Properties of water
Water (H 2O) is a polar inorganic chemical that is a tasteless and odorless liquid that is almost colorless except for a slight tint of blue at room temperature.
It is known as the "universal solvent" and the "solvent of life," and it is by far the most studied chemical compound.
It is the most abundant substance on Earth's surface and the only common substance that may be found as a solid, liquid, or gas.
It is also the universe's third most abundant chemical (behind molecular hydrogen and carbon monoxide) Water molecules are highly polar and establish hydrogen bonds with one another.
Because of its polarity, it can dissociate ions in salts and link to other polar chemicals like alcohols and acids, allowing them to dissolve.
It has numerous unique properties due to its hydrogen bonding, including a solid form that is less dense than its liquid form, a boiling point of 100 °C for its molar mass, and a large heat capacity.
Water is amphoteric, which means it may function as an acid or a base depending on the pH of the solution it's in; it swiftly forms both H+ and OH ions.
Because of its amphoteric nature, it undergoes self-ionization.
Because the product of the activities, or roughly, the concentrations of H+ and OH, is a constant, their concentrations are inversely proportional to one another.
Water is a chemical compound with the formula H 2O, in which two hydrogen atoms are covalently bound to a single oxygen atom in each molecule.
At room temperature and pressure, water is a tasteless, odorless liquid.
Water exhibits modest absorption bands at wavelengths around 750 nm, giving it a blue color.
This is immediately visible in a water-filled bath or washbasin with a white lining.
Ice crystals larger than a grain of sand, such as those found in glaciers, appear blue as well.
Water is largely a liquid under normal conditions, unlike other comparable hydrides of the oxygen family, which are mostly gaseous.
Hydrogen bonding is responsible for water's distinctive features.
Water molecules are constantly moving about one another, and hydrogen bonds are constantly breaking and reforming at speeds faster than 200 femtoseconds (2 1013 seconds).
These linkages, on the other hand, are strong enough to provide water many of its special qualities, some of which make it essential to life.
Water, ice, and vapor
The liquid phase is the most frequent in the Earth's atmosphere and surface, and it is the form that is commonly signified by the word "water".
Ice is a solid phase of water that can be hard, amalgamated crystals like ice cubes or loosely collected granular crystals like snow.
Other crystalline and amorphous phases of ice exist in addition to the conventional hexagonal crystalline ice.
Water vapor is the liquid phase of water that is gaseous (or steam).
Miniscule droplets of water floating in the air generate visible steam and clouds.
Water is also a supercritical fluid.
The critical pressure is 22.064 MPa and the critical temperature is 647 K.
In nature, this only happens in the most extreme of circumstances.
The hottest regions of deep water hydrothermal vents, where water is heated to the critical temperature by volcanic plumes and the critical pressure are created by the weight of the ocean at the extreme depths where the vents are located, are plausible examples of naturally occurring supercritical water.
This pressure is attained at a depth of roughly 2200 meters, which is substantially lower than the ocean's average depth (3800 meters).
Water has a high specific heat capacity of 4184 J/(kg) at 25 °C, which is the second-highest among all heteroatomic species (after ammonia), as well as the high heat of vaporization (40.65 kJ/mol or 2257 kJ/kg at the normal boiling point), both of which are due to extensive hydrogen bonding between its molecules.
Water's peculiar qualities allow it to manage Earth's climate by buffering massive temperature variations.
Since 1970, the seas have acquired the majority of the excess energy stored in the climate system.
At 0 °C, water's specific enthalpy of fusion (also known as latent heat) is 333.55 kJ/kg: melting ice requires the same amount of energy as warming ice from 160 °C to its melting point or heating the same amount of water by around 80 °C.
Only ammonia has a higher concentration than the rest of the common chemicals.
On glacier and drift ice, this feature confers resistance to melting.
Ice was and still is widely used to keep food from spoiling before and after the invention of artificial refrigeration.
Ice has a particular heat capacity of 2030 J/(kg) at 10 °C, while steam has a heat capacity of 2080 J/(kg) at 100 °C.
Water has a density of roughly 1 gram per cubic centimeter (62 lb/cu ft), which was used to define the gram at first.
The density varies with temperature, but not in a linear manner: as the temperature rises, the density climbs to a peak of 3.98 °C (39.16 °F) and then falls; this is rare.
Regular hexagonal ice is similarly less dense than liquid water; water loses around 9% of its density when it freezes.
These effects result from the reduction of thermal motion caused by cooling, which allows water molecules to establish more hydrogen bonds, preventing them from colliding.
While the breakdown of hydrogen bonds caused by heating permits water molecules to pack closer together despite increased thermal motion (which tends to expand a liquid) in the 0-4 °C range, water expands when the temperature rises above 4 °C.
The water around 4 °C (39 °F) is roughly 4% less dense than water near the boiling point.
Under increased pressure, ice transforms into ice II, ice III, high-density amorphous ice (HDA), and very-high-density amorphous ice, all of which have a higher density than liquid water (VHDA).
If water were most dense near the freezing point, the very cold water at the surface of lakes and other bodies of water would sink in the winter, lakes would freeze from the bottom up, and all life in them would perish.
Additionally, because water is a good thermal insulator (owing to its heat capacity), certain frozen lakes may not thaw completely in the summer.
The water below is insulated by the coating of ice that floats on top.
Water at a temperature of around 4 °C (39 °F) sinks to the bottom, maintaining the temperature of the water at the bottom (see diagram).
The density of salt water is determined by the amount of dissolved salt and the temperature.
The oceans still have ice floating in them; otherwise, they would freeze from the bottom up.
The salt content of oceans, on the other hand, lowers the freezing point by around 1.9°C (see here for an explanation) and brings the temperature of the density maximum of water back to the original freezing point of 0°C.
This is why, in ocean water, the downward convection of colder water is not obstructed by the expansion of water as the temperature approaches the freezing point.
The cold water around the freezing point of the oceans continues to sink.
As a result, organisms that live at the bottom of freezing oceans like the Arctic Ocean live in water that is 4 degrees Celsius colder than those that reside at the bottom of frozen lakes and rivers.
The ice that develops when the surface of saltwater freezes (around 1.9 °C for normal salinity seawater, 3.5 percent), is practically salt-free and has nearly the same density as freshwater ice.
In a process known as brine rejection, the salt that is "frozen out" of the ice floats on the surface, adding to the salinity and density of the seawater just below it.
Convection causes the denser saltwater to sink, and the replacement seawater follows suit.
At 1.9 °C, this generates virtually freshwater ice on the surface.
The growing ice sinks to the bottom due to the increased density of the seawater beneath it.
On a wide scale, the process of brine rejection and sinking cold salty water causes ocean currents to form to move such water away from the poles, resulting in the thermohaline circulation, a global system of currents.
Water, being a moderately polar substance, will tend to be miscible with high-polarity liquids such as ethanol and acetone, whereas low-polarity chemicals, such as hydrocarbons, will likely be immiscible and weakly soluble.
Water vapor is entirely miscible with air as a gas.
In comparison to total air pressure, the greatest water vapor pressure that is thermodynamically stable with the liquid (or solid) at a given temperature is quite low.
For example, if the vapor's partial pressure is 2% of atmospheric pressure and the air is cooled from 25 °C to 22 °C, water will begin to condense at around 22 °C, defining the dew point and resulting in fog or dew.
The fog dissipates in the morning due to the reverse process.
When the humidity in a room is raised, such as by taking a hot shower or bath, while the temperature remains constant, the vapor quickly exceeds the pressure required for phase transition and condenses out as minute water droplets, generally referred to as steam.
When water (or ice, if cool enough) is exposed to saturated air, the vapor pressure of water in the air is at equilibrium with the vapor pressure owing to (liquid) water; water (or ice, if cool enough) will not lose mass through evaporation.
Because the amount of water vapor in the air is so little, relative humidity (the ratio of water vapor partial pressure to saturated partial vapor pressure) is far more useful.
Super-saturated vapor pressure is defined as vapor pressure above 100 percent relative humidity, which can occur when air is abruptly cooled, for as by ascending suddenly in an updraft.
The compressibility at 0 °C, at the limit of zero pressure, is 5.11010 Pa1.
Around 45 °C, the compressibility hits a minimum of 4.41010 Pa1 near the zero-pressure limit, before increasing with increasing temperature.
Compressibility reduces with increasing pressure, with 3.91010 Pa1 at 0 °C and 100 megapascals (1,000 bar).
Because water has a low compressibility, even at 4 km depth, where pressures are 40 MPa, there is only a 1.8 percent reduction in volume.
It can stay fluid until it reaches its homogenous nucleation point of 231 K (42 °C; 44 °F).
As the stabilization energy of hydrogen bonding is exceeded by intermolecular repulsion, the melting point of ordinary hexagonal ice falls slightly under moderately high pressures, by 0.0073 °C (0.0131 °F)/atm or about 0.5 °C (0.90 °F)/70 atm, but as ice transforms into its polymorphs (see crystalline states of ice) above 209.9 MPa (2,072 at (triple point of Ice VII).
Although pure water with no external ions is a superb electrical insulator, even "deionized" water contains ions.
In the liquid state, auto-ionization occurs when two water molecules combine to generate one hydroxide anion (OH) and one hydronium cation (H 3O+).
Because of auto-ionization, pure liquid water has a similar intrinsic charge carrier concentration to the semiconductor germanium and a three-order-of-magnitude greater intrinsic charge carrier concentration than the semiconductor silicon at ambient temperatures.
As a result, water cannot be considered a completely dielectric material or electrical insulator, but rather a limited conductor of ionic charge.
Water is such a good solvent that it almost always has some sort of solute dissolved in it, most commonly a salt.
If even a trace of such impurity is present in water, the ions can transfer charges back and forth, allowing the water to conduct electricity more efficiently.
At 25 degrees Celsius, water has a theoretical maximum electrical resistance of 18.2 Mcm (182 km).
This figure matches what is commonly found in reverse osmosis, ultra-filtered, and deionized ultra-pure water systems used in semiconductor manufacturing factories, for example.
In otherwise ultra-pure water, a salt or acid impurity level over 100 parts per trillion (ppt) begins to substantially diminish its resistivity by up to several km.
At 25.00 °C, sensitive equipment may detect an electrical conductivity of 0.05501 0.0001 S/cm in pure water.
Water can also be electrolyzed into oxygen and hydrogen gases, although this is a very slow process in the absence of dissolved ions since the very little current is conducted.
Protons are the major charge carriers in ice (see proton conductor).
Ice was once thought to have a minuscule but measurable conductivity of 11010 S/cm, however that conductivity is now thought to be almost completely due to surface imperfections, and without them, ice is an insulator with impossibly low conductivity.