In the nervous system, a synapse[1] is a structure that permits a neuron (or nerve cell) to pass an electrical or chemical signal to another neuron or to the target effector cell.
Synapses are essential to the transmission of nervous impulses from one neuron to another. Neurons are specialized to pass signals to individual target cells, and synapses are the means by which they do so. At a synapse, the plasma membrane of the signal-passing neuron (the presynaptic neuron) comes into close apposition with the membrane of the target (postsynaptic) cell. Both the presynaptic and postsynaptic sites contain extensive arrays of molecular machinery that link the two membranes together and carry out the signaling process. In many synapses, the presynaptic part is located on an axon and the postsynaptic part is located on a dendrite or soma. Astrocytes also exchange information with the synaptic neurons, responding to synaptic activity and, in turn, regulating neurotransmission.[2] Synapses (at least chemical synapses) are stabilized in position by synaptic adhesion molecules (SAMs) projecting from both the pre- and post-synaptic neuron and sticking together where they overlap; SAMs may also assist in the generation and functioning of synapses.[3]
https://en.wikipedia.org/wiki/Synapse
monoamine releaser, is a drug that induces the release of a monoamine neurotransmitter from the presynaptic neuron into the synapse, leading to an increase in the extracellular concentrations of the neurotransmitter.
https://en.wikipedia.org/wiki/Monoamine_releasing_agent
https://en.wikipedia.org/wiki/Monoamine_releasing_agent
https://en.wikipedia.org/wiki/Membrane_transport_protein
https://en.wikipedia.org/wiki/Reverse_transport
https://en.wikipedia.org/wiki/Vesicular_monoamine_transporter_2
https://en.wikipedia.org/wiki/2C-E
https://en.wikipedia.org/wiki/Benzylpiperazine
https://en.wikipedia.org/wiki/Norpropylhexedrine
https://en.wikipedia.org/wiki/Cathine
https://en.wikipedia.org/wiki/Serotonin
https://en.wikipedia.org/wiki/Tyramine
https://en.wikipedia.org/wiki/Propylhexedrine
https://en.wikipedia.org/wiki/Nitrous_oxide
https://en.wikipedia.org/wiki/Nitric_oxide
https://en.wikipedia.org/wiki/Purging_(gas)
https://en.wikipedia.org/wiki/Pyrophoricity
https://en.wikipedia.org/wiki/Argon
https://en.wikipedia.org/wiki/Hexane
https://en.wikipedia.org/wiki/Maleic_anhydride
https://en.wikipedia.org/wiki/Dehydrogenation
https://en.wikipedia.org/wiki/Slag
https://en.wikipedia.org/wiki/Ferroalloy
https://en.wikipedia.org/wiki/Electric_arc_furnace
https://en.wikipedia.org/wiki/Steelmaking
https://en.wikipedia.org/wiki/Foundry
Hydrogen gas is a reducing agent when it reacts with non-metals and an oxidizing agent when it reacts with metals.
https://en.wikipedia.org/wiki/Reducing_agent
An anode is an electrode of a polarized electrical device through which conventional current enters the device. This contrasts with a cathode, an electrode of the device through which conventional current leaves the device. A common mnemonic is ACID, for "anode current into device".[1] The direction of conventional current (the flow of positive charges) in a circuit is opposite to the direction of electron flow, so (negatively charged) electrons flow out the anode of a galvanic cell, into an outside or external circuit connected to the cell. For example, the end of a household battery marked with a "-" (minus) is the anode.
In both a galvanic cell and an electrolytic cell, the anode is the electrode at which the oxidation reaction occurs. In a galvanic cell the anode is the wire or plate having excess negative charge as a result of the oxidation reaction. In an electrolytic cell, the anode is the wire or plate upon which excess positive charge is imposed.[2] As a result of this, anions will tend to move towards the anode where they will undergo oxidation.
Historically, the anode of a galvanic cell was also known as the zincode because it was usually composed of zinc.[3][4]: pg. 209, 214
https://en.wikipedia.org/wiki/Anode
Hydrogen (whose reduction potential is 0.0) acts as an oxidizing agent because it accepts an electron donation from the reducing agent lithium (whose reduction potential is -3.04), which causes Li to be oxidized and hydrogen to be reduced.
- H2(g) + F2(g) → 2 HF(g)[b]
Hydrogen acts as a reducing agent because it donates its electrons to fluorine, which allows fluorine to be reduced.
https://en.wikipedia.org/wiki/Reducing_agent
The Great Oxidation Event (GOE), also called the Great Oxygenation Event, the Oxygen Catastrophe, the Oxygen Revolution, the Oxygen Crisis, or the Oxygen Holocaust,[2] was a time interval during the Paleoproterozoic era when the Earth's atmosphere and the shallow ocean first experienced a rise in the amount of oxygen.[3] This began approximately 2.460–2.426 Ga (billion years) ago, during the Siderian period, and ended approximately 2.060 Ga, during the Rhyacian.[4] Geological, isotopic, and chemical evidence suggests that biologically-produced molecular oxygen (dioxygen, O2) started to accumulate in Earth's atmosphere and changed it from a weakly reducing atmosphere practically free of oxygen into an oxidizing atmosphere containing abundant oxygen,[5] with oxygen levels being as high as 10% of their present atmospheric level by the end of the GOE.[6]
The sudden injection of toxic oxygen into an anaerobic biosphere may have caused the extinction of many existing anaerobic species on Earth. Although the event is inferred to have constituted a mass extinction,[7] due in part to the great difficulty in surveying microscopic species' abundances, and in part to the extreme age of fossil remains from that time, the Great Oxidation Event is typically not counted among conventional lists of "great extinctions", which are implicitly limited to the Phanerozoic eon. In any case, isotope geochemical data from sulfate minerals have been interpreted to indicate a decrease in the size of the biosphere of >80% associated with changes in nutrient supplies at the end of the GOE.[8]
The GOE is inferred to have been caused by cyanobacteria producing the oxygen, which may have enabled the subsequent development of multicellular life-forms like the Francevillian biota.[9]
https://en.wikipedia.org/wiki/Great_Oxidation_Event
Atmospheric pressure, also known as barometric pressure (after the barometer), is the pressure within the atmosphere of Earth. The standard atmosphere (symbol: atm) is a unit of pressure defined as 101,325 Pa (1,013.25 hPa), which is equivalent to 1013.25 millibars,[1] 760 mm Hg, 29.9212 inches Hg, or 14.696 psi.[2] The atm unit is roughly equivalent to the mean sea-level atmospheric pressure on Earth; that is, the Earth's atmospheric pressure at sea level is approximately 1 atm.
In most circumstances, atmospheric pressure is closely approximated by the hydrostatic pressure caused by the weight of air above the measurement point. As elevation increases, there is less overlying atmospheric mass, so atmospheric pressure decreases with increasing elevation. Because the atmosphere is thin relative to the Earth's radius—especially the dense atmospheric layer at low altitudes—the Earth's gravitational acceleration as a function of altitude can be approximated as constant and contributes little to this fall-off. Pressure measures force per unit area, with SI units of pascals (1 pascal = 1 newton per square metre, 1 N/m2). On average, a column of air with a cross-sectional area of 1 square centimetre (cm2), measured from the mean (average) sea level to the top of Earth's atmosphere, has a mass of about 1.03 kilogram and exerts a force or "weight" of about 10.1 newtons, resulting in a pressure of 10.1 N/cm2 or 101 kN/m2 (101 kilopascals, kPa). A column of air with a cross-sectional area of 1 in2 would have a weight of about 14.7 lbf, resulting in a pressure of 14.7 lbf/in2.
https://en.wikipedia.org/wiki/Atmospheric_pressure
https://en.wikipedia.org/wiki/Cyanobacteria
https://en.wikipedia.org/wiki/Carbon_monoxide
https://en.wikipedia.org/wiki/Banded_iron_formation
https://en.wikipedia.org/wiki/Lithium_aluminium_hydride
A substance is pyrophoric (from Greek: πυροφόρος, pyrophoros, 'fire-bearing') if it ignites spontaneously in air at or below 54 °C (129 °F) (for gases) or within 5 minutes after coming into contact with air (for liquids and solids).[1] Examples are organolithium compounds and triethylborane. Pyrophoric materials are often water-reactive as well and will ignite when they contact water or humid air. They can be handled safely in atmospheres of argon or (with a few exceptions) nitrogen. Class D fire extinguishers are designated for use in fires involving pyrophoric materials. A related concept is hypergolicity, in which two compounds spontaneously ignite when mixed.
https://en.wikipedia.org/wiki/Pyrophoricity
https://en.wikipedia.org/wiki/Triethylborane
https://en.wikipedia.org/wiki/Tetrahydrofuran
https://en.wikipedia.org/wiki/Coking
https://en.wikipedia.org/wiki/Dentistry
https://en.wikipedia.org/wiki/Induction_furnace
https://en.wikipedia.org/wiki/Dust
https://en.wikipedia.org/wiki/Blast_furnace
https://en.wikipedia.org/wiki/Smelting
https://en.wikipedia.org/wiki/Reducing_agent
The atmosphere of Earth is the layer of gases, known collectively as air, retained by Earth's gravity that surrounds the planet and forms its planetary atmosphere. The atmosphere of Earth creates pressure, absorbs most meteoroids and ultraviolet solar radiation, warms the surface through heat retention (greenhouse effect), allowing life and liquid water to exist on the Earth's surface, and reduces temperature extremes between day and night (the diurnal temperature variation).
https://en.wikipedia.org/wiki/Atmosphere_of_Earth
- Lithium aluminium hydride (LiAlH4), a very strong reducing agent
- Red-Al (NaAlH2(OCH2CH2OCH3)2), a safer and more stable alternative to lithium aluminum hydride
- Hydrogen without or with a suitable catalyst; e.g. a Lindlar catalyst
- Sodium amalgam (Na(Hg))
- Sodium-lead alloy (Na + Pb)
- Zinc amalgam (Zn(Hg)) (reagent for Clemmensen reduction)
- Diborane
- Sodium borohydride (NaBH4)
- Ferrous compounds that contain the Fe2+ ion, such as iron(II) sulfate
- Stannous compounds that contain the Sn2+ ion, such as tin(II) chloride
- Sulfur dioxide (sometimes also used as an oxidizing agent), Sulfite compounds
- Dithionates, e.g. Na2S2O6
- Thiosulfates, e.g. Na2S2O3 (mainly in analytical chemistry)[9]
- Iodides, such as potassium iodide (KI) (mainly in analytical chemistry)
- Hydrogen peroxide (H
2O
2) – mostly an oxidant but can occasionally act as a reducing agent (typically in analytical chemistry.) - Hydrazine (Wolff-Kishner reduction)
- Diisobutylaluminium hydride (DIBAL-H)
- Oxalic acid (C
2H
2O
4) - Formic acid (HCOOH)
- Ascorbic acid (C6H8O6)
- Reducing sugars, such as erythrose, see Aldose
- Phosphites, hypophosphites, and phosphorous acid
- Dithiothreitol (DTT) – used in biochemistry labs to avoid SS-bonds
- Carbon monoxide (CO)
- Cyanides in hydrochemical metallurgical processes
- Carbon (C)
- Tris-2-carboxyethylphosphine hydrochloride (TCEP)
- Corrosion – Gradual destruction of materials by chemical reaction with its environment
- Electrochemistry – Branch of chemistry
- Electrolyte – Ionic solids whose dissociation in water free up ions carrying the electrical current in solution
- Electron acceptor – Chemical entity capable of accepting electrons
- Electron donor – Chemical entity capable of donating electrons to another entity
- Electrosynthesis – Synthesis of chemical compounds in an electrochemical cell
- Organic reduction – Redox reaction that takes place with organic compounds
- Oxidizing agent – Chemical compound used to oxidize another substance in a chemical reaction
- Redox – Chemical reaction in which oxidation states of atoms are changed
- Reducing equivalent – Chemical reaction in which oxidation states of atoms are changed
Common reducing agents
 | This section needs additional citations for verification. (October 2016) |
See also
https://en.wikipedia.org/wiki/Reducing_agent
Sodium amalgam, commonly denoted Na(Hg), is an alloy of mercury and sodium. The term amalgam is used for alloys, intermetallic compounds, and solutions (both solid solutions and liquid solutions) involving mercury as a major component. Sodium amalgams are often used in reactions as strong reducing agents with better handling properties compared to solid sodium. They are less dangerously reactive toward water and in fact are often used as an aqueous suspension.
Sodium amalgam was used as a reagent as early as 1862.[1] A synthesis method was described by J. Alfred Wanklyn in 1866.[2]
Structure and compositions
No particular formula is assigned to "sodium amalgam". Na5Hg8 and Na3Hg are well defined compounds. In sodium amalgams, the Hg-Hg distances are expanded to around 5 Å vs. about 3 Å for mercury itself.[3] Usually amalgams are classified on the weight percent of sodium. Amalgams with 2% Na are solids at room temperature, whereas some more dilute amalgams remain liquid.[4]
Preparation
| External video | |
|---|---|
| Demonstration and commentary on the preparation of sodium amalgam | |
 via YouTube | |
 Mixing sodium with mercury (31 December 2018), by NileRed |
Metallic sodium dissolves in mercury exothermically, i.e. with the release of heat, therefore, formation of sodium amalgam is famously dangerous for generating sparks.[5] The process causes localised boiling of the mercury and for this reason the formation is usually conducted in a fume hood and often performed using air-free techniques, such as synthesis under anhydrous liquid paraffin. Sodium amalgam may be prepared in the laboratory by dissolving sodium metal in mercury or the reverse.[6][7][8] Sodium amalgams can be purchased from chemical supply houses.
Uses
Sodium amalgam has been used in organic chemistry as a powerful reducing agent, which is safer to handle than sodium itself. It is used in Emde degradation, and also for reduction of aromatic ketones to hydrols.[9]
A sodium amalgam is used in the design of the high pressure sodium lamp providing sodium to produce the proper color, and mercury to tailor the electrical characteristics of the lamp.
Mercury cell electrolysis
Sodium amalgam is a by-product of chlorine made by mercury cell electrolysis. In this cell, brine (concentrated sodium chloride solution) is electrolysed between a liquid mercury cathode and a titanium or graphite anode. Chlorine is formed at the anode, while sodium formed at the cathode dissolves into the mercury, making sodium amalgam. Normally this sodium amalgam is drawn off and reacted with water in a "decomposer cell" to produce hydrogen gas, concentrated sodium hydroxide solution, and mercury to be recycled through the process. In principle, all the mercury should be completely recycled, but inevitably a small portion goes missing. Because of concerns about this mercury escaping into the environment, the mercury cell process is generally being replaced by plants which use a less toxic cathode.
References
- Bachmann, W. E. (February 1933). "The Mechanism of Reduction by Sodium Amalgam and Alcohol. I. The Reduction of Aromatic Ketones to Hydrols". Journal of the American Chemical Society. 55 (2): 770–774. doi:10.1021/ja01329a051. ISSN 0002-7863.
External links
https://en.wikipedia.org/wiki/Sodium_amalgam
In metallurgy, a flux (from Latin fluxus 'flow') is a chemical cleaning agent, flowing agent, or purifying agent. Fluxes may have more than one function at a time. They are used in both extractive metallurgy and metal joining.
Some of the earliest known fluxes were sodium carbonate, potash, charcoal, coke, borax,[1] lime,[2] lead sulfide[3] and certain minerals containing phosphorus. Iron ore was also used as a flux in the smelting of copper. These agents served various functions, the simplest being a reducing agent, which prevented oxides from forming on the surface of the molten metal, while others absorbed impurities into slag, which could be scraped off molten metal.[4]
Fluxes are also used in foundries for removing impurities from molten nonferrous metals such as aluminium, or for adding desirable trace elements such as titanium.
As cleaning agents, fluxes facilitate soldering, brazing, and welding by removing oxidation from the metals to be joined. In some applications molten flux also serves as a heat-transfer medium, facilitating heating of the joint by the soldering tool or molten solder.
As cleaning agents, fluxes facilitate soldering, brazing, and welding by removing oxidation
from the metals to be joined. In some applications molten flux also
serves as a heat-transfer medium, facilitating heating of the joint by
the soldering tool or molten solder.
https://en.wikipedia.org/wiki/Flux_(metallurgy)
Potash (/ˈpɒtæʃ/) includes various mined and manufactured salts that contain potassium in water-soluble form.[1] The name derives from pot ash, plant ashes or wood ash soaked in water in a pot, the primary means of manufacturing potash before the Industrial Era. The word potassium is derived from potash.[2]
Potash is produced worldwide in amounts exceeding 90 million tonnes (40 million tonnes K2O equivalent) per year, with Canada being the largest producer, mostly for use in fertilizer. Various kinds of fertilizer-potash constitute the single greatest industrial use of the element potassium in the world. Potassium was first derived in 1807 by electrolysis of caustic potash (potassium hydroxide).[3]
https://en.wikipedia.org/wiki/Potash
https://en.wikipedia.org/wiki/Glass_transition
https://en.wikipedia.org/wiki/Elastomer
https://en.wikipedia.org/wiki/Damping
Roasting is a process of heating a sulfide ore to a high temperature in the presence of air. It is a step in the processing of certain ores. More specifically, roasting is often a metallurgical process involving gas–solid reactions at elevated temperatures with the goal of purifying the metal component(s). Often before roasting, the ore has already been partially purified, e.g. by froth flotation. The concentrate is mixed with other materials to facilitate the process. The technology is useful in making certain ores usable but it can also be a serious source of air pollution.[1]
Roasting consists of thermal gas–solid reactions, which can include oxidation, reduction, chlorination, sulfation, and pyrohydrolysis. In roasting, the ore or ore concentrate is treated with very hot air. This process is generally applied to sulfide minerals. During roasting, the sulfide is converted to an oxide, and sulfur is released as sulfur dioxide, a gas. For the ores Cu2S (chalcocite) and ZnS (sphalerite), balanced equations for the roasting are:
- 2 Cu2S + 3 O2 → 2 Cu2O + 2 SO2
- 2 ZnS + 3 O2 → 2 ZnO + 2 SO2
The gaseous product of sulfide roasting, sulfur dioxide (SO2) is often used to produce sulfuric acid. Many sulfide minerals contain other components such as arsenic that are released into the environment.
Up until the early 20th century, roasting was started by burning wood on top of ore. This would raise the temperature of the ore to the point where its sulfur content would become its source of fuel, and the roasting process could continue without external fuel sources. Early sulfide roasting was practiced in this manner in "open hearth" roasters, which were manually stirred (a practice called "rabbling") using rake-like tools to expose unroasted ore to oxygen as the reaction proceeded.
This process released large amounts of acidic, metallic, and other toxic compounds. Results of this include areas that even after 60–80 years are still largely lifeless, often exactly corresponding to the area of the roast bed, some of which are hundreds of metres wide by kilometres long. Roasting is an exothermic process.[2][3]
https://en.wikipedia.org/wiki/Roasting_(metallurgy)
Roasting operations
The following describe different forms of roasting:[4]
Oxidizing roasting
Oxidizing roasting, the most commonly practiced roasting process, involves heating the ore in excess of air or oxygen, to burn out or replace the impurity element, generally sulfur, partly or completely by oxygen. For sulfide roasting, the general reaction can be given by:
- 2MS (s) + 3O2 (g) -> 2MO (s) + 2SO2 (g)
Roasting the sulfide ore, until almost complete removal of the sulfur from the ore, results in a dead roast.[5]
Volatilizing roasting
Volatilizing roasting, involves oxidation at elevated temperatures of the ores, to eliminate impurity elements in the form of their volatile oxides. Examples of such volatile oxides include As2O3, Sb2O3, ZnO and sulfur oxides. Careful control of the oxygen content in the roaster is necessary, as excessive oxidation can form non-volatile oxides.
Chloridizing roasting
Chloridizing roasting transforms certain metal compounds to chlorides through oxidation or reduction. Some metals such as uranium, titanium, beryllium and some rare earths are processed in their chloride form. Certain forms of chloridizing roasting may be represented by the overall reactions:
The first reaction represents the chlorination of a sulfide ore involving an exothermic reaction. The second reaction involving an oxide ore is facilitated by addition of elemental sulfur. Carbonate ores react in a similar manner as the oxide ore, after decomposing to their oxide form at high temperatures.[6]
Sulfating roasting
Sulfating roasting oxidizes certain sulfide ores to sulfates in a supply of air to enable leaching of the sulfate for further processing.[citation needed]
Magnetic roasting
Magnetic roasting involves controlled roasting of the ore to convert it into a magnetic form, thus enabling easy separation and processing in subsequent steps. For example, controlled reduction of haematite (non magnetic Fe2O3) to magnetite (magnetic Fe3O4).
Reduction roasting
Reduction roasting partially reduces an oxide ore before the actual smelting process.
Sinter roasting
Sinter roasting involves heating the fine ores at high temperatures, where simultaneous oxidation and agglomeration of the ores take place. For example, lead sulfide ores are subjected to sinter roasting in a continuous process after froth flotation to convert the fine ores to workable agglomerates for further smelting operations.
References
- Ghosh, Ahindra; Ray, Hem Shanker (1991). Principles of Extractive Metallurgy. ISBN 9788122403220.
https://en.wikipedia.org/wiki/Roasting_(metallurgy)
https://en.wikipedia.org/wiki/Refractory
https://en.wikipedia.org/wiki/Flash_smelting
https://en.wikipedia.org/wiki/Magnetation_(iron_ore)
https://en.wikipedia.org/wiki/Magnetic_separation
https://en.wikipedia.org/wiki/Base_metal
A base metal is a common and inexpensive metal, as opposed to a precious metal such as gold or silver.[1] In numismatics, coins often derived their value from the precious metal content; however, base metals have also been used in coins in the past and today.[2]
Specific definitions
In contrast to noble metals, base metals may be distinguished by oxidizing or corroding relatively easily and reacting variably with diluted hydrochloric acid (HCl) to form hydrogen. Examples include iron, nickel, lead and zinc. Copper is also considered a base metal because it oxidizes relatively easily, although it does not react with HCl.
In mining and economics, the term base metals refers to industrial non-ferrous metals excluding precious metals. These include copper, lead, nickel and zinc.[3]
The U.S. Customs and Border Protection agency is more inclusive in its definition of commercial base metals. Its list includes—in addition to copper, lead, nickel, and zinc—the following metals: iron and steel, aluminium, tin, tungsten, molybdenum, tantalum, cobalt, bismuth, cadmium, titanium, zirconium, antimony, manganese, beryllium, chromium, germanium, vanadium, gallium, hafnium, indium, niobium, rhenium, and thallium, and their alloys.[4]
https://en.wikipedia.org/wiki/Base_metal
The philosopher's stone, or more properly philosophers' stone (Arabic: حجر الفلاسفة, romanized: ḥajar al-falāsifa; Latin: lapis philosophorum), is a mythic alchemical substance capable of turning base metals such as mercury into gold (chrysopoeia, from the Greek χρυσός khrusos, "gold", and ποιεῖν poiēin, "to make") or silver. It is also called the elixir of life, useful for rejuvenation and for achieving immortality;[1] for many centuries, it was the most sought-after goal in alchemy. The philosopher's stone was the central symbol of the mystical terminology of alchemy, symbolizing perfection at its finest, enlightenment, and heavenly bliss. Efforts to discover the philosopher's stone were known as the Magnum Opus ("Great Work").[2]
https://en.wikipedia.org/wiki/Philosopher%27s_stone
Immortality is the concept of eternal life.[2][failed verification] Some modern species may possess biological immortality.[citation needed]
Some scientists, futurists, and philosophers have theorized about the immortality of the human body, with some suggesting that human immortality may be achievable in the first few decades of the 21st century with the help of certain technologies such as mind uploading (digital immortality).[3] Other advocates believe that life extension is a more achievable goal in the short term, with immortality awaiting further research breakthroughs. The absence of aging would provide humans with biological immortality, but not invulnerability to death by disease or injury. Whether the process of internal immortality is delivered within the upcoming years depends chiefly on research (and in neuron research in the case of internal immortality through an immortalized cell line) in the former view and perhaps is an awaited goal in the latter case.[4]
What form an unending human life would take, or whether an immaterial soul exists and possesses immortality, has been a major point of focus of religion,[citation needed] as well as the subject of speculation and debate. In religious contexts, immortality is often stated to be one of the promises of divinities to human beings who perform virtue or follow divine law.[citation needed]
https://en.wikipedia.org/wiki/Immortality
Biological immortality (sometimes referred to as bio-indefinite mortality) is a state in which the rate of mortality from senescence is stable or decreasing, thus decoupling it from chronological age. Various unicellular and multicellular species, including some vertebrates, achieve this state either throughout their existence or after living long enough. A biologically immortal living being can still die from means other than senescence, such as through injury, poison, disease, predation, lack of available resources, or changes to environment.
This definition of immortality has been challenged in the Handbook of the Biology of Aging,[1] because the increase in rate of mortality as a function of chronological age may be negligible at extremely old ages, an idea referred to as the late-life mortality plateau. The rate of mortality may cease to increase in old age, but in most cases that rate is typically very high.[2]
The term is also used by biologists to describe cells that are not subject to the Hayflick limit on how many times they can divide.
https://en.wikipedia.org/wiki/Biological_immortality
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