH, chlorine result. When using calcium chloride or potassium chloride, the products contain calcium or potassium instead of sodium. Related processes are known that use molten NaCl to give chlorine and sodium metal or condensed hydrogen chloride to give hydrogen and chlorine; the process has a high energy consumption, for example over 4 billion kWh per year in West Germany in 1985. Because the process gives equivalent amounts of chlorine and sodium hydroxide, it is necessary to find a use for these products in the same proportion. For every mole of chlorine produced, one mole of hydrogen is produced. Much of this hydrogen is used to produce hydrochloric acid or ammonia, or is used in the hydrogenation of organic compounds.
Three production methods are in use. While the mercury cell method produces chlorine-free sodium hydroxide, the use of several tonnes of mercury leads to serious environmental problems. In a normal production cycle a few hundred pounds of mercury per year are emitted, which accumulate in the environment. Additionally, the chlorine and sodium hydroxide produced via the mercury-cell chloralkali process are themselves contaminated with trace amounts of mercury; the membrane and diaphragm method use no mercury, but the sodium hydroxide contains chlorine, which must be removed. The performance of these devices is governed by the considerations of electrochemical engineering; the most common chloralkali process involves the electrolysis of aqueous sodium chloride in a membrane cell. Saturated brine is passed into the first chamber of the cell where the chloride ions are oxidised at the anode, losing electrons to become chlorine gas: 2Cl− → Cl2 + 2e−At the cathode, positive hydrogen ions pulled from water molecules are reduced by the electrons provided by the electrolytic current, to hydrogen gas, releasing hydroxide ions into the solution: 2H2O + 2e− → H2 + 2OH−The ion-permeable ion exchange membrane at the center of the cell allows the sodium ions to pass to the second chamber where they react with the hydroxide ions to produce caustic soda.
The overall reaction for the electrolysis of brine is thus: 2NaCl + 2H2O → Cl2 + H2 + 2NaOHA membrane cell is used to prevent the reaction between the chlorine and hydroxide ions. If this reaction were to occur the chlorine would disproportionate to form chloride and hypochlorite ions: Cl2 + 2OH− → Cl− + ClO− + H2OAbove about 60 °C, chlorate can be formed: 3Cl2 + 6OH− → 5Cl− + ClO3− + 3H2OBecause of the corrosive nature of chlorine production, the anode must be made from a non-reactive metal such as titanium, whereas the cathode can be made from a more oxidized metal such as nickel. In the diaphragm cell process, there are two compartments separated by a permeable diaphragm made of asbestos fibers. Brine is introduced into flows into the cathode compartment. To the Membrane Cell, chloride ions are oxidized at the anode to produce chlorine, at the cathode, water is split into caustic soda and hydrogen; the diaphragm prevents the reaction of the caustic soda with the chlorine. A diluted caustic brine leaves the cell.
The caustic soda must be concentrated to 50% and the salt removed. This is done using an evaporative process with about three tonnes of steam per tonne of caustic soda; the salt separated from the caustic brine can be used to saturate diluted brine. The chlorine contains oxygen and must be purified by liquefaction and evaporation. In the mercury-cell process known as the Castner–Kellner process, a saturated brine solution floats on top of a thin layer of mercury; the mercury is the cathode, where sodium is produced and forms a sodium-mercury amalgam with the mercury. The amalgam is continuously drawn out of the cell and reacted with water which decomposes the amalgam into sodium hydroxide and mercury; the mercury is recycled into the electrolytic cell. Chlorine bubbles out of the cell. Mercury cells are being phased out due to concerns about mercury poisoning from mercury cell pollution such as occurred in Canada and Japan; the interests of chloralkali product manufacturers are represented at regional and international levels by associations such as Euro Chlor and The World Chlorine Council.
Electrolysis can be done with beakers, one containing a brine solution and one containing pure water connected by a salt bridge. Anodes are made ideally from platinum group metals. Since corrosion is less severe at the cathode, it can be stainless silver. Gas diffusion electrode Electrochemical engineering Solvay process Bommaraju, Tilak V.. "Brine Electrolysis." Electrochemistry Encyclopedia. Cleveland: Case Western Rsserve University
A metal is a material that, when freshly prepared, polished, or fractured, shows a lustrous appearance, conducts electricity and heat well. Metals are malleable or ductile. A metal may be an alloy such as stainless steel. In physics, a metal is regarded as any substance capable of conducting electricity at a temperature of absolute zero. Many elements and compounds that are not classified as metals become metallic under high pressures. For example, the nonmetal iodine becomes a metal at a pressure of between 40 and 170 thousand times atmospheric pressure; some materials regarded as metals can become nonmetals. Sodium, for example, becomes a nonmetal at pressure of just under two million times atmospheric pressure. In chemistry, two elements that would otherwise qualify as brittle metals—arsenic and antimony—are instead recognised as metalloids, on account of their predominately non-metallic chemistry. Around 95 of the 118 elements in the periodic table are metals; the number is inexact as the boundaries between metals and metalloids fluctuate due to a lack of universally accepted definitions of the categories involved.
In astrophysics the term "metal" is cast more to refer to all chemical elements in a star that are heavier than the lightest two and helium, not just traditional metals. A star fuses lighter atoms hydrogen and helium, into heavier atoms over its lifetime. Used in that sense, the metallicity of an astronomical object is the proportion of its matter made up of the heavier chemical elements. Metals are present in many aspects of modern life; the strength and resilience of some metals has led to their frequent use in, for example, high-rise building and bridge construction, as well as most vehicles, many home appliances, tools and railroad tracks. Precious metals were used as coinage, but in the modern era, coinage metals have extended to at least 23 of the chemical elements; the history of metals is thought to begin with the use of copper about 11,000 years ago. Gold, iron and brass were in use before the first known appearance of bronze in the 5th millennium BCE. Subsequent developments include the production of early forms of steel.
Metals are lustrous, at least when freshly prepared, polished, or fractured. Sheets of metal thicker than a few micrometres appear opaque; the solid or liquid state of metals originates in the capacity of the metal atoms involved to lose their outer shell electrons. Broadly, the forces holding an individual atom’s outer shell electrons in place are weaker than the attractive forces on the same electrons arising from interactions between the atoms in the solid or liquid metal; the electrons involved become delocalised and the atomic structure of a metal can be visualised as a collection of atoms embedded in a cloud of mobile electrons. This type of interaction is called a metallic bond; the strength of metallic bonds for different elemental metals reaches a maximum around the center of the transition metal series, as these elements have large numbers of delocalized electrons. Although most elemental metals have higher densities than most nonmetals, there is a wide variation in their densities, lithium being the least dense and osmium the most dense.
Magnesium and titanium are light metals of significant commercial importance. Their respective densities of 1.7, 2.7 and 4.5 g/cm3 can be compared to those of the older structural metals, like iron at 7.9 and copper at 8.9 g/cm3. An iron ball would thus weigh about as much as three aluminium balls. Metals are malleable and ductile, deforming under stress without cleaving; the nondirectional nature of metallic bonding is thought to contribute to the ductility of most metallic solids. In contrast, in an ionic compound like table salt, when the planes of an ionic bond slide past one another, the resultant change in location shifts ions of the same charge into close proximity, resulting in the cleavage of the crystal; such a shift is not observed in a covalently bonded crystal, such as a diamond, where fracture and crystal fragmentation occurs. Reversible elastic deformation in metals can be described by Hooke's Law for restoring forces, where the stress is linearly proportional to the strain. Heat or forces larger than a metal's elastic limit may cause a permanent deformation, known as plastic deformation or plasticity.
An applied force may be a compressive force, or a shear, bending or torsion force. A temperature change may affect the movement or displacement of structural defects in the metal such as grain boundaries, point vacancies and screw dislocations, stacking faults and twins in both crystalline and non-crystalline metals. Internal slip and metal fatigue may ensue; the atoms of metallic substances are arranged in one of three common crystal structures, namely body-centered cubic, face-centered cubic, hexagonal close-packed. In bcc, each atom is positioned at the center of a cube of eight others. In fcc and hcp, each atom is surrounded by twelve others; some metals adopt different structures depending on the temperature. The
Gold extraction refers to the processes required to extract gold from its ores. This may require a combination of comminution, mineral processing, hydrometallurgical, pyrometallurgical processes to be performed on the ore. Gold mining from alluvium ores was once achieved by techniques associated with placer mining such as simple gold panning and sluicing, resulting in direct recovery of small gold nuggets and flakes. Placer mining techniques since the mid to late 20th century have only been the practice of artisan miners. Hydraulic mining was used in the Californian gold rush, involved breaking down alluvial deposits with high-pressure jets of water. Hard rock ores have formed the basis of the majority of commercial gold recovery operations since the middle of the 20th century where open pit and or sub-surface mining techniques are used. Once the ore is mined it can be treated as a whole ore using a dump leaching or heap leaching processes; this is typical of oxide deposits. The ore is crushed and agglomerated prior to heap leaching.
High grade ores and ores resistant to cyanide leaching at coarse particle sizes, require further processing in order to recover the gold values. The processing techniques can include grinding, concentration and pressure oxidation prior to cyanidation; the smelting of gold began sometime around 6000 - 3000 BC. According to one source the technique began to be in use in Syria. In ancient Greece, Heraclitus wrote on the subject. According to de Lecerda and Salomons mercury was first in use for extraction at about 1000 BC, according to Meech and others, mercury was used in obtaining gold until the latter period of the first millennia. A technique known to Pliny the Elder was extraction by way of crushing and applying heat, with the resultant material powdered; the solubility of gold in a water and cyanide solution was discovered in 1783 by Carl Wilhelm Scheele, but it was not until the late 19th century, that an industrial process was developed. The expansion of gold mining in the Rand of South Africa began to slow down in the 1880s, as the new deposits being found tended to be pyritic ore.
The gold could not be extracted from this compound with any of the available chemical processes or technologies. In 1887, John Stewart MacArthur, working in collaboration with brothers Dr Robert and Dr William Forrest for the Tennant Company in Glasgow, developed the MacArthur-Forrest Process for the extraction of gold ores. By suspending the crushed ore in a cyanide solution, a separation of up to 96 percent pure gold was achieved; the process was first used on a large scale at the Witwatersrand in 1890, leading to a boom of investment as larger gold mines were opened up. In 1896, Bodländer confirmed that oxygen was necessary for the process, something, doubted by MacArthur, discovered that hydrogen peroxide was formed as an intermediate; the method known as heap leaching was first proposed in 1969 by the United States Bureau of Mines, was in use by the 1970s. Gold occurs principally as a native metal alloyed to a greater or lesser extent with silver, or sometimes with mercury. Native gold can occur as sizeable nuggets, as fine grains or flakes in alluvial deposits, or as grains or microscopic particles embedded in rock minerals.
Ores in which gold occurs in chemical composition with other elements are comparatively rare. They include calaverite, nagyagite and krennerite. Gravity concentration has been the most important way of extracting the native metal using pans or washing tables. Amalgamation with mercury was used to enhance recovery by adding it directly to the riffle tables, mercury is still used in small diggings across the world. However, froth flotation processes may be used to concentrate the gold. In some cases when the gold is present in the ore as discrete coarse particles, a gravity concentrate can be directly smelted to form gold bars. In other cases when the gold is present in the ore as fine particles or is not sufficiently liberated from the host rock, the concentrates are treated with cyanide salts, a process known as cyanidation leaching, followed by recovery from the leach solution. Recovery from solution involves adsorption on activated carbon stripping the gold from the carbon and passing the pregnant solution through electrowinning and onto the smelting process.
Froth flotation is applied when the gold present in an ore is associated with sulfide minerals such as pyrite, chalcopyrite or arsenopyrite, when such sulfides are present in large quantities in the ore. In this case, concentration of the sulfides results in concentration of gold values. Recovery of the gold from the sulfide concentrates requires further processing by roasting or wet pressure oxidation; these pyrometallurgical or hydrometallurgical treatments are themselves followed by cyanidation and carbon adsorption techniques for final recovery of the gold. Sometimes gold is present as a minor constituent in a base metal concentrate, is recovered as a by-product during production of the base metal. For example, it can be recovered in the anode slime during the electrorefining process; the results of laboratory and bench experimental studies can be used with sufficient accuracy to implement adhesive gold dressing of fine and nano-size particles within a “few hundred microns – a few tens of nanometers”.
If the gold can not be concentrated for smelting it is leached by an aqueous solution: The cyanide process is the industry standard. Thiosulfate leaching has been proven to be effective on ores with high soluble c
Jöns Jacob Berzelius
Baron Jöns Jacob Berzelius, named by himself and contemporary society as Jacob Berzelius, was a Swedish chemist. Berzelius is considered, along with Robert Boyle, John Dalton, Antoine Lavoisier, to be one of the founders of modern chemistry. Berzelius began his career as a physician but his researches in physical chemistry were of lasting significance in the development of the subject, he is noted for his determination of atomic weights. In 1803 Berzelius demonstrated the power of an electrochemical cell to decompose chemicals into pairs of electrically opposite constituents. Berzelius's work with atomic weights and his theory of electrochemical dualism led to his development of a modern system of chemical formula notation that could portray the composition of any compound both qualitatively and quantitatively, his system abbreviated the Latin names of the elements with one or two letters and applied subscripts to designate the number of atoms of each element present in both the acidic and basic ingredients.
Berzelius himself isolated several new elements, including cerium and thorium. Berzelius’s interest in mineralogy fostered his analysis and preparation of new compounds of these and other elements; the mineral berzelianite was named after him. He was a strict empiricist and insisted that any new theory be consistent with the sum of chemical knowledge, he developed classical analytical techniques, investigated isomerism and catalysis, phenomena that owe their names to him. He became a member of the Royal Swedish Academy of Sciences in 1808 and served from 1818 as its principal functionary, the perpetual secretary, he is known in Sweden as "the Father of Swedish Chemistry". Berzelius Day is celebrated on 20 August in honour of him. Berzelius was born in the parish of Väversunda in Östergötland in Sweden, his father was a school teacher in his mother a homemaker. Berzelius lost both his parents at an early age. Relatives in Linköping took care of him, there he attended the school today known as Katedralskolan.
He enrolled at Uppsala University, where he learned the profession of medical doctor from 1796 to 1801. He worked as an apprentice with a physician in the Medevi mineral springs. During this time, he conducted analysis of the spring water. For his medical studies, he examined the influence of galvanic current on several diseases and graduated as M. D. in 1802. He worked as physician near Stockholm until the mine-owner Wilhelm Hisinger discovered his analytical abilities and provided him with a laboratory. Between 1808 and 1836, Berzelius worked together with Anna Sundström. In 1807, Berzelius was appointed professor in pharmacy at the Karolinska Institute. In 1808, he was elected a member of the Royal Swedish Academy of Sciences. At this time, the Academy had been stagnating for several years, since the era of romanticism in Sweden had led to less interest in the sciences. In 1818, Berzelius was elected the Academy's secretary and held the post until 1848. During Berzelius' tenure, he is credited with revitalising the Academy and bringing it into a second golden era.
He was elected a Foreign Honorary Member of the American Academy of Arts and Sciences in 1822. In 1827, he became correspondent of the Royal Institute of the Netherlands, in 1830 associate member. In 1837, he was elected a member of the Swedish Academy, on chair number 5. Not long after arriving to Stockholm he wrote a chemistry textbook for his medical students, from which point a long and fruitful career in chemistry began. In 1813, he published an essay on the proportions of elements in compounds; the essay commenced with a general description, introduced his new symbolism, examined all the known elements, included a table of specific weights, finished with a selection of compounds written in his new formalisation. In 1818, he compiled a table of relative atomic weights, where oxygen was set to 100, which included all of the elements known at the time; this work provided evidence in favour of the atomic theory proposed by John Dalton: that inorganic chemical compounds are composed of atoms combined in whole number amounts.
In discovering that atomic weights are not integer multiples of the weight of hydrogen, Berzelius disproved Prout's hypothesis that elements are built up from atoms of hydrogen. Berzelius's atomic weight tables was first published in a German translation of his Textbook of Chemistry in 1826. In order to aid his experiments, he developed a system of chemical notation in which the elements were given simple written labels—such as O for oxygen, or Fe for iron—with proportions noted by numbers; this is the same system used today, the only difference being that instead of the subscript number used today, Berzelius used a superscript. Berzelius is credited with identifying the chemical elements silicon, selenium and cerium. Students working in Berzelius's laboratory discovered lithium and vanadium. Berzelius discovered silicon by repeating an experiment performed by Thénard. In the experiment, Berzelius reacted silicon tetrafluoride with potassium metal and purified its product by washing it until it became a brown powde