Water
“If there is magic on the planet, it is contained in water.” ~ American anthropologist Loren Eiseley
Hydrogen and oxygen are among the most abundant elements in the universe. Water exists in outer space. Much of the water in space is a byproduct of star formation.
Water formed early in the history of the universe; as quickly as supernovas spread oxygen into space, allowing it to combine with ubiquitous hydrogen gas.
In most planets, including Earth, vast reservoirs underground belie a more modest surface presence. The Moon appears bone dry but stores considerable H2O within.
Even Mercury, with daytime surface temperatures reaching 400 °C, hot enough to melt lead, has ice tucked within.
History
Philosophers from antiquity regarded water as a basic element, typifying all liquids. Scientists did not discard that view until the back half of the 18th century.
English chemist Henry Cavendish discovered hydrogen in 1766, calling it “factitious air.” Cavendish synthesized water in 1781 by detonating hydrogen, which combined with the oxygen in the air to form water as a byproduct.
Lavoisier termed both oxygen (1779) and hydrogen (1783) and determined water as a combination of oxygen and hydrogen, thus dispelling water as an element unto itself.
Properties
“Water shows exceptional properties, different from all other liquids.” ~ Austrian chemist Katrin Amann-Winkel et al
Pure water is an odorless, tasteless liquid, with a bluish tint observable only at a deep volume. Water is transparent in the visible spectrum, but strongly absorbs both infrared and ultraviolet light.
Water is so commonplace that it is regarded as a typical liquid. Water is anything but.
Water is one of lightest gases. Liquid water is much denser than expectable. Water’s solid phase – ice – is surprisingly featherweight in light of its liquid form.
Water can be extremely slippery and exceedingly sticky simultaneously. This slippery/sticky behavior is how the feel of water is sussed. There is no direct sensation of moisture.
Every body of water, from droplets on to more voluminous realms, is an intricate viscoelastic system, with an underlying extended structure governed by quantum forces; whence water’s enduring mystery.
“Although water is one of the simplest molecules, it is in reality a very complex liquid displaying more than 64 counterintuitive anomalies, most of which have not been adequately explained.” ~ Italian physicist Francesco Mallamace et al
Water & Life
“Life as we know it depends absolutely on the presence of liquid water, and the existence of the liquid state sets the boundary conditions for life itself.” ~ Andrew Clarke
The quirks of water are what makes life possible. The high cohesion of water molecules gives it high freezing and melting points, making its liquid phase predominate: the phase upon which life directly relies. Water’s thermodynamics – its large heat capacity and high thermal conductivity – vitally help organisms control body temperature.
The high latent heat of water defines an evaporation which resists dehydration while yielding considerable evaporative cooling. Water’s unique hydration properties critically affect important biological macromolecules – proteins and nucleic acids – that determine their 3d structures, and hence their biological functions.
Water being a small molecule, an excellent solvent, and having high relative permittivity, all contribute biological utility. The easy ionization and proton exchange of H2O contributes to a rich repertoire of biological interactions.
Water’s thermal connectivity and density characteristics means that all of a body of freshwater, not just its surface, must be close to 4 °C before freezing can start. The freezing of lakes, rivers, and seas is top-down. This insulates water from further freezing, permitting bottom dwellers to survive.
Water reflectivity aids rapid thawing. Water’s density-driven thermal convection affords seasonal mixing in deep temperate waters, carrying life-essential oxygen into the depths.
The compressibility of water lowers sea level ~40 meters, yielding 5% more land on Earth. Water’s high surface tension, and its expansion on freezing, engenders the erosion of rocks, providing soil upon which life on land may thrive.
The tremendous heat capacity of oceans allows them to act as global thermal reservoirs, such that sea temperatures vary only 1/3rd as much as land temperatures, and so moderate the world’s climate.
The quantum properties of water facilitate life in innumerable ways, from the molecular to the planetary scale.
Polarity
A water molecule is in sum electrically neutral, but the atomic arrangement of a water molecule is not linear, so there is a skew of electrical charges: a slight negative charge at the oxygen atom interface, while the hydrogen atoms are slightly positive. Hence, water is a polar molecule, with an electric dipole moment. This creates a weak and fleeting affinity for hydrogen bonding. That affinity has a cumulative effect, in making water molecules slightly cohesive. Water’s flowing effect owes to hydrogen bonding, which is somewhat like static cling.
By having one hydrogen atom in the drink and the other in the breeze, about a quarter of the water at the interface between a water body and atmosphere straddles the liquid-gas phases. This layer of molecular ambiguity at the interface is extremely thin – 0.3 nanometers – and has no effect on the behavior of water molecules below because of their submerged location.
At extremely low temperatures, a water molecule has quantum tunneling ability forbidden in the classical world. The oxygen and hydrogen atoms are delocalized: present in all 6 symmetrically equivalent positions simultaneously. This only occurs at the quantum level, with no parallel in everyday experience.
“The average kinetic energy of water protons at almost absolute zero temperature is about 30% less than it is in bulk liquid or solid water. This is in complete disagreement with accepted models based on the energies of its vibrational modes.” ~ Russian nuclear physicist Alexander Kolesnikov
Bonding
“What makes water molecules unique is that they can form hydrogen bridge bonds.” ~ English physicist Gareth Parkinson
The molecular bonding of water has a strangely strong looseness: constantly rearranging restless hydrogen into irregular molecular arrangements owing to the polar alignment of H2O. Water molecules have the unique ability to form short-lived connections that create fleeting networks. These transitory networks are crucial to water’s peculiarity in many ways.
The oxygen in H2O seeks connections with the hydrogen in neighboring water molecules, consistently maintaining tetrahedral coordination between close neighbors. At any given time, an oxygen atom in water is bonded to roughly 3.6 hydrogen atoms, with ever-changing interaction patterns on a picosecond time scale. The hydrogen bonds that link water molecules together break and form several thousands of billions of times per second.
Yet, measured at the level of attoseconds (10–18 seconds), water is not H2O. It is instead H1.5O. This was demonstrated in a 1995 British neutron-scattering experiment that found 25% fewer protons in water than expected.
The protons in hydrogen were sometimes not detectable by neutron probes. The reason is quantum entanglement. This phenomenon also occurs in other hydrogen compounds.
Water’s bonding network is extensive. A single ion can tweak the bonds of several million water molecules over a distance exceeding 20 nanometers, causing the liquid to stiffen. This networked entanglement is central to many of water’s unique properties, beginning with surface tension.
Surface Tension
“You may have inner tranquility, but you can’t escape surface tension.” ~ American evolutionary biologist Louise Roth
Water’s high surface tension owes generally to H2O polarity, which creates molecular bonding affinity, but specifically results from hydrogen bonding at the water-air interface, where adsorption of hydroxide ions occurs. Rounded raindrops reflect it. Ice is an unusually strong molecular solid because of it. Earth’s rapid water cycle is abetted by surface tension.
Nonpolar volatile molecules like carbon dioxide (CO2) and methane (CH4) can’t form droplets. Instead, they float airborne as a fine mist, unable to rain.
Surface tension aids capillary action: the ability to flow against gravity when confined, such as in a thin tube. Owing to hydrogen bonding networks that form when water is squeezed, water molecules move faster as their density increases.
Vascular plants (tracheophytes) rely upon water’s facile capillary action for liquid transport between roots, wood, stems, and leaves. Without capillary action, tracheophytes could not have evolved.
For thousands of years it was known that oil may float on water, but that the converse was not true. But it is possible for water to float on oil, depending upon the size of the water droplet and the surface tension of the oil. Mineral oils have a low surface tension, and so will not support water; but commercial vegetable oil will.
Solvency
Water is an ionizing agent. Because of its polarity, the positive and negative poles exert forces that can pull apart other molecules. As such, water is a solvent.
Rocks are slightly soluble. Hence, water erodes solid landscapes, fashioning valleys and canyons, ferrying inorganic nutrients in its flow.
Sugar dissolves in water because sugar’s slight polarity is matched to covalent hydrogen bonds. This polarity prevents sugar molecules from regaining integrity. Because water is a polar molecule, other polar molecules are generally water-soluble.
Salt (sodium chloride (NaCl)) crystals enveloped by water dissolve because its 2 elements succumb to water’s polarity: positive sodium ionization and negative chlorine ionization. This is seduction by electrical deception. It takes several water molecules to simulate the ionic appeal that divorces sodium from chloride in salt’s natural crystal ionized form.
Insoluble ions are not so easily seduced: their ionic bonds remain faithful, unpersuaded by H2O’s partial charges. By such chemical stoicism, insoluble substances resist water’s siren song of solvency.
Salt and sugar are the leading actor and actress in life’s metabolism theater. A lot of organic molecules besides salt and sugar are water-soluble. Others, such as oils, are not.
The difference in organic solubility comes from attachments. If atoms other than carbon (C) and hydrogen (H) are present, a molecule has partial charges that can be emulated by water.
Hydrocarbon oil, naturally occurring as petroleum (crude oil), is steadily insoluble. But the sugar glucose (C6H12O6), thanks to ionic oxygen in the mix, readily succumbs to dissolution.
Positively charged ions (cations) are loath to pair up, strongly preferring anions for their partner. But when liquored up by water, cations do couple.
Forms
Water’s seeming simplicity is deceptive. The pairing of an oxygen atom with 2 of hydrogen has variants which abstrusely affect water’s behaviors. Further, the thermodynamics of water are entangled with the molecular constructions which this sublime fluid takes. At both the atomic and molecular levels, water is a system like no other.
Atomic Forms
Water has different forms, both as isomers and isotopes.
Isomers
Water has 2 distinct nuclear-spin isomers of its hydrogen atoms: para-water and ortho-water. The isomers are stable, though interconversion does mysteriously occur sometimes. Ambient-temperature water is a mixture of these 2 isomers.
Whereas the nuclear spin of the hydrogen atoms in para-water sum to 0 via asymmetric wavefunctions, ortho-water nuclear spins symmetrically sum to 1. The atomic twists and turns of para-water let it react 23% faster than the stolider ortho-water. Para-water is also able to more strongly attract its reaction partner than ortho-water, which leads to enhanced chemical reactivity.
Isotopes
Water has several isotopes, especially hydrogen, though oxygen too has isotopes. Known oxygen isotopes range in mass number from 12 to 28. Oxygen has 3 stable isotopes: 16O, 17O, and 18O.
16O is the most abundant: 99.762%. 16O is the principal form by stellar evolution, as 16O can be made by stars that were initially fueled entirely by hydrogen.
17O and 18O are produced later in the stellar life cycle. 17O comes by burning hydrogen into helium during the CNO cycle.
18O is common in the helium-rich zones of stars. 18O is typically produced when 14N (nitrogen), abundantly derived from CNO combustion, captures a 4He (helium) nucleus.
Hydrogen has 2 stable isotopes: protium and deuterium. Protium has a nucleus of a single proton. Deuterium (aka heavy hydrogen) has a nucleus comprising a proton and a neutron. Both protium and deuterium sport a single electron.
Water is typically a mixture of protium and deuterium, naturally varying by source. More than 99.98+% of the hydrogen in ocean water is protium.
Recently evaporated ocean water, including rainwater, river water, and snow, tends to have the lighter isotopes of hydrogen and oxygen. Such waters evaporate faster than heavier varieties.
Heavy water, properly termed deuterium oxide (D2O), has a richer deuterium content. 0.0156% of the hydrogen in ocean water is deuterium.
Rats avoid heavy water by its smell. Humans are generally unaware of the difference (from ordinary water), other than heavy water sometimes has a sweet flavor, or causes a burning sensation.
Conversely, light water is deuterium depleted. Light water has been found beneficial for mammals with cancer.
There are differences in the bonds between light and heavy water, at the atomic and molecular levels.
The distance between deuterium and oxygen nuclei in D2O is 3% shorter than the distance between hydrogen and oxygen in an H2O molecule. Conversely, the hydrogen bonds between one molecule of water and another are 4% longer in heavy water than light water.
Despite their differences, light and heavy water behave much alike. Their melting temperatures differ by less than 4 °C, and they boil at even closer temperatures. This owes to competing nuclear quantum effects (NQEs) being offset along different molecular axes.
One NQE is associated with motion perpendicular to the plane of a water molecule. Another NQE relates to motion parallel to the hydrogen bond.
The difference in nuclear quantum effects between H2 and D2 is that one NQE increases while the other lessens. Hence, a small quantum net effect between the different hydrogen forms, and there is similarity in phase transitions.
Another oddity of water is its lack of chemical purity. Water is never just H2O. A small fraction of any body of H2O splits into positively charged hydrons: hydrogen ions (H+), which are protons without bonded electrons, and negatively charged hydroxyl groups (OH–).
Hydrons latch onto water molecules, forming hydronium ions (H3O+). So-called “pure water” at room temperature balances equal numbers of positive hydronium ions and negative hydroxyl groups, creating a neutral solution (pH = 7).
Molecular Forms
“Water has the remarkable property of being more compressible in winter than in summer.” ~ English chemist John Canton in 1764
Water molecules take 2 structural forms, based upon the instant inclinations of the hydrogen bonds. One structure is tetrahedral, the other disordered, in a myriad of ways.
“Despite its simplicity, water tends to form tetrahedral order locally by directional hydrogen bonding. This structuring is known to be responsible for a vast array of unusual properties.” ~ Japanese physicist Hajime Tanaka et al
Water is densest at 4 °C, when its tetrahedral ordering is at its apex. Heating reduces the number of tetrahedral structures, resulting in more disorder.
In water cooler than 4 °C, disordered areas are more densely packed than the tetrahedrons. Agitated by heat, warming above 4 °C spreads out molecules in disordered regions, making the water less dense.
In 1762, Scottish chemist Joseph Black discovered latent heat: that materials may absorb heat without changing temperature. Black also noted that different substances have dissimilar heat capacities.
Directly related to latent heat, heat capacity is the ratio of heat absorbed by a substance to its temperature change. In other words, heat capacity is the measure of how much the temperature of a certain substance rises in response to heating: a ratio of heat energy/temperature. The term specific heat capacity is a measure of heat capacity per unit mass of a material.
Heat capacity is an extensive property: proportional to the size of the system. In the instance of H2O, heat capacity proportionally varies by how big the body of water is.
Water has an exceptionally high specific heat capacity: it takes a lot of energy to heat water. Water is laden with latent heat. This is because much of the extra heat applied to H2O converts molecular regions from tetrahedral form to disordered structures, rather than increasing the kinetic energy of the water molecules, which would raise the temperature.
Water’s specific heat capacity is at a minimum at 35 °C, increasing as its temperature falls or rises. The heat capacity of most other liquids rises continuously with temperature.
Between 0–35 °C, increasing water temperature steadily removes regions of ordered, tetrahedral structure, reducing the ability to absorb heat. Above 35 °C, so few of the tetrahedral regions are left that water behaves like a regular liquid.
The strong attraction between water molecules keeps them more tightly packed than most other liquids. As such, water is especially difficult to compress. Compressibility is at its lowest when the higher-density disordered structures dominate.
With most liquids, compressibility rises continuously with temperature. Water is anomalous. As water hots up, the dense, disordered regions become more prevalent; areas which are more difficult to compress. Still, heat forces molecules within disordered regions further apart, making them more compressible. Disordered region expansion takes precedence over structural shifts beyond 46 °C.
The speed of sound in water increases up to 74 °C, after which it starts to fall again. This owes to the interplay between water’s density and compressibility profiles, which stem from the changing balance between the 2 structural forms.
Exceptionally, water molecules diffuse more easily at higher pressures. That is because high pressure converts more molecules to a compressed disordered structure, making them more mobile. Water becomes less viscous at higher pressure.
Increasing pressure increases the amount by which water expands when heated. Rising temperature expands disordered regions more rapidly than ordered, tetrahedral ones, and high pressure forces fluctuations in disordered regions.
Properties such as viscosity, melting point, and boiling point are significantly different in heavy water (D2O) from H2O. The heavier isotopes change the quantum mechanical properties of water molecules, altering the balancing act between tetrahedral and disordered regions.
Phases
States of matter are also called phases. Whereas state stresses stasis, phase suggests fluidity via energy level.
A molecular system may be arranged in a certain variety of configurations depending upon total energy in the system. Constituent molecules have the liberty of greater possible configurations at higher temperatures/energies and can freely move about. This is the gaseous phase.
At lower temperatures, configurations are more constrained. At this more ordered phase, the material system is liquid.
If the temperature/energy drops to a certain threshold, the molecules are confined to a specific configuration, forming a solid.
This picture is common for relatively simple molecules, which have 3 distinct states: gaseous, liquid, and solid. such CO2 is exemplary. For more complex or slippery molecules, including H2O, there are greater possibilities for combinations, giving rise to more phases.
Liquid crystals are a beautiful illustration of this: complex organic molecules which can liquidly flow within a solid-like crystalline structure.
Water is one of the few chemical compounds that exists alternately as a gas, liquid, or solid within the temperature range of ordinary life, with an intricate quantum dance in transitioning from one state to another. Whereas water exhibits its phase range within 100 ºC, most other compounds have a much broader range for their temperature transitions.
According to its structural properties, water should be a gas at the temperature and pressure where water is liquid. Consequently, liquid water behaves much differently than other liquids.
Water has more than the 3 well-known phases (gas, liquid, solid), and may selectively exist in more than 1 phase simultaneously.
Liquid water uniquely exhibits 2 distinct states, with different rates of thermal expansion, surface tension, and refractive index (a measure of how light travels through it), as well as other characteristics. The thermal crossover for these states occurs at ~50 ±10 °C.
“Water at room temperature can’t decide in which of the two forms it should be, high or low density, which results in local fluctuations between the two. Water is not a complicated liquid, but two simple liquids with a complicated relationship.” ~ Swedish physiochemist Lars Pettersson
Water boils at 100 °C. At a critical temperature and pressure (374 °C and 22.064 MPa), found naturally only in the hottest regions of hydrothermal vents, water becomes supercritical: the gas and liquid phases merge into a homogeneous fluid phase, with properties from both phases.
Water’s thermal conductivity is not constant. Water cooled to 225 K starts to more efficiently conduct heat. Below this temperature, liquid water undergoes sharp but continuous structural changes, becoming highly ordered: very much like ice.
Water cooled to less than –135 °C turns syrupy but clear: glassy water. This phase is well past freezing.
Freezing
“The fact that water has previously been warmed contributes to its freezing quickly.” ~ Aristotle
Water nominally freezes at 0 °C, but, like most other liquids, may exist as a supercooled liquid: staying fluid below its freezing point. Liquid water at –40 °C has been found in clouds. Water may remain a liquid to –48 °C. Supercooled water freezes if disturbed.
The freezing of water is controlled not only by its temperature, but also by its size. The nucleation of ice in small droplets is strongly size-dependent.
Freezing is a process of water molecules forming a bonded network. In transition to freezing, water physically changes into tetrahedrons, with each water molecule loosely bonded to 4 others.
Freezing water droplets garb themselves with pointy tips, sporting fractal shapes as they crystallize. Depending upon temperature and pressure, ice may take various crystalline forms while water molecules cling to each other via hydrogen bonds.
Aristotle, and others after him, noted that hot water freezes faster than cold water. Yet namesake credit is given for its 1963 rediscovery by Tanzanian scientist Erasto Mpemba, who first noticed the Mpemba effect when he was 13, while making ice cream in haste: mixing boiling milk with sugar, rather than waiting for the milk to cool first, as instructed.
Mpemba asked his physics teacher for an explanation. The teacher told Mpemba that he must have been confused, as what he supposedly saw was impossible. Mpemba persisted in his inquiry.
Though the mechanics of the Mpemba effect remain something of a mystery, the likely explanation lies in the strength and arrangement of hydrogen bonds, which are affected by temperature, and partly depend upon nearby water molecules.
Ice
There are 17 known forms of ice. Many of them form under extreme pressure, such as in the interiors of frozen planets; 2 take on a quartz-like crystalline structure.
Whereas most solids sink into their liquid forms because the solids are denser than the liquids, ice floats on top of water. Ice, like water, has unusual structural properties.
Water attains its maximum density at 4 °C. Liquids typically condense as they cool but water expands upon freezing.
Whereas molecules of water vapor behave rather independently, ice forms a strong structural lattice. Liquid water molecules have weaker interactions with each other than they do in ice. Water also exists in a liquid-crystal state when adjacent to hydrophilic surfaces, which tend to interact with water or be dissolved by it.
Ice possesses quantum properties opposite those of other crystals. Typically, crystalline molecules shrink as they cool, but the shrinking stops before reaching absolute zero (0 K), owing to zero-point energy.
In theoretical physics, zero-point energy is the lowest possible energy that a particulate matter may have: the energy of its ground state. Zero-point energy is inherently tied to the uncertainty principle.
Typically, less massive atoms have more zero-point energy. Lighter nuclei need more room to move than heavier elements, which translates into larger crystals for lighter elements as they cool.
At higher temperatures, this quantum effect becomes less pronounced. Hence volume differences decrease as temperature rises.
The opposite occurs with ice. D2O (heavy water) occupies more volume than H2O ice. This volume difference increases as temperature rises.
“In order to access and measure quantum mechanical effects in matter, we usually need to go to very low temperatures, but in water ice some zero-point effects actually become more relevant as the temperature increases.” ~ physicist Marivi Fernández-Serra
The structural integrity of ice is illustrated in its melting. Ice melt is not continuous. Instead, ice discontinuously liquefies layer by layer.
Ice’s structure varies at different scales. Macroscopically, ice crystals form a hexagonal lattice, which is why all snowflakes are 6-sided. In contrast, in crystallites of up to 100,000 molecules, ice crystals form from water in layers of hexagonal and cubic arrays.
Ice formation in clouds – nucleation – is a critical aspect of precipitation development. The semi-disordered cubic-hex stacks of ice that form from cloud droplets facilitate faster development of rain than if ice nucleation occurred solely like snowflakes (hexagonally). Earth would be much different if ice formation physics was otherwise, as the water cycle would be hindered.
Evaporation
“Evaporation turns out to be a process driven by very small temperature differences. Often, only 10/10,000th parts of Kelvin are enough to make it happen.” ~ Polish physicist Daniel Jakubczyk
Evaporation is critical to life, and fundamental to the world’s climate by way of the water cycle.
For evaporating droplets, the temperature before evaporation starts and after it is accomplished is the same. Tiny temperature fluctuations between these moments tells evaporation’s tale.
The heat flux between a droplet and its surroundings plays a key role in evaporation. If the thermal flux between a hot pan and water drops were efficient, the water would instantly evaporate in boiling off. Instead, this flux is hindered, as droplets slide on an insulating layer of water vapor. This thermal layer is thick enough to restrain the heat flux. The thickness of the thermal layer depends upon surrounding conditions, not droplet size.
Organic Role
Water acts as a catalyst in many biochemical reactions by stabilizing charge states on a surface. Via condensation, large organic molecules, which have a skeleton of carbon and hydrogen, are built up by forming covalent bonds and eliminating water. These organic molecules can be broken down by adding water, as covalent bonds are split by hydrolysis. In the hydrolysis ionization reaction, both the organic molecule and the water molecule are split.
Water’s stickiness is instrumental in engendering life, as water’s properties assist both the formation and concentration of organic molecules (anabolism), and their transition to simpler forms (catabolism). Thus, water abets life at the molecular level by facilitating metabolism.
Cells are 60% water and 40% macromolecules: proteins, carbohydrates, and other components. Water confined within a cell slows by a factor of 10 compared to pure water alone. This impacts the intricate cellular machinery that is self-organized and precisely choreographed to transform energy, manufacture molecular products, and be responsive to its environment.
“While it is tempting to focus on the macromolecules, it is becoming increasingly clear that the water is not simply an innocent bystander, but instead plays an active and sometimes decisive role in mediating the processes carried out by the cell’s machinery.” ~ Canadian chemist Kevin Kubarych
Water has a high heat capacity, so it buffers aqueous systems from dramatic temperature changes. As water evaporates, it cools the liquid that remains. Frozen water forms an ice lattice, taking up more space. Hence, ice floats, providing a blanket of insulation.
At the planetary scale, water offers some stability by mediating temperature extremes; another benefit to life.
Fear of Water
“The hydrophobic interaction is one of the most important driving forces in Nature and is key to processes such as protein folding and the self-assembly of lipid membranes.” ~ Dutch chemical physicist Huib Bakker
Water-repellent molecules are termed hydrophobic; literally: “fear of water.” This molecular fear puts the surrounding water on guard.
Hydrophobic groups in molecules enhance the ordering of the surrounding hydrogen-bond network of water. These fearful groups form ideal templates around which a water network can fold, leading to greater local tetrahedral order, like water on its way to freezing.
A hydration shell forms around a hydrophobic group. Water network folding creates ridges 0.3 nm high, which is the intermolecular distance of water molecules. Ridge angles are 104.5°, corresponding to the bond angle of a water molecule. This is typical of hydrogen-bond networks in water.
Temperature matters. The hydration shell surrounding a hydrophobic group melts as water warms. The hydrogen-bond network gradually weakens. Disorder gains the upper hand as water hots up.
Size matters. Alcohols with hydrophobic chains longer than 1 nm are contrary in their effect on hydrogen bonding from hydrophobic groups half that size.
While small hydrophobic structures enhance order, large ones cause surrounding water molecules to have more broken hydrogen bonds, as the curvature of the large hydrophobic surface does not fit the 3d spatial arrangement of water hydrogen bonds. At temperatures above 80 °C, water molecules surrounding such plump hydrophobic groups are less ordered than their nearby brethren in bulk liquid.
Water Memory
“Homeopathy is wholly capable of satisfying the therapeutic demands of this age better than any other system or school of medicine.” ~ American physician Charles Menninger
In 1796, German physician Samuel Hahnemann claimed a medicinal treatment based on a “law of similars,” where “like cures like.” This notion gave rise to homeopathy: medicines made from substances causing similar symptoms in healthy people, prepared by serial dilutions until little or none of the supposed active ingredient remains.
In 1988, French immunologist Jacques Benveniste published an article in the magazine Nature, to the open skepticism of the publisher, of an experiment reportedly showing that water has a “memory” of the compound last diluted in it, even after dilution is repeated until no molecule of the diluted compound remains. Such a claim would provide scientific support for homeopathy, a treatment based upon such dilution. But Benveniste’s results were irreproducible.
Nature magazine called the idea of water memory “scientifically unacceptable, although this doesn’t yet seem to have affected the commercial success of homeopathy.” Nor perhaps should it.
Homeopathy as a purely medicinal treatment may be harmlessly ineffective, but placebos can be powerful medicine. A skilled, seemingly knowledgeable speaker imparts to someone willing to believe that healing power is within the patient’s grasp. That life-affirming belief alone positively affects the spirit, and thus the immune system. (While adults may self-deceive themselves to health, babies cannot. Treating infants homeopathically has an abysmal track record.) Beyond belief, homeopathy is bunk.
“Homeopathy not only doesn’t work; it couldn’t possibly work. It is inconsistent with our basic knowledge of physics, chemistry, and biology.” ~ American physician Harriet Hall
