Properties of water

Within the Earth's atmosphere and surface, the liquid phase is the most common and is the form that is generally denoted by the word "water". The solid phase of water is known as ice and commonly takes the structure of hard, amalgamated crystals, such as ice cubes, or loosely accumulated granular crystals, like snow. Aside from common hexagonal crystalline ice, other crystalline and amorphous phases of ice are known. The gaseous phase of water is known as water vapor (or steam). Visible steam and clouds are formed from minute droplets of water suspended in the air.

Water also forms a supercritical fluid. The critical temperature is 647 K and the critical pressure is 22.064 MPa. In nature, this only rarely occurs in extremely hostile conditions. A likely example of naturally occurring supercritical water is in the hottest parts of deep water hydrothermal vents, in which water is heated to the critical temperature by volcanic plumes and the critical pressure is caused by the weight of the ocean at the extreme depths where the vents are located. This pressure is reached at a depth of about 2200 meters: much less than the mean depth of the ocean (3800 meters).[27]

Heat of vaporization of water from melting to critical temperature

Water has a very high specific heat capacity of 4184 J/(kg·K) at 20 °C (4182 J/(kg·K) at 25 °C) —the second-highest among all the heteroatomic species (after ammonia), as well as a high heat of vaporization (40.65 kJ/mol or 2257 kJ/kg at the normal boiling point), both of which are a result of the extensive hydrogen bonding between its molecules. These two unusual properties allow water to moderate Earth's climate by buffering large fluctuations in temperature. Most of the additional energy stored in the climate system since 1970 has accumulated in the oceans.[28]

The specific enthalpy of fusion (more commonly known as latent heat) of water is 333.55 kJ/kg at 0 °C: the same amount of energy is required to melt ice as to warm ice from −160 °C up to its melting point or to heat the same amount of water by about 80 °C. Of common substances, only that of ammonia is higher. This property confers resistance to melting on the ice of glaciers and drift ice. Before and since the advent of mechanical refrigeration, ice was and still is in common use for retarding food spoilage.

The specific heat capacity of ice at −10 °C is 2030 J/(kg·K)[29] and the heat capacity of steam at 100 °C is 2080 J/(kg·K).[30]

Density of ice and water as a function of temperature

The density of water is about 1 gram per cubic centimetre (62 lb/cu ft): this relationship was originally used to define the gram.[31] The density varies with temperature, but not linearly: as the temperature increases, the density rises to a peak at 3.98 °C (39.16 °F) and then decreases;[32] the initial increase is unusual because most liquids undergo thermal expansion so that the density only decreases as a function of temperature. The increase observed for water from 0 °C (32 °F) to 3.98 °C (39.16 °F) and for a few other liquids[lower-alpha 4] is described as negative thermal expansion. Regular, hexagonal ice is also less dense than liquid water—upon freezing, the density of water decreases by about 9%.[35][lower-alpha 5]

These peculiar effects are due to the highly directional bonding of water molecules via the hydrogen bonds: ice and liquid water at low temperature have comparatively low-density, low-energy open lattice structures. The breaking of hydrogen bonds on melting with increasing temperature in the range 0–4 °C allows for a denser molecular packing in which some of the lattice cavities are filled by water molecules.[32][36] Above 4 °C, however, thermal expansion becomes the dominant effect,[36] and water near the boiling point (100 °C) is about 4% less dense than water at 4 °C (39 °F).[35][lower-alpha 6]

Under increasing pressure, ice undergoes a number of transitions to other polymorphs with higher density than liquid water, such as ice II, ice III, high-density amorphous ice (HDA), and very-high-density amorphous ice (VHDA).[37][38]

Temperature distribution in a lake in summer and winter

The unusual density curve and lower density of ice than of water is essential for much of the life on earth—if water were most dense at the freezing point, then in winter the cooling at the surface would lead to convective mixing. Once 0 °C are reached, the water body would freeze from the bottom up, and all life in it would be killed.[35] Furthermore, given that water is a good thermal insulator (due to its heat capacity), some frozen lakes might not completely thaw in summer.[35] As it is, the inversion of the density curve leads to a stable layering for surface temperatures below 4 °C, and with the layer of ice that floats on top insulating the water below,[39] even e.g., Lake Baikal in central Siberia freezes only to about 1 m thickness in winter. In general, for deep enough lakes, the temperature at the bottom stays constant at about 4 °C (39 °F) throughout the year (see diagram).[35]

WOA surface density

The density of saltwater depends on the dissolved salt content as well as the temperature. Ice still floats in the oceans, otherwise, they would freeze from the bottom up. However, the salt content of oceans lowers the freezing point by about 1.9 °C[40] (see here for explanation) and lowers the temperature of the density maximum of water to the former freezing point at 0 °C. This is why, in ocean water, the downward convection of colder water is not blocked by an expansion of water as it becomes colder near the freezing point. The oceans' cold water near the freezing point continues to sink. So creatures that live at the bottom of cold oceans like the Arctic Ocean generally live in water 4 °C colder than at the bottom of frozen-over fresh water lakes and rivers.

As the surface of saltwater begins to freeze (at −1.9 °C[40] for normal salinity seawater, 3.5%) the ice that forms is essentially salt-free, with about the same density as freshwater ice. This ice floats on the surface, and the salt that is "frozen out" adds to the salinity and density of the seawater just below it, in a process known as brine rejection. This denser saltwater sinks by convection and the replacing seawater is subject to the same process. This produces essentially freshwater ice at −1.9 °C[40] on the surface. The increased density of the seawater beneath the forming ice causes it to sink towards the bottom. On a large scale, the process of brine rejection and sinking cold salty water results in ocean currents forming to transport such water away from the Poles, leading to a global system of currents called the thermohaline circulation.

Red line shows saturation

Water is miscible with many liquids, including ethanol in all proportions. Water and most oils are immiscible usually forming layers according to increasing density from the top. This can be predicted by comparing the polarity. Water being a relatively polar compound will tend to be miscible with liquids of high polarity such as ethanol and acetone, whereas compounds with low polarity will tend to be immiscible and poorly soluble such as with hydrocarbons.

As a gas, water vapor is completely miscible with air. On the other hand, the maximum water vapor pressure that is thermodynamically stable with the liquid (or solid) at a given temperature is relatively low compared with total atmospheric pressure. For example, if the vapor's partial pressure is 2% of atmospheric pressure and the air is cooled from 25 °C, starting at about 22 °C, water will start to condense, defining the dew point, and creating fog or dew. The reverse process accounts for the fog burning off in the morning. If the humidity is increased at room temperature, for example, by running a hot shower or a bath, and the temperature stays about the same, the vapor soon reaches the pressure for phase change and then condenses out as minute water droplets, commonly referred to as steam.

A saturated gas or one with 100% relative humidity is when the vapor pressure of water in the air is at equilibrium with vapor pressure due to (liquid) water; water (or ice, if cool enough) will fail to lose mass through evaporation when exposed to saturated air. Because the amount of water vapor in the air is small, relative humidity, the ratio of the partial pressure due to the water vapor to the saturated partial vapor pressure, is much more useful. Vapor pressure above 100% relative humidity is called supersaturated and can occur if the air is rapidly cooled, for example, by rising suddenly in an updraft.[lower-alpha 7]

The compressibility of water is a function of pressure and temperature. At 0 °C, at the limit of zero pressure, the compressibility is 5.1×10⁻¹⁰ Pa⁻¹. At the zero-pressure limit, the compressibility reaches a minimum of 4.4×10⁻¹⁰ Pa⁻¹ around 45 °C before increasing again with increasing temperature. As the pressure is increased, the compressibility decreases, being 3.9×10⁻¹⁰ Pa⁻¹ at 0 °C and 100 megapascals (1,000 bar).[41]

The bulk modulus of water is about 2.2 GPa.[42] The low compressibility of non-gasses, and of water in particular, leads to their often being assumed as incompressible. The low compressibility of water means that even in the deep oceans at 4 km depth, where pressures are 40 MPa, there is only a 1.8% decrease in volume.[42]

The bulk modulus of water ice ranges from 11.3 GPa at 0 K up to 8.6 GPa at 273 K.[43] The large change in the compressibility of ice as a function of temperature is the result of its relatively large thermal expansion coefficient compared to other common solids.

The solid/liquid/vapour triple point of liquid water, ice I_h and water vapor in the lower left portion of a water phase diagram.

The temperature and pressure at which ordinary solid, liquid, and gaseous water coexist in equilibrium is a triple point of water. Since 1954, this point had been used to define the base unit of temperature, the kelvin,[44][45] but, starting in 2019, the kelvin is now defined using the Boltzmann constant, rather than the triple point of water.[46]

Due to the existence of many polymorphs (forms) of ice, water has other triple points, which have either three polymorphs of ice or two polymorphs of ice and liquid in equilibrium.[45] Gustav Heinrich Johann Apollon Tammann in Göttingen produced data on several other triple points in the early 20th century. Kamb and others documented further triple points in the 1960s.[47][48][49]

The melting point of ice is 0 °C (32 °F; 273 K) at standard pressure; however, pure liquid water can be supercooled well below that temperature without freezing if the liquid is not mechanically disturbed. It can remain in a fluid state down to its homogeneous nucleation point of about 231 K (−42 °C; −44 °F).[51] The melting point of ordinary hexagonal ice falls slightly under moderately high pressures, by 0.0073 °C (0.0131 °F)/atm[lower-alpha 8] or about 0.5 °C (0.90 °F)/70 atm[lower-alpha 9][52] as the stabilization energy of hydrogen bonding is exceeded by intermolecular repulsion, but as ice transforms into its polymorphs (see crystalline states of ice) above 209.9 MPa (2,072 atm), the melting point increases markedly with pressure, i.e., reaching 355 K (82 °C) at 2.216 GPa (21,870 atm) (triple point of Ice VII[53]).

Pure water containing no exogenous ions is an excellent electronic insulator, but not even "deionized" water is completely free of ions. Water undergoes autoionization in the liquid state when two water molecules form one hydroxide anion (OH⁻
) and one hydronium cation (H
₃O⁺
). Because of autoionization, at ambient temperatures pure liquid water has a similar intrinsic charge carrier concentration to the semiconductor germanium and an intrinsic charge carrier concentration three orders of magnitude greater than the semiconductor silicon, hence, based on charge carrier concentration, water can not be considered to be a completely dielectric material or electrical insulator but to be a limited conductor of ionic charge.[54]

Because water is such a good solvent, it almost always has some solute dissolved in it, often a salt. If water has even a tiny amount of such an impurity, then the ions can carry charges back and forth, allowing the water to conduct electricity far more readily.

It is known that the theoretical maximum electrical resistivity for water is approximately 18.2 MΩ·cm (182 kΩ·m) at 25 °C.[55] This figure agrees well with what is typically seen on reverse osmosis, ultra-filtered and deionized ultra-pure water systems used, for instance, in semiconductor manufacturing plants. A salt or acid contaminant level exceeding even 100 parts per trillion (ppt) in otherwise ultra-pure water begins to noticeably lower its resistivity by up to several kΩ·m.

In pure water, sensitive equipment can detect a very slight electrical conductivity of 0.05501 ± 0.0001 μS/cm at 25.00 °C.[55] Water can also be electrolyzed into oxygen and hydrogen gases but in the absence of dissolved ions this is a very slow process, as very little current is conducted. In ice, the primary charge carriers are protons (see proton conductor).[56] Ice was previously thought to have a small but measurable conductivity of 1×10⁻¹⁰ S/cm, but this conductivity is now thought to be almost entirely from surface defects, and without those, ice is an insulator with an immeasurably small conductivity.[32]

A diagram showing the partial charges on the atoms in a water molecule

An important feature of water is its polar nature. The structure has a bent molecular geometry for the two hydrogens from the oxygen vertex. The oxygen atom also has two lone pairs of electrons. One effect usually ascribed to the lone pairs is that the H–O–H gas-phase bend angle is 104.48°,[57] which is smaller than the typical tetrahedral angle of 109.47°. The lone pairs are closer to the oxygen atom than the electrons sigma bonded to the hydrogens, so they require more space. The increased repulsion of the lone pairs forces the O–H bonds closer to each other. [58]

Another consequence of its structure is that water is a polar molecule. Due to the difference in electronegativity, a bond dipole moment points from each H to the O, making the oxygen partially negative and each hydrogen partially positive. A large molecular dipole, points from a region between the two hydrogen atoms to the oxygen atom. The charge differences cause water molecules to aggregate (the relatively positive areas being attracted to the relatively negative areas). This attraction, hydrogen bonding, explains many of the properties of water, such as its solvent properties.[59]

Although hydrogen bonding is a relatively weak attraction compared to the covalent bonds within the water molecule itself, it is responsible for several of the water's physical properties. These properties include its relatively high melting and boiling point temperatures: more energy is required to break the hydrogen bonds between water molecules. In contrast, hydrogen sulfide (H
₂S), has much weaker hydrogen bonding due to sulfur's lower electronegativity. H
₂S is a gas at room temperature, despite hydrogen sulfide having nearly twice the molar mass of water. The extra bonding between water molecules also gives liquid water a large specific heat capacity. This high heat capacity makes water a good heat storage medium (coolant) and heat shield.

Dew drops adhering to a spider web

Water molecules stay close to each other (cohesion), due to the collective action of hydrogen bonds between water molecules. These hydrogen bonds are constantly breaking, with new bonds being formed with different water molecules; but at any given time in a sample of liquid water, a large portion of the molecules are held together by such bonds.[60]

Water also has high adhesion properties because of its polar nature. On clean, smooth glass the water may form a thin film because the molecular forces between glass and water molecules (adhesive forces) are stronger than the cohesive forces. In biological cells and organelles, water is in contact with membrane and protein surfaces that are hydrophilic; that is, surfaces that have a strong attraction to water. Irving Langmuir observed a strong repulsive force between hydrophilic surfaces. To dehydrate hydrophilic surfaces—to remove the strongly held layers of water of hydration—requires doing substantial work against these forces, called hydration forces. These forces are very large but decrease rapidly over a nanometer or less.[61] They are important in biology, particularly when cells are dehydrated by exposure to dry atmospheres or to extracellular freezing.[62]

Rain water flux from a canopy. Among the forces that govern drop formation: Surface tension, Cohesion (chemistry), Van der Waals force, Plateau–Rayleigh instability.

This paper clip is under the water level, which has risen gently and smoothly. Surface tension prevents the clip from submerging and the water from overflowing the glass edges.

Temperature dependence of the surface tension of pure water

Water has an unusually high surface tension of 71.99 mN/m at 25 °C[63] which is caused by the strength of the hydrogen bonding between water molecules.[64] This allows insects to walk on water.[64]

Because water has strong cohesive and adhesive forces, it exhibits capillary action.[65] Strong cohesion from hydrogen bonding and adhesion allows trees to transport water more than 100 m upward.[64]

Presence of colloidal calcium carbonate from high concentrations of dissolved lime turns the water of Havasu Falls turquoise.

Water is an excellent solvent due to its high dielectric constant.[66] Substances that mix well and dissolve in water are known as hydrophilic ("water-loving") substances, while those that do not mix well with water are known as hydrophobic ("water-fearing") substances.[67] The ability of a substance to dissolve in water is determined by whether or not the substance can match or better the strong attractive forces that water molecules generate between other water molecules. If a substance has properties that do not allow it to overcome these strong intermolecular forces, the molecules are precipitated out from the water. Contrary to the common misconception, water and hydrophobic substances do not "repel", and the hydration of a hydrophobic surface is energetically, but not entropically, favorable.

When an ionic or polar compound enters water, it is surrounded by water molecules (hydration). The relatively small size of water molecules (~ 3 angstroms) allows many water molecules to surround one molecule of solute. The partially negative dipole ends of the water are attracted to positively charged components of the solute, and vice versa for the positive dipole ends.

In general, ionic and polar substances such as acids, alcohols, and salts are relatively soluble in water, and nonpolar substances such as fats and oils are not. Nonpolar molecules stay together in water because it is energetically more favorable for the water molecules to hydrogen bond to each other than to engage in van der Waals interactions with non-polar molecules.

An example of an ionic solute is table salt; the sodium chloride, NaCl, separates into Na⁺
cations and Cl⁻
anions, each being surrounded by water molecules. The ions are then easily transported away from their crystalline lattice into solution. An example of a nonionic solute is table sugar. The water dipoles make hydrogen bonds with the polar regions of the sugar molecule (OH groups) and allow it to be carried away into solution.

The quantum tunneling dynamics in water was reported as early as 1992. At that time it was known that there are motions which destroy and regenerate the weak hydrogen bond by internal rotations of the substituent water monomers.[68] On 18 March 2016, it was reported that the hydrogen bond can be broken by quantum tunneling in the water hexamer. Unlike previously reported tunneling motions in water, this involved the concerted breaking of two hydrogen bonds.[69] Later in the same year, the discovery of the quantum tunneling of water molecules was reported.[70]

Water is relatively transparent to visible light, near ultraviolet light, and far-red light, but it absorbs most ultraviolet light, infrared light, and microwaves. Most photoreceptors and photosynthetic pigments utilize the portion of the light spectrum that is transmitted well through water. Microwave ovens take advantage of water's opacity to microwave radiation to heat the water inside of foods. Water's light blue colour is caused by weak absorption in the red part of the visible spectrum.[3][71]

The repulsive effects of the two lone pairs on the oxygen atom cause water to have a bent, not linear, molecular structure,[73] allowing it to be polar. The hydrogen–oxygen–hydrogen angle is 104.45°, which is less than the 109.47° for ideal sp³ hybridization. The valence bond theory explanation is that the oxygen atom's lone pairs are physically larger and therefore take up more space than the oxygen atom's bonds to the hydrogen atoms.[74] The molecular orbital theory explanation (Bent's rule) is that lowering the energy of the oxygen atom's nonbonding hybrid orbitals (by assigning them more s character and less p character) and correspondingly raising the energy of the oxygen atom's hybrid orbitals bonded to the hydrogen atoms (by assigning them more p character and less s character) has the net effect of lowering the energy of the occupied molecular orbitals because the energy of the oxygen atom's nonbonding hybrid orbitals contributes completely to the energy of the oxygen atom's lone pairs while the energy of the oxygen atom's other two hybrid orbitals contributes only partially to the energy of the bonding orbitals (the remainder of the contribution coming from the hydrogen atoms' 1s orbitals).

In liquid water there is some self-ionization giving hydronium ions and hydroxide ions.

2 H
₂O ⇌ H
₃O⁺
+ OH⁻

The equilibrium constant for this reaction, known as the ionic product of water, ${\displaystyle K_{\rm {w}}=[{\rm {H_{3}O^{+}}}][{\rm {OH^{-}}}]}$, has a value of about 10⁻¹⁴ at 25 °C. At neutral pH, the concentration of the hydroxide ion (OH⁻
) equals that of the (solvated) hydrogen ion (H⁺
), with a value close to 10⁻⁷ mol L⁻¹ at 25 °C.[75] See data page for values at other temperatures.

The thermodynamic equilibrium constant is a quotient of thermodynamic activities of all products and reactants including water:

${\displaystyle K_{\rm {eq}}={\frac {a_{\rm {H_{3}O^{+}}}\cdot a_{\rm {OH^{-}}}}{a_{\rm {H_{2}O}}^{2}}}}$

However for dilute solutions, the activity of a solute such as H₃O⁺ or OH⁻ is approximated by its concentration, and the activity of the solvent H₂O is approximated by 1, so that we obtain the simple ionic product ${\displaystyle K_{\rm {eq}}\approx K_{\rm {w}}=[{\rm {H_{3}O^{+}}}][{\rm {OH^{-}}}]}$

The action of water on rock over long periods of time typically leads to weathering and water erosion, physical processes that convert solid rocks and minerals into soil and sediment, but under some conditions chemical reactions with water occur as well, resulting in metasomatism or mineral hydration, a type of chemical alteration of a rock which produces clay minerals. It also occurs when Portland cement hardens.

Water ice can form clathrate compounds, known as clathrate hydrates, with a variety of small molecules that can be embedded in its spacious crystal lattice. The most notable of these is methane clathrate, 4 CH
₄·23H
₂O, naturally found in large quantities on the ocean floor.

Rain is generally mildly acidic, with a pH between 5.2 and 5.8 if not having any acid stronger than carbon dioxide.[76] If high amounts of nitrogen and sulfur oxides are present in the air, they too will dissolve into the cloud and raindrops, producing acid rain.

Water is amphoteric: it has the ability to act as either an acid or a base in chemical reactions.[84] According to the Brønsted-Lowry definition, an acid is a proton (H⁺
) donor and a base is a proton acceptor.[85] When reacting with a stronger acid, water acts as a base; when reacting with a stronger base, it acts as an acid.[85] For instance, water receives an H⁺
ion from HCl when hydrochloric acid is formed:

HCl
(acid) + H
₂O
(base) ⇌ H
₃O⁺
+ Cl⁻

In the reaction with ammonia, NH
₃, water donates a H⁺
ion, and is thus acting as an acid:

NH
₃
(base) + H
₂O
(acid) ⇌ NH⁺
₄ + OH⁻

Because the oxygen atom in water has two lone pairs, water often acts as a Lewis base, or electron-pair donor, in reactions with Lewis acids, although it can also react with Lewis bases, forming hydrogen bonds between the electron pair donors and the hydrogen atoms of water. HSAB theory describes water as both a weak hard acid and a weak hard base, meaning that it reacts preferentially with other hard species:

H⁺

(Lewis acid) + H
₂O
(Lewis base) → H
₃O⁺

Fe³⁺

(Lewis acid) + H
₂O
(Lewis base) → Fe(H
₂O)³⁺
₆

Cl⁻

(Lewis base) + H
₂O
(Lewis acid) → Cl(H
₂O)⁻
₆

When a salt of a weak acid or of a weak base is dissolved in water, water can partially hydrolyze the salt, producing the corresponding base or acid, which gives aqueous solutions of soap and baking soda their basic pH:

Na
₂CO
₃ + H
₂O ⇌ NaOH + NaHCO
₃

Water's Lewis base character makes it a common ligand in transition metal complexes, examples of which include metal aquo complexes such as Fe(H
₂O)²⁺
₆ to perrhenic acid, which contains two water molecules coordinated to a rhenium center. In solid hydrates, water can be either a ligand or simply lodged in the framework, or both. Thus, FeSO
₄·7H
₂O consists of [Fe₂(H₂O)₆]²⁺ centers and one "lattice water". Water is typically a monodentate ligand, i.e., it forms only one bond with the central atom.[86]

Some hydrogen-bonding contacts in FeSO₄^.7H₂O. This metal aquo complex crystallizes with one molecule of "lattice" water, which interacts with the sulfate and with the [Fe(H₂O)₆]²⁺ centers.

As a hard base, water reacts readily with organic carbocations; for example in a hydration reaction, a hydroxyl group (OH⁻
) and an acidic proton are added to the two carbon atoms bonded together in the carbon-carbon double bond, resulting in an alcohol. When the addition of water to an organic molecule cleaves the molecule in two, hydrolysis is said to occur. Notable examples of hydrolysis are the saponification of fats and the digestion of proteins and polysaccharides. Water can also be a leaving group in S_N2 substitution and E2 elimination reactions; the latter is then known as a dehydration reaction.

Water contains hydrogen in the oxidation state +1 and oxygen in the oxidation state −2.[87] It oxidizes chemicals such as hydrides, alkali metals, and some alkaline earth metals.[88][89] One example of an alkali metal reacting with water is:[90]

2 Na + 2 H
₂O → H
₂ + 2 Na⁺
+ 2 OH⁻

Some other reactive metals, such as aluminum and beryllium, are oxidized by water as well, but their oxides adhere to the metal and form a passive protective layer.[91] Note that the rusting of iron is a reaction between iron and oxygen[92] that is dissolved in water, not between iron and water.

Water can be oxidized to emit oxygen gas, but very few oxidants react with water even if their reduction potential is greater than the potential of O
₂/H
₂O. Almost all such reactions require a catalyst.[93] An example of the oxidation of water is:

4 AgF
₂ + 2 H
₂O → 4 AgF + 4 HF + O
₂

Water can be split into its constituent elements, hydrogen, and oxygen, by passing an electric current through it.[94] This process is called electrolysis. The cathode half reaction is:

2 H⁺
+ 2
e⁻
→ H
₂

The anode half reaction is:

2 H
₂O → O
₂ + 4 H⁺
+ 4
e⁻

The gases produced bubble to the surface, where they can be collected or ignited with a flame above the water if this was the intention. The required potential for the electrolysis of pure water is 1.23 V at 25 °C.[94] The operating potential is actually 1.48 V or higher in practical electrolysis.