“All life is chemistry.” ~ Belgian chemist Jan van Helmon
“Organic chemistry is the chemistry of carbon compounds. Biochemistry is the study of carbon compounds that crawl.” ~ Mike Adam
Life sustains itself by chemical energy. Transformative sustenance transpires in water.
Clean water is vital for all known life. Mars may have never evolved life because the salinity of its water has long exceeded levels by which life could arise, survive, or thrive.
Humans are 65–90% water, scaffolded by carbon-containing organic molecules. 99% of human body mass comprises 6 elements: oxygen (65%), carbon (18%), hydrogen (10%), nitrogen (3%), calcium (1.4%), and phosphorous (1.1%). By contrast, aluminum and silicon are abundant in Earth’s crust but rare in life forms.
Carbon’s unsurpassed flexibility lends itself to complexity. Sturdy nitrogen provides stability. Oxygen and hydrogen readily react, and so make excellent elements for cellular activity. Being readily reactive, calcium ions play a central role in many cellular functions, and in intercellular communication. As a store of energy, phosphorous is vital to metabolism. 2 more elements are noteworthy for their organic roles: potassium and sulfur. Potassium (0.4%) assists in homeostasis and cellular communication. Through its surfeit of electrons which willingly shuffle, sulfur (0.3%) helps catalyze reactions.
All told, cells make a living using just 30 different monomers (molecule types).
“The characteristics of a cell rest on the structure of its molecular components.” ~ French biologist François Jacob
“The molecular components of organisms are remarkably uniform in the nature of the components, as well as in the ways in which they are assembled and used.” ~ Spanish American biologist Francisco Ayala
A functional group comprises slightly different configurations of the same organic compound, all with similar behaviors. A functional group is typically a molecular subset of a larger molecule. The same functional group produces selfsame or similar reactions regardless of the molecule of which it is part.
An organic compound’s functional group is indicative of the entire molecule’s reactive properties. Reactions of an organic compound can be predicted by knowing the kind(s) of functional group(s) it has.
Carbon molecules are commonly classified based upon their functional group. Proteins, the building blocks of life, are complex carbon-based macromolecules comprising many functional groups. In doing their job, proteins selectively employ a certain functional group as a tool.
The term moiety is often used as a synonym for functional group, but a moiety is a specific segment of a molecule, regardless of functional group. A moiety may be part of a functional group or encompass parts of different functional groups.
“The meaning of the “p” in “pH” is unknown.” ~ Wikipedia
pH is a measure of how base or acidic an aqueous solution is. In a chemical reaction, an acid is a molecule or ion capable of donating a cation of protium (1H+); a base is accepting of 1H+.
(The term proton is sometimes used for 1H+. This usage comes from the 1923 concept of acids and bases independently developed by chemists Johannes Brønsted and Martin Lowry. Of course, using proton here is inexact. The term hydron is also sometimes used, though generally disfavored, as hydron is the general name for H+, which is the cationic form of hydrogen, regardless of isotope. Hence, hydron represents the nucleus of 1H+, 2H+ (D+) (deuterium isotope), and 3H+ (T+) (tritium isotope). It may seem strange that chemistry has no exact term for 1H+. There have been numerous characterizations of acids and bases, including a 1923 electronic theory of acid-base reactions by Gilbert Lewis. The Brønsted–Lowry conceptualization is a theory for which exceptions are known. Much of the intricacy of chemistry remains to be discovered.)
Water is neutral (pH = 7). Acids have a pH < 7, while bases have a >7 pH.
pH was coined in 1924. The “H” in pH stands for hydrogen, but the meaning of “p” is a long-standing mystery. p = power is a common guess. pH: the power of hydrogen.
While acidity is related to the concentration of hydrogen ions, it is not ionic concentration per se that confers pH, but instead the activity factor of a solution, which is the tendency of hydrogen ions to interact with other elements in a solution. Acid-base reactions are tangled enterprises.
pH is a major aspect of organic chemistry. Acids or bases promote most organic reactions or are at least involved at some stage in a reaction pathway.
After the needed action is accomplished, pH hyperactivity needs calming down; in a word: neutralized.
Acids and bases are ionic go-getters. It would make sense that neutralizing them has something to do with taking a charge off: deionization (the removal of ions). The ease of neutralization is related to the electronegativity of the elements involved. But something much fatter than electrons easing up is involved. A proton jumps between ions to neutralize. In aqueous neutralization involving an acid and salt, the salt is but a spectator. Proton hops turn acidic ions (H3O+) into water (H2O).
More generally, proton transfer either turns an acidic H3O+ (hydronium) ion into a water molecule or ionizes a water molecule to the base HO– (hydroxide). Transfer takes 1–2 picoseconds, with the proton hopping along a particular pathway.
The local H-bond water network provides a pathway for proton transfer; or not. Without a permissive aqueous dynamic and structure, proton transfer is stymied. A proton may hop its way to a dead end, where it is locally trapped for an extended period.
In a gas reaction, ions clump by attraction. Without any solvent at all, proton transfers neutralize.
Acids and bases mixed in the right proportions neutralize by proton transfer: acids donate protons that bases accept.
“Life exists in the universe only because the carbon atom possesses certain exceptional properties.” ~ English physicist James Jeans
Carbon (C) is the only element that can form chains and rings on its own, using single or double bonds. Thus, carbon acts as a backbone in manifold molecular construction.
4 of carbon’s 6 electrons are available to form strong covalent bonds. This uniquely flagrant friendliness makes carbon the most flexible and stable element for complex compound construction. Hence, life is built on carbon. Carbon provides the chemical skeleton of the 4 major types of organic compounds: nucleic acids (aka polynucleotides (RNA, DNA)), carbohydrates (sugars), proteins, and lipids (fats).
An organic compound, which necessarily has a carbon backbone, is an assemblage (polymer) of smaller molecular subunits (monomers), altogether forming a macromolecule. Polymerization is the process by which repeating subunits (monomers) are bound into chains of various lengths, forming polymers. Amino acids (monomers) are polymerized into proteins.
Coupled to their energetically ready construction, the complexity and flexibility of macromolecules lets them work in a wide variety of roles: functioning as structural components, molecular messengers, enzymes, energy sources, nutrient storage, and storehouses of genetic information.
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While carbon normally bonds with 4 other atoms, other configurations are possible. Mellitene (C12H18) is based upon a hexagonal ring of 6 carbon atoms offering 6 arms for 6 carbon hydrogen atoms that each have 3 hydrogen atoms attached. Leftover electrons zip around the middle of the ring, strengthening core carbon bonds and thereby stabilizing the molecule.
Pure carbon in crystalline form appears as diamonds: a lattice allotrope with remarkable optical characteristics, and the hardest natural material known.
An allotrope is an element existing in multiple forms, with atoms bonded together in different ways. The structural variations characterizing allotropy apply only to distinct forms of an element within the same phase, and only to elements, not compounds.
A more pedestrian carbon allotrope than mellitene is graphite, the stuff of pencil lead. While diamond won’t conduct electricity, graphite will. Yet graphite is the most stable form of carbon.
Because of carbon’s inherent promiscuity, there is a greater variety of carbon compounds than all other elements combined. Plastics and petrol, which are derived from erstwhile plant matter, are carbon polymers.
Atmospheric carbon most naturally occurs as carbon dioxide (CO2). While carbon dioxide is currently a relatively minuscule component of the atmosphere, at about 0.039%, Earth’s early atmosphere was rich in CO2.
CO2 is transparent to visible light, as are N2 and O2, the other notable atmospheric gases. But whereas nitrogen and oxygen are also insensitive in the infrared spectrum, carbon dioxide absorbs infrared light.
In trapping infrared, CO2 prevents some of that radiation from escaping into space. Hence atmospheric carbon dioxide acts as a greenhouse gas, helping keep the planet below warm.
Water vapor has a powerful greenhouse effect, absorbing light in a broader spectrum than carbon dioxide. H2O absorbs essentially everything at wavelengths longer than 500 cm–1. CO2 centers its absorption around 667 cm–1.
Absorption dynamics depend upon quantum mechanical properties specific to a molecule’s geometry. Radiation at specific frequency ranges jiggles the bond angles of a molecule, affecting molecular vibrational mode, allowing a molecule to hold on to extra energy.
It was the carbon in carbon dioxide, not the oxygen, that was the ready fuel for emergent life. The magic trick was transforming the carbon in CO2 into energy.
“Photosynthesis powers life on our planet.” ~ American physicist Franklin Fuller
The chemical equation for photosynthesis is: 6 CO2 +
12 H2O + sunlight energy → C6H12O6 + 6 H2O + 6 O2.
A photosynthesizing organism derives energy from sunlight via charge separation; forging atmospheric carbon dioxide and water into energy-storing sugars by freeing electrons to work chemical transformations via photonic energy. Charge separation is the process of an atomic electron being excited to a higher energy level by absorbing a photon, and thereby leaving its home to join a nearby electron acceptor molecule.
Chloroplasts – plant organelles that conduct photosynthesis – have chlorophyll that absorb various wavelengths of light, with peak inputs around 430 nanometers (blue light) and 660 nm (red). Chlorophyll is green because it shies from absorbing mid-spectrum green light (526–606 nm), which is thereby reflected.
O2 is a byproduct of photosynthesis; an oddity considering that oxygen is itself energy-rich by its easy reactivity. But throwing off water’s O2 makes sense in light of the photosynthesis system.
In plant photosynthesis, oxygen stubbornly sticks to the carbon in CO2, which becomes the backbone molecule upon which the target sugar (glucose: C6H12O6) is built. In the same reaction as evaporation, water molecules are split to liberate oxygen (O2) into the atmosphere.
The rate of charge separation for evaporation in photosynthesis is enhanced by coherent dynamics that are vibronic in nature. This allows photosynthesis to be optimal at the quantum level.
Hydrogen is the watery element elicited in photosynthesis; valued for its donation in the electron transport chain: the molecular processing system for deriving energy.
Carbon has a dual role in photosynthesis. It is the proton acceptor (HCO3–) in the very beginning of the photosynthetic reaction, and the terminal electron acceptor (CO2) at the end. This varied employment of different species in a reaction series illustrates the wondrous flexibility of carbon.
A hydrocarbon is any variety of organic compound consisting entirely of carbon and hydrogen. Owing to the inherent molecular flexibilities of carbon and hydrogen, hydrocarbons assume a vast variety of structures.
Saturated hydrocarbons (alkanes) are the simplest hydrocarbon species, composed entirely of single bonds. Petroleum fuels are derived from alkanes. Alkanes with 1 or more rings of carbon atoms are cycloalkanes.
Unsaturated hydrocarbons have 1 or more double or triple bonds between carbon atoms. Alkenes are double-bonded hydrocarbons. Hydrocarbons with triple bonds are alkynes.
Hydrocarbons come in a variety of phases at room temperature: gaseous (e.g. methane (CH4)), liquid (e.g. hexane (C6H14) and octane (C8H18)), wax or low-melting solid (e.g. paraffin (CnH2n+2, from C20H42 to C40H82)), and polymers (e.g. the ubiquitous plastic polyethylene (always following the formula: (C2H4)nH2), which is so durable as to resist biodegradation).
The more complex the hydrocarbon, the more stable.
Hydrocarbons are hydrophobic, as are lipids.
Lipids are a broad molecular group, engorged with nomenclature. The lipid group includes fats, waxes, sterols, glycerides, phospholipids, and fat-soluble vitamins.
Although the terms lipid and fat are sometimes used as synonyms, fats are a lipid subgroup.
All fats are derivatives of glycerol, combined with fatty acids. Glycerol is a simple alcohol (polyol) compound, comprising 3 hydroxyl groups. A hydroxyl is a functional group comprising an oxygen atom covalently bonded to a single hydrogen atom. Fatty acids comprise a carboxylic acid with a long aliphatic tail (chain). An aliphatic compound is a group of hydrocarbons that do not link together to form a ring.
Triglycerides – the common form of fat in animals – have 3 fatty acids. Humans store unused energy (calories) in triglycerides. Overeating high-carbohydrate foods packs on saturated triglycerides.
Chemical composition complexities aside, the biologically relevant facts of fats are that they act as an energy store, and that fats come in 2 forms: saturated and unsaturated.
A saturated fat has only single bonds between carbon atoms. The term saturated is used because the fat is chock full of hydrogen atoms.
In contrast, unsaturated fat has at least 1 double bond within the fatty acid chain. A fat molecule with only 1 double bond is monounsaturated. Molecules of fat with more than 1 double bond are polyunsaturated.
Hydrogen atoms are eliminated where double bonds are formed. Thus, metabolically, unsaturated fats hold a bit less energy (i.e., fewer calories) than an equivalent portion of saturated fat.
Unsaturated fats are regarded as healthier for human consumption; the more unsaturated, the better. Unsaturated fats metabolize more cleanly.
Nature equipped biological systems to recognize the shapes and other characteristics of molecules and process them accordingly. Metabolically, a saturated fat is treated to a different set of biochemical reactions than an unsaturated fat.
Each form of fat has a different shape. Fats are metabolized by geometry.
The more double bonds there are, the more readily a fat will react with oxygen. The tendency to rancidity is greater the more unsaturated the fat is. So, to preserve shelf life, fatty foods are commercially processed by hydrogenation: saturating the fat at high temperature, thus breaking the double bonds and slimming them to single bonds, with hydrogen attached.
“You will die but the carbon will not; its career does not end with you. It will return to the soil, and there a plant may take it up again in time, sending it once more on a cycle of plant and animal life.” ~ Polish polymath Jacob Bronowski
The carbon cycle is the gaseous cycling of carbon exchange among the pedosphere (soil), hydrosphere (water bodies), atmosphere, and biosphere (living ecological systems) of the Earth. Carbon deposits and exchanges among the different spheres differ.
The atmosphere has a small active exchange of CO2. The pedosphere is both a carbon sink and primary cycler of carbon.
Autotrophs, predominantly plants, move carbon through the cycle, fixing CO2 into organic compounds. These organic compounds are consumed by hungry heterotrophs (e.g. animals), and then decompose into a carbon sink. Carbon is returned to the atmosphere as CO2 by heterotrophic cellular respiration, such as by breathing.
Recesses below ground and in the deep ocean are vast carbon storehouses, with exchange becoming active only by violent disturbances: volcanic eruptions and larger-scale geological movements (metamorphism), such as plate tectonics.
Long before life emerged, carbon was cycled from the atmosphere into the geosphere as carbonate sediments. The evolution of photosynthesis accelerated the carbon cycle, which gained an even greater boost as plants colonized land.
Organic carbon from plant matter became buried as coal and in marine sediments, resulting in a dramatic decline in atmospheric CO2, while plant exhaust (oxygen) became prominent.
The interplay of carbon and oxygen have been a central driver in the emergence and development of life on Earth, and their cycles remain pivotal in the biological fortune of the planet.
Between 24,000 years ago and today, large ice sheets covering most of Canada, and parts of Europe and Asia, melted away. Global sea level rose by 120 meters. Earth warmed 5 °C. Rainfall and vegetation patterns shifted, sometimes abruptly.
Milankovitch cycling, including orbital shifts, set these changes into motion. But a complex feedback system, in which the carbon cycle was integral, governed the transition from glacial to interglacial state.
17,500 ya, CO2 levels started to rise from the ice age level of 180 parts per million (ppm), reaching 265 ppm 10,000 ya. Over the next 10,000 years, atmospheric carbon dioxide rose another 20 ppm, until the rapid ramp originating with the onset of the industrial revolution in the 1850s.
In 2010, atmospheric CO2 was 380 ppm. It rose to 405 ppm in 2017. By 2080, atmospheric CO2 should be over 700 ppm.
In the Ocean
The ocean has the largest store of actively exchanged carbon. The seas near the surface constantly exchange carbon with the atmosphere. The depths are not so yielding.
The deep ocean holds 50 times more CO2 than the atmosphere. It does so by a biological pump.
Phytoplankton fix CO2 at the sea surface by photosynthesis, forming organic carbon. When they die, the phytoplankton sink. The dissolved organic matter (DOM) – including carbon – is consumed by other organisms, and respired back into CO2, which forms dissolved carbonate.
The biological pump also depletes nutrients such as nitrogen and silica from the surface waters. Nitrogen is needed by all phytoplankton, while silica is in demand by diatoms. All these nutrients become trapped in the deep sea by the shallow surface layer.
Ocean circulation returns this nutrient-rich water to the surface in the Southern Ocean, and at upwelling regions. When this happens, biological productivity at the surface is enhanced, and CO2 is released into the atmosphere, if degassing outpaces the biological pump.
Seagrass meadows are concentrated stores of carbon dioxide. A typical forest stores 30,000 tonnes of CO2 per square kilometer. Coastal seagrass beds store up to 83,000 tonnes per km2.
Although seagrass meadows are less than 0.2% of the world’s oceans, they hold more than 10% of all carbon buried annually in the sea. 90% of the carbon storage is in the soil, which accumulates as the meadow thrives.
Seagrass meadows can last for centuries; but not any longer. Seagrass beds are among the most threatened ecosystem worldwide, thanks to dredging and degrading water quality.
29% of all seagrass meadows have already been destroyed, with further yearly loss of at least 1.5% of the historic total. CO2 emissions for seagrass meadow destruction could be as much as 25% of that which comes from deforestation.
‘Dryland systems have high rates of carbon turnover compared to other biomes.” ~ French ecologist Philippe Ciais
Semi-arid biomes, from deserts to shrublands, drive variability in the carbon cycle. The greening of these lands from sporadic or seasonal rainfall has an outsized effect on carbon cycle dynamics. As the climate warms, drylands become particularly sensitive to precipitation.
Forests comprise the great storehouse of terrestrial carbon: 86% of the aboveground carbon. Forest soil has 73% of the Earth’s soil carbon.
Forests are extremely active in the carbon cycle. Forests store massive quantities of carbon in the soil, through the root systems of trees and the microbes which support vegetation in mutualistic relationships.
The quality of the soil determines how much carbon is retained. Fertile soils sequester 30% of the carbon they take in during photosynthesis. In contrast, nutrient-poor soils may only retain 6% of that carbon.
Logging accelerates carbon release into the atmosphere in multiple ways. Besides killing the trees that act as carbon keepers, logging cuts off the food supply for subterranean life in the soil. Deforestation bankrupts the prospect of the ground acting as a carbon store.
The organic matter in soil is 60% carbon. If the withdrawal of carbon from the soil went up by just 0.3%, the release into the atmosphere would equal a year’s worth of emissions worldwide from human use of fossil fuels.
Scientists long thought that the soil locked in carbon; a comforting fiction. The supposedly stable carbon keepers in soil are instead relatively volatile.
For over a century, it was supposed that soil carbon became locked in large compound molecules, highly resistant to microbial assault. Only recently has it been discovered that those presumed molecules – so-called humic soil substances – simply don’t exist in any significant quantity. Instead, the persistence of organic matter in soil is controlled by the surrounding microbial ecosystem.
Plants thrive on CO2. The more there is in the air, the faster and larger plants can grow, if other conditions conducive to growth exists (such as quality soil and proper rainfall).
Plants, especially ones with extensive roots, such as trees, have steadfast friends in the soil: symbiotic microbes. Soil microbial ecosystems rely upon plant life. Plant loss releases CO2, as soil microbes lose their livelihoods, and their lives.
The carbon cycle is also invigorated by the soil in forests. As trees thrive from more CO2, so too soil life, so much so that the microbes start consuming and releasing the carbon long locked up in the accumulated plant matter from generations long past.
At 700 ppm CO2, trees are pumping a bit more oxygen into the air is grossly outweighed by soil carbon release, which is as much as 15% at the surface. That’s 50 years of worldwide fossil fuel burning at the current rate.
While the outlines are clear, the carbon cycle is not well understood at the detail level. This is because isotopic variations of CO2 have distinct significances in the carbon cycle. Different carbon cycle processes affect different carbon dioxide isotopes, and the different isotopes have varying impacts on global warming.
In 2010 there were 1 billion vehicles on the road. 300 million more were added by 2018. The number of vehicles is projected to rise to at least 3 billion by 2050. That will put air pollution, and atmospheric CO2 release, into overdrive.
It is a safe bet that deforestation worldwide will continue apace if not accelerate. With more cars doing their part to increase atmospheric CO2, and fewer trees, the forest carbon bank will be paying its interest in a positive feedback cycle to accelerate atmospheric CO2 well beyond current projections, which do not take into account the actual dynamics of soil in the carbon cycle. Human mass intervention in the carbon cycle is clearly an experiment of global consequence.
“All we are is a lot of talking nitrogen.” ~ American playwright Arthur Miller (We are also a lot of farting nitrogen. Nitrogen is the main constituent of flatus (20–90%), followed by carbon dioxide (10–30%).)
Scottish physician Daniel Rutherford is credited with discovering nitrogen in 1772, though he did not identify it. He simply killed a mouse to show that it couldn’t breathe once the ambient oxygen had been used up.
Rutherford’s discovery was the leftover gas besides oxygen: mostly nitrogen, with residual carbon dioxide; what he called noxious air (fixed air). Rutherford’s noxious air provided him further proof of his conviction to phlogiston theory.
Imaginative German alchemist Johann Joachim Becher concocted the phlogiston theory in 1667. By his account, phlogiston was an odorless, tasteless, colorless, massless fire element, contained within combustible substances, and released during combustion.
Once burned, a dephlogisticated substance was held to be in its ‘true’ form, the calx. Phlogiston theory purported to explain both combustion and the rusting of metals, processes now collectively known as oxidation.
If carbon is the king of organic chemistry, nitrogen is queen. Nitrogen (N) is an essential element in building amino acids and nucleic acids, which are respectively the building blocks of proteins and genomes. Nitrogen is the 4th most abundant element in organisms, and the 7th most common in the cosmos.
While essential and abundant, nitrogen only reluctantly plays its organic role. Nitrogen has 5 electrons in its outer shell, and so is trivalent in most compounds. The triple bond of molecular nitrogen is a tough bond to break.
For nitrogen to become organically employed, it must be “fixed.” Nitrogen fixation is the ability of an organism to transform atmospheric N2 into usable nitrogen species, such as ammonia (NH3). Only prokaryotes, supposedly simple soil bacteria, are capable of nitrogen fixation.
Any supposition of simplicity is deceptive. Wholesale nitrogen fixation occurs only after bacteria become intimate with a plant by a welcomed invasion. The process is the most complex coordination between 2 species that is known.
Diazotroph bacteria mastered the knack of fixing nitrogen thanks to nitrogenase, an enzyme that assists in the necessary chemical reactions. Some higher plants and termites joined with diazotrophs in symbiotic relationships, providing a home in return for the ability to fix nitrogen.
The primary role of plants in the nitrogen cycle is assimilation of biologically-accessible nitrogen (NO–3 & NO+4) into plant biomass, which is locked up until the plant dies.
Protists and fungi incorporate organic nitrogen into their biomass, but do not actively excrete nitrogen compounds; but animals do.
The primary nitrogen excretory product from animals is ammonia (NH3), derived from the digestion of proteins. Because ammonia is highly toxic, the urinary and excretory systems of animals are complex regulatory mechanisms that afford quick and efficient elimination.
All known life needs fixed nitrogen. Because nitrogen is physiologically essential, organisms evolved various mechanisms for regulating its uptake and excretion. These mechanisms vary widely in different life forms.
“In many ecosystems worldwide, nitrogen is the element whose supply rate from the environment is most limited. Because competition is fierce for this resource, nitrogen supply controls the behaviour of many organisms and shapes the structure and function of whole ecosystems.” ~ American ecologist Edward Schuur
The nitrogen cycle is the gaseous cycling of nitrogen between the environment and life, from elemental to biologically accessible. Much of the nitrogen that organisms need to grow is supplied by recycling. New deposits of nitrogen are minuscule compared to quantities recycled. But nitrogen infusion is vital for newly forming ecosystems, and for balancing natural nitrogen loss from an ecosystem into streams or back into the atmosphere.
Nitrogen gas (N2) comprises 78% of the atmosphere, and so the air forms a large reservoir pool. The soil makes a small exchange pool.
This is quite a contrast to the carbon cycle, where the atmosphere is the exchange pool, and the soil a sink. What is much the same between the 2 cycles is the prominent part played by plants, which, in collaboration with soil bacteria, fix nitrogen into an organically usable form.
On land, nitrogen cycling comes in the organic back end: from the decomposition of organic matter. Denitrifying bacteria convert soil nitrates into gaseous N2.
The largest pool of fixed nitrogen is dissolved in the oceans. Dissolved organic nitrogen (DON) accounts for 60% of the reactive nitrogen in the ocean.
DON fuels formation of organic molecules from carbon dioxide, especially in oceanic regions low in inorganic nutrients, such as open-ocean gyres: large systems of rotating currents. Most DON is cycled by phytoplankton in less than a year.
Marine sedimentary rock is a rich source of nitrogen: 10 times that found in igneous rock, which explains why plants have such a tough time establishing themselves after a volcanic eruption. 75% of Earth’s crust is covered by sedimentary and related rock types.
The nitrogen cycle is naturally a well-orchestrated system, balanced on a planetary scale. For Earth as a whole, the nitrogen fixed equals the nitrogen returned to the atmosphere. If this were not true, atmospheric nitrogen would become depleted.
It is crucial for healthy ecosystems that a balanced nitrogen cycle be maintained. Too much new nitrogen introduced into the cycle – such as excessive application of nitrogen-based fertilizers and other pollution that releases nitrogen into the atmosphere – and world ecology is profoundly affected. This has been happening at least since the onset of the Industrial Revolution in the mid-19th century, beginning with widespread industrial burning, the subsequent commercialization of internal combustion engines, and capped off with nitrogen-rich fertilizers.
Nitrogen fertilizer to promote plant growth has increased crop yields while severely polluting the environment with excess nitrogen. Because of the inefficiencies of nitrogen uptake, only about 10% to 15% of the nitrogen applied via fertilizer is used by crops. The rest is released into the environment as pollution. Burning fossil fuels – the fossils being those of plants – also releases nitrogen.
Excess nitrogen from human intervention has altered the Earth’s nitrogen cycle in the soil, oceans, and atmosphere. Atmospheric depositing of reactive nitrogen into terrestrial ecosystems globally doubled from 1968 to 2007. The pace is accelerating.
Excess reactive nitrogen in water depletes oxygen supply by redox reactions. Runoff from fertilizer into bodies of water has created a growing number of “dead zones” around the world. Only primitive anaerobic life, which needs no oxygen, can survive in these aquatic dead zones.
As a result of fertilizer runoff, half of the lakes in the United States are eutrophic: have excess nutrients that result in poor water quality, including turbidity and oxygen depletion, from the bottom of the water body on up. Eutrophic water is conducive to algal blooms, which further reduce the oxygen in the water.
“During peace time, a scientist belongs to the world, but during war time, he belongs to his country.” ~ Fritz Haber
German chemist Fritz Haber invented the Haber process in 1909, allowing nitrogen fixation by reaction of nitrogen gas and hydrogen gas, catalyzed by enriched iron or ruthenium; a process used industrially to produce ammonia. The Haber process was an important step for the industrial production of both fertilizers and explosives.
For his role in creating chemical weapons in World War I, Haber is considered the father of chemical warfare: developing and deploying chlorine and other poisonous gases on the enemy. Haber personally supervised the first gas attacks. The Kaiser promoted him to captain; a rare gift for a scientist too old to enlist.
Haber’s wife of almost 2 decades, also a chemist, committed suicide over his enthusiasm for his work. His son would later commit suicide, in shame over his father’s accomplishments.
To further his career prospects, Haber, born into a Hasidic family, renounced Judaism and became a Lutheran.
It did him no good when the Nazis came to power. Haber fled Germany in 1933 but left a legacy. Haber developed the gas Zyklon B, which was later used to exterminate Jews in the Nazi death camps.
Oxygen (O) is a colorless, odorless gas, though molecular oxygen appears pale blue. Like nitrogen, oxygen’s most common molecular form is diatomic (O2). But unlike sturdy nitrogen, oxygen is highly reactive: forming oxide compounds with almost all other elements. The electron orbitals of ground state dioxygen (O2) have 2 degenerates: 2 unpaired electrons with antibonding orbitals. The O–O molecular bond is weaker than the N–N molecular bond.
Diradicals are molecular species with 2 electrons occupying 2 equal energy (degenerate) orbits. Diradicals vary by the spins of covalently bonded electrons. O2 and H2C (methylene) are exemplary diradicals.
Diradicals have 3 possible arrangements: singlet, doublet, and triplet.
The oxygen molecule’s ground state is triplet. Both electrons in the degenerate orbitals spin up. The parallel spins of the unpaired electrons make gaseous dioxygen paramagnetic (reactive to magnetic fields), which is quite unusual for a gas. Liquid oxygen is magnetic.
Singlet oxygen has 1 spin-up electron and 1 spin-down electron in 1 orbital, with an equal energy orbital empty. The singlet configuration has several species, all relatively high energy. Hence singlet O2 is much more reactive than triplet O2.
If the O2‘s ground state was singlet instead of triplet, life would be impossible, and any accumulation of organic matter unlikely. Singlet oxygen is used as an industrial-strength pesticide in buildings, exterminating even the most persistent critters.
The human immune system produces singlet O2 for weaponry: reactive oxygen species (ROS). Plants have an ROS response to attacking pathogens: strengthening the cell wall with superoxide or hydrogen peroxide to imprison the infection.
In a reaction powered by sunshine, singlet oxygen is formed from water during photosynthesis. Carotenoids in chloroplasts absorb energy from singlet oxygen using tetraterpenoids: a molecular skeletal structure comprising 40 carbon atoms. Carotenoids convert the highly reactive singlet to triplet ground state before it inflicts harm on tissues.
The ability to detoxify ROS evolved in the earliest life, prior to photosynthesis and aerobic respiration; otherwise, aerobic organisms would have poisoned themselves.
Photolysis of ozone by short-wavelength (high-energy) light in the troposphere produces singlet O2. The troposphere is the lowest layer of the Earth’s atmosphere, 17 kilometers thick in the middle latitudes, with about 75% of the atmosphere’s mass and 99% of its water vapor and aerosols.
Doublet O2, with one electron unpaired, is a simple free radical, and highly reactive.
O2‘s odd triplet ground state prevents molecular oxygen from reacting directly with many other molecules, which are often in the singlet state. But triplet oxygen will readily react with doublet molecules, such as radicals, to form a new radical. And the univalent pathway of adding electrons 1 at a time in series reactions is quite common.
Oxygen’s ready reactivity explains its toxicity, by too-enthusiastic production of intermediates, especially singlet oxygen, and, from water, hydrogen peroxide (H2O2) and the hydroxyl radical (*OH). Free-radical oxygen also enhances the lethality of ionizing radiation.
There are biochemical defenses against oxygen toxicity. Antioxidants are a rear-guard to minimize tissue damage. A free-radical chain reaction triggered by singlet O2 can be broken by antioxidants that react with the chain-propagating radicals.
Vitamin E (α-tocopherol) is one antioxidant biocompound. Vitamin E deficiency can cause muscular dystrophy and reproductive failure in humans.
Chronic, low-level, cumulative oxygen toxicity is the biochemical cause of aging.
Discovery of Oxygen
English theologian Joseph Priestley is often credited with discovering oxygen, though claims can be made for Antoine Lavoisier and Carl Wilhelm Scheele.
Priestley gained a scientific reputation for inventing soda water. He also wrote on electricity, and discovered several airs, including “dephlogisticated air,” now known as oxygen. Priestley was a staunch adherent of Johann Joachim Becher’s phlogiston theory: that combustible matter contained a hidden fire element, phlogiston.
Priestley’s attribution of oxygen as dephlogisticated air was confused, because, under the theory of phlogiston, oxygen would be the gaseous venue by which phlogiston can be released: a phlogisticating air, not dephlogisticated air.
Priestley’s discovery was made by heating mercury oxide by sunlight, and having mice, and later himself, breath the results (vaporous mercury and oxygen). His 1776 published description of isolated dephlogisticated air did not identify oxygen per se.
Antoine Lavoisier determined air as a mixture of gases, primarily nitrogen and oxygen. Lavoisier demonstrated oxygen as the agent of rusting metals and explained oxygen’s role in plant and animal respiration.
Lavoisier’s explanation of combustion disproved the phlogiston theory that Priestley held dear. Priestley never accepted Lavoisier’s outrageous speculations on oxygen, respiration, and rust. Instead, Priestley religiously defended phlogiston theory for the rest of his life.
Priestley, with a Presbyterian cast of mind, also maintained that humans have no free will, that instead conditions create dynamics with inevitable outcomes (predeterminism). According to Priestley, everything in Nature, including men’s minds, are subject to the law of causation, but because a benevolent God created all, perfection of man and the world would come in due time. For Priestley, evil arose only from an imperfect understanding of the world.
Swedish pharmaceutical chemist Carl Wilhelm Scheele was called “hard-luck Scheele” by American writer and biochemistry professor Isaac Asimov because Scheele made several chemical discoveries before others who are generally given the credit. Slow to publish, Scheele was scooped on oxygen, hydrogen, chlorine, barium, tungsten, and molybdenum. Scheele isolated oxygen about 2 years before Priestly, but Priestly published first.
Like Priestley, Scheele, who called oxygen “fire air,” described the gas using phlogistical terms, as he considered his discovery as confirming, not overturning, phlogiston theory.
Though Scheele did not comprehend the import of his oxygen discovery, others did. Scheele wrote Lavoisier about his findings. Lavoisier grasped the significance. Ironically, Scheele’s report was pivotal in invalidating the long-held theory of phlogiston.
“One giant leap for mankind.” ~ American astronaut Neal Armstrong while standing on the Moon in 1969.
The Apollo space program was NASA’s last step to having an American walk on the moon before the 1960s expired; a national goal set in 1961 by President Kennedy at the height of the Cold War, as a rallying distraction to the diplomatic follies of his and the Soviet Union’s administrations.
The distraction proved expensive. $24 billion was spent getting there; at the time, by far the largest commitment of resources ever made by any modern nation in peacetime. At its apex, the Apollo program employed 400,000 people, and required support from over 20,000 corporations and universities.
NASA management was always concerned about cutting payload. To this end, only pure oxygen was circulated in a spacecraft. Nitrogen, 78% of ordinary air, was considered deadweight.
NASA partly acknowledged the risk of using pure O2 in a 1966 technical report: “in pure oxygen [flames] will burn faster and hotter without the dilution of atmospheric nitrogen to absorb some of the heat or otherwise interfere.”
As soon as O2 absorbs heat, the molecule atomizes. Each atom raises hell by stealing electrons from nearby atoms; sizzling larceny that hots up any fire.
Worse, it takes but the slightest stimulation to spark an O2 orgy. Some NASA engineers fretted that static electricity from the Velcro on the astronauts’ suits might cause spontaneous combustion.
Yet NASA’s 1966 report concluded: “inert gas has been considered as a means of suppressing flammability… Inert additives are not only unnecessary but also increasingly complicated.”
In space, where there is no atmospheric pressure, just enough gas to breath is all that is needed. But on the ground, owing to atmospheric pressure, technicians had to pump the simulators with prodigious quantities of pure oxygen to keep the simulator walls from crumpling.
Mere months after the O2 report, during a capsule training simulation, an unexplained spark ignited a fire that cremated the 3 astronauts inside within seconds; whereupon NASA management decided that inert gases had something going for them after all.
Oxygen has been the volatile actor in life’s evolution. Before the origin of life, in the Hadean eon, atmospheric oxygen was negligible. Similarly, oxygen also exists in space, scattered about in minuscule quantities.
Anaerobic organisms arose that produced oxygen as exhaust, gaining usable carbon in a CO2 conversion catalyzed by sunshine: photosynthesis. Such phototropic life appeared within a billion years of Earth’s formation. Planetary oxidation had begun by 3.4 bya.
But not without counterforce. Methanogens thrived in the nickel-rich seas billions of years ago, belching methane into the air. The methane reacted with atmospheric oxygen, creating carbon dioxide and water.
Eventually, the methane party wound down. Oceanic nickel levels began to drop 2.7 bya. Nickel levels halved by 2.5 bya. The heyday of the methanogens had passed.
The mantle of the early Earth was so hot that dynamic tectonic plate flow with subduction did not begin until 3.0–2.7 bya. Oxidized material from Earth’s surface began being recycled into the mantle. Pressure at depth resulted in oxygen release into the atmosphere.
Relatively rapid oxidation, albeit fractional, resulted in Earth’s first extinction event. Much microbial life had evolved to survive in an oxygen-poor environment. The ascent of O2 spurred adaptation to greater oxygen tolerance. Over 2.5 bya, cyanobacteria evolved, and begat the slow oxygenation of the planet, by inhaling carbon dioxide and exhaling oxygen.
It took 500 million years for the atmosphere to begin oxidizing. The oxygen produced by photosynthesis readily reacts with ferrous iron and other elements to form precipitates, such as insoluble ferric oxide (e.g., rust).
1,500 different minerals were found on Earth prior to life arising, generated by dynamic mantle and crust processes during the first 2 billion years. Oxidation of Earth created 2,500 new minerals, many of those being oxidized and weathered products of predecessor minerals.
As the atmospheric oxygen level rose, every mineral that could be oxidized was. Once the weathering of iron-rich rocks abated, photosynthetic cyanobacteria belched so much oxygen so quickly as to overwhelm the planet’s ability to soak it all in.
Free oxygen poured into the air and oceans, radically altering geochemical dynamics, as well as the biochemical evolution of life. Rising atmospheric oxygen facilitated the evolution of eukaryotes: a significant step from purely prokaryotic life.
Life added to Earth’s mineral stock. Over 4,400 different mineral species have been cataloged. 400 have been added since eukaryotes arose.
Aerobic respiration may have presaged the oxygen surplus from photosynthesizers. 2.9 bya, the most ancient aerobic process produced pyridoxal (C8H9NO3), the active form of vitamin B6, and an oxygen-based enzyme, manganese catalase. The enzyme detoxifies hydrogen peroxide by breaking it down into oxygen and water. These early aerobic organisms may have got the oxygen needed for pyridoxal production by busting up hydrogen peroxide (H2O2), which might have come from glacial ice being bombarded by ultraviolet radiation, which generates generous amounts of H2O2.
UV levels were very high at the time, as the atmosphere had yet to form its later ultraviolet shield. The UV shield that eventually formed was an ozone layer in the upper atmosphere, a byproduct of atmospheric oxygen proliferation.
Despite unfiltered radiation from the Sun, early life prodigiously evolved. A few factors were in its favor.
1st, the Sun burned less brightly. 2nd, early life was in the oceans, which provided some protection from UV. 3rd, primordial bacteria developed protective mechanisms to limit DNA damage from UV radiation.
Without atmospheric O2, there is no O3 (ozone). Though the ozone layer accounts for only 0.00001% of the volume of atmospheric gas, its accumulation in the upper stratosphere, and its ability to absorb 99% of incoming ultraviolet rays, provides a blanket of protection for life on the surface.
Atmospheric oxygen stops water loss from the planet. Hydrogen released from water bumps into oxygen in the air before it wafts into space and is recaptured in rain droplets.
What started as a primitive organism waste product became organic fuel. The energy that can be derived from fermentation, or the early-evolved methane-sulfate reactions, are puny compared to the potency of aerobic respiration.
Nothing else could have powered multicellular life. All plants and animals depend upon oxygen for at least part of their life cycle. Only the early risers, microscopic life, managed to eke out an existence before the oxygen bloom.
The proof is prehistoric. 300 mya, atmospheric oxygen levels were 66% higher than today. The higher oxygen levels greatly affected the species that adapted to intake oxygen as an energy source. Paleozoic amoebas were 100 to 1,000 times larger than they are now.
During the atmospheric oxygen bloom, insects supersized. Dragonflies had wingspans of 70 centimeters. There were millipedes over a meter long.
Oxygen transport in vertebrates is through the bloodstream. But insects move air through their bodies by trachea: an internal network of channels. So, for insects, a higher oxygen level readily supports rapid evolution of a larger body.
While life adjusted to an oxygenated atmosphere, its toxicity at the cellular level remains a challenge. Both plant and animal tissues create anoxic conditions to generate stem cells.
“Oxygen is a diffusible signal involved in the control of stem cell activity.” ~ Italian botanist Francesco Licausi et al
“To a large extent, the structure, behavior and unique qualities of each living being are a consequence of the proteins they contain.” ~ American molecular biologists Kathleen Park Talaro & Barry Chess
A protein is a linear polymer comprising amino acids; an intricate, living macromolecule, typically folded into a globe or fiber as a product of energetic economy. Hyperactive proteins are called enzymes.
“The assembly of amino acids into a protein is done one at a time in linear and specific order.” ~ American biologist Peter Ward & American geobiologist Joe Kirschvink
Organisms are largely built of proteins. The concept of the set of proteins in a cell comprises a cell’s proteome.
Proteins are the workhorses of biochemistry, the chief actors for cells: for structure (e.g. membranes) and internal organization, cell identification, transportation (of molecules within, and in and out of cells), communication (signaling, both within and between cells), enzymatic action, gene regulation, defense, and cell movement.
Proteins are commonly capable of a variety of functions. But they often need to be focused to a specific task, else they might create cellular havoc, possibly leading to disease.
This is where cofactors come in. A cofactor is a chemical compound that binds to a protein as a requisite for tailored activity. Enzymes are typically activated by cofactors, which essentially act as helper molecules.
Proteins are characterized by 4 levels of structure: chemical (amino acid composition) (aka primary); fold pattern (aka secondary); 3d shape (aka tertiary); and assembly (aka quaternary): an assemblage of several protein molecules, such as with polypeptide chains. These characteristics are genetically defined. Each is significant in protein functioning.
“Proteins are machines that have structures and motions.” ~ American molecular biologist Peter Wright
Proteins remember their situation, and so are prepared for what needs to be done next. For example, by binding to a certain ion, a protein activates. In this way, the state of a protein at any instant embodies a memory of its past.
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“Even tiny changes in structure can greatly affect the properties of a compound.” ~ German chemist Jochen Küpper
Most molecules occur in several shapes, each of which may behave differently from another. Spatial arrangement can also affect reaction rate.
Such diversity of functionality by topology is at the heart of protein functioning. These multifaceted macromolecules evolved to perform different complex functions. One aspect of this comes in the way that proteins fold and unfold to activate and perform tasks.
“Cells achieve their steady-state size by adding a constant volume between birth and division, regardless of their size at birth.” ~ Korean Canadian molecular biologist Suckjoon Jun et al
To realize their biological potential cells must attain a specific size. They do so through certain proteins responsible for growth which are also involved in cell division. The proteins know exactly how large a cell should be by understanding the genetic blueprints. Some eukaryotic cells edit the instructions that these proteins get to alter desired cell size.
“It’s a very robust mechanism.” ~ Suckjoon Jun
“Over time, Nature improved protein folding so that more complex structures were able to develop.” ~ German molecular biomechanist Frauke Gräter
Proteins are produced in a linear sequence. The last step in protein production is an elaborate folding to a shape that is energetically in repose.
Chaperonin are proteins that provide scaffolding for the first folding of new proteins. In a high-speed origami, a protein assumes its final resting structure in a series of rapid incremental steps. This process takes only a few seconds, thanks to active guidance by chaperonin.
A protein partly unfolds to work. How a protein folds/unfolds, and at what speed, has much to do with its performance. The intricate 3d conformation of a protein is essential to proper functioning.
A major force behind protein folding and polypeptide interaction is the avoidance of water by hydrophobic amino acids. Internal friction, which reflects the energy landscape of the protein, plays an important role in the dynamics of folding, and the ability of a protein to function properly. The folding process of a protein is regulated by a formidable network of other proteins, and other, smaller molecules.
“There is a sea of small molecules – ligands – with which proteins live in a cell. These ligands, which are very small molecules only about 100 daltons in size, are critical in determining the behavior of folding macromolecules on the order of 100 kilodaltons in size, that is 1,000 times larger. It’s like the mouse telling the elephant what to do.” ~ American biochemists Lila Gierasch & Scott Garman
Many proteins have varying degrees of folding. 30% of human proteins have unfolded portions.
Whether a portion of a protein is folded or not, or even appears disordered, is scant indication of what it may do. A protein operates orchestrally; all portions have some part to play. The odd bits may be the piece that give a protein its versatility.
Many diseases result from misfolded proteins (prions), notably neurodegenerative diseases associated with aging, such as Parkinson’s, Huntington’s, and Alzheimer’s. Prions refold from a harmless form into one that is malicious and contagious, initiating a chain reaction that creates a prion aggregation.
RfaH is a protein which activates genes that allow E. coli bacteria to launch a successful attack on a host, inciting disease. It does so by folding itself into different shapes, and by doing so is able to perform vastly different tasks. Specifically, RfaH can fiddle with both genetic transcription and translation – coupling the two together – by smartly changing its shape.
Dihydrofolate reductase (DHFR) is an enzyme common to E. coli bacteria and humans and everything in between. DHFR structure is almost identical throughout all life forms; conserved through evolution, as its function is critical for the synthesis of DNA. But how DHFR unfolds to expose amino acids differs between bacteria and primates. This distinction in atomic dynamics makes a significant difference.
DHFR motion evolved to fit the cellular environment. Human DHFR is so well-tuned for its own cells that it won’t work in bacteria: the product molecules in E. coli are packed too tight for human DHFR to function.
“Man’s health and well-being depends upon, among many things, the proper functioning of the myriad proteins that participate in the intricate synergisms of living systems.” ~ American biochemist Stanford Moore
The activities of proteins are regulated in a variety of ways, though the agent almost always involves a ribozyme, which is an RNA-based enzyme.
Proteins have an active site, where substrates bind and undergo a chemical reaction. The active site is not the only place which affects protein functionality.
Protein activity can be regulated by an effector molecule binding at a protein’s allosteric site: a site that is not the active site. Enzymes are typical effector molecules at allosteric sites, though an enzyme too is subject to allosteric regulation, as enzymes are themselves proteins.
Allostery refers to protein regulation at an allosteric site. Allosteric regulation is the action of effectors at the allosteric site, either enhancing a protein’s activity (allosteric activator), or tamping it down (allosteric inhibitor). Allostery is especially important in cellular communication, where getting the right message across is critical.
Because eukaryotic cells are highly compartmentalized, localized control of allosteric regulation is widespread in eukaryotes. Enzymatic activity is itself controlled in a process called localization.
Enzymes are continually active. They evolved for flexibility. An enzyme’s effect can be controlled by small binding sequences that localize its activity. Enzymes can be very selective about the 3d configuration of the molecules they interact with.
Localization by proximity control is readily engineered because binding sequences to effect such control are typically simple. Because the biomechanics are so easy, much genetic regulation is allosteric.
Generalists & Specialists
Functionally, proteins are either specialists or generalists. Evolution begat specialization. The earliest cells were chock full of jack-of-all-trades proteins.
Generalists dabble in various chemical reactions. They carry out tasks which are less crucial, such as making vitamins (B12), which cells need in limited quantities.
Specialist enzymes stick to specific substrates. Their devotion means that they are more trusted to essential functions, such as translating genetic instructions into new proteins, and the constant chore of producing cellular energy.
A specialist protein senses when a mammal’s lungs are full of air, and so plays a critical role in regulating breathing. The same protein is instrumental in the sense of touch and proprioception.
The importance of generalists is not to be understated, particularly in their having a broader vista than their dedicated brethren. In eliminating toxic substances, a cell is better off with enzymes that can recognize and deal with more than one kind of danger.
“Due to specific demands, sensory systems evolve independently in different species.” ~ American neurobiologist Marco Gallio et al
An animal’s tolerances and perils vary greatly, depending upon a creature’s adaptation to its environment. Some cells can withstand intense heat, or, conversely, cold, which would damage the average cell; likewise for mechanical stresses and chemical agents.
All animals feel pain, and they do so via the same protein, termed TRPA. TRPA acts as a stress detector (nociceptor). Nociception is detection of stimuli – whether chemical, mechanical, or thermal – which are hazardous to a cell. Via TRPA activity, nociception initiates pain to warn an animal.
“TRPA is remarkably conserved across animal evolution.” ~ Marco Gallio
While TRPA is essentially the same for all animals, its detection acumen varies vastly, depending upon what defines danger for the host in which it works. While the consummate generalist in sensing danger and sounding an alarm, TRPA is specifically adapted to know exactly what tolerances a cell has. How such exquisite expertise can be attained or held by a protein is mysterious.
Even proteins that share the same topology may perform a different range of tasks. The orientation of a protein’s amino acids affects electron transfer efficiency.
Channeling energy differentially alters the branching to neighboring amino acids. A simple electron channel change can cascade into a considerable functional difference.
A protein is a complex polymer built up from chains of polypeptide molecules. A polypeptide is formed from numerous peptides.
A peptide is a short polymer of amino acid monomers, linked by peptide bonds. The shortest peptide is 2 amino acids joined by a single peptide bond: a dipeptide.
An amino acid is a molecule comprising: a carboxylic acid group, an amine group, and a side chain specific to the amino acid. Amino acids have as key elements: carbon, hydrogen, oxygen, and nitrogen.
An alpha-amino acid has both the amine and the carboxylic acid groups attached to the first (alpha-) carbon atom. The 22 proteinogenic (protein-building) amino acids are all α-amino acids.
The side chain of an amino acid is a hydrocarbon. A hydrocarbon is a molecule comprising only carbon and hydrogen. The side chain alkyl group is historically designated by R. Methane (CH4) is the simplest alkane.
An amino acid side chain may have a string of carbon atoms. The number of carbon atoms defines the size of an alkane.
A carboxylic acid is a polar molecule: –CO2H, connected to a hydrocarbon. A carboxylic acid completes itself with a side chain. The minus in –CO2H indicates a negative charge.
An amine is a derivative of ammonia (NH3). To be an amine, 1 or more of nitrogen’s links to hydrogen is replaced by a bond to an alkane (an amino acid side chain). The nitrogen atom has a lone pair. An amine group is thus: NH2R, NHR2, or even NR3.
Amino acids are simple and easily constructed. The building blocks of life are chemically quite accessible.
20 different amino acids form the basic building blocks of peptides. Some polypeptides have all 20 amino acids, some fewer.
There are 22 known amino acids, but only 20 are commonly used in extant life forms. Amino acid number 22, the most recently discovered, is very ancient; used by methane-producing archaea (methanogens) for energy conversion.
The human body can synthesize 12 of the 20 standard amino acids but cannot make the other 8. These 8 must be supplied by dietary intake, and so are termed essential amino acids.
There are also conditionally essential amino acids, which some humans can sufficiently synthesize but others cannot, and so must get adequate amounts from their diet. The difference between essential and non-essential amino acids is not fully understood, as some amino acids can be derived from others. For example, the various human amino acids containing sulfur can be converted internally, but not synthesized de novo in humans.
There are specific proteins for many of the hundreds of millions of species that have lived on Earth, and for different plant parts and body organs. The human body synthesizes ~100,000 different proteins.
The vast diversity of proteins comes primarily from variety in amino acid R groups, which yield markedly different chemical and functional properties. Proteins often have over 100 amino acid units.
There are more ways to arrange 20 units in a chain of 100 than there are atoms in the universe; whence the practically unlimited information storage capacity of proteins.