Water has the chemical formula H2O, making it an inorganic substance. It is the primary chemical component of the Earth’s hydrosphere and the fluids of all known living things (in which it serves as a solvent. It is translucent, flavourless, odourless, and almost colourless. In spite of not supplying food, energy, or organic micronutrients, it is essential for all known forms of life. Its molecules are made up of two hydrogen atoms joined by covalent bonds and have the chemical formula H2O. The angle at which the hydrogen atoms are joined to the oxygen atom is 104.45°. The liquid condition of H2O at normal pressure and temperature is known as “water” as well.


Water occurs because the environment on Earth is pretty near to the triple point of water.

Seas and oceans account for the majority (approximately 96.5%) of the planet’s total water volume, which makes up roughly 71% of its surface.[4] There are negligible amounts of water in the groundwater (1.7%), glaciers and ice caps of Antarctica and Greenland (1.7%), clouds (which are made up of ice and liquid water suspended in air), and precipitation (0.001%) of the atmosphere. The water cycle, which includes evaporation, transpiration (evapotranspiration), condensation, precipitation, and runoff, is a continuous process that typically results in water reaching the sea.

The global economy depends heavily on water. Agriculture uses over 70% of the freshwater that people use For many regions of the world, fishing in salt and fresh water bodies has been and still is a significant source of sustenance.

Properties of Water

Water (H2O) is a polar inorganic chemical that is almost colourless at normal temperature with the exception of a little blue undertone. It is referred to as the “universal solvent” and the “solvent of life” and is by far the chemical compound that has been investigated the most It is the most prevalent material on Earth’s surface and the only one that can be found there as a solid, liquid, and gas at the same time. Aside from carbon monoxide and molecular hydrogen, it is the third most prevalent molecule in the cosmos.

Water molecules are highly polar and form hydrogen bonds with one another. Due to its polarity, it may dissolve other polar chemicals like alcohols and acids by bonding with them and dissociating the ions in salts. Its several distinctive characteristics, including possessing a solid form that is less dense than its liquid form, a comparatively high boiling point of 100 °C for its molar mass, and a high heat capacity, are all brought about by its hydrogen bonding. Water is amphoteric, which means that depending on the pH of the solution it is in, it may display traits of an acid or a base; it quickly creates both H+ and OH ions.[c] It goes through a process of self-ionization due to its amphoteric nature. The concentrations of H+ and OH are inversely proportional to one another because the sum of their activities, or roughly, their concentrations, is a constant.

Physical Properties

Water is a chemical substance having the molecular formula H 2O. Two hydrogen atoms are covalently bound to one oxygen atom in each water molecule. At room temperature and pressure, water has no taste or dour. Liquid water appears to be blue because of modest absorption bands at wavelengths of around 750 nm. This is clearly visible in a white washbasin or bathtub that is full with water. Glaciers and other large ice crystals also have a blue appearance.

Physical Properties of water

Water is essentially a liquid under normal circumstances, in contrast to other comparable hydrides of the oxygen family, which are often gaseous. Water has this special quality because of hydrogen bonding. The hydrogen atoms in water molecules move around each other continually.

Water, Ice, and Vapor

The liquid phase is the most prevalent and the form that is often indicated by the word “water” inside the Earth’s atmosphere and surface. The term “ice” refers to the solid form of water, which often has the shape of ice cubes or loosely piled granular crystals (like snow). Other crystalline and amorphous phases of ice are known in addition to the typical hexagonal crystalline ice. Water vapour (sometimes known as steam) is the name for the gaseous state of water. Water droplets that are so small that are suspended in the air give forth visible steam and clouds.

Water may create a supercritical fluid as well. The critical pressure is 22.064 MPa and the critical temperature is 647 K. In nature, this only sometimes happens under highly adverse circumstances. The hottest areas of deep water hydrothermal vents, where water is heated to the critical temperature by volcanic plumes and the critical pressure is brought on by the weight of the ocean at the extreme depths where the vents are located, are likely examples of naturally occurring supercritical water. This pressure is attained at a depth of around 2200 meters, which is much less than the ocean’s typical depth of 3800 meters.

Density of Saltwater and Ice

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 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.

Density of Saltwater and Ice

The ice that develops as saltwater begins to freeze (around 1.9 °C for seawater with a typical salinity, 3.5%) is basically salt-free and has a density similar to that of freshwater ice. In a process known as brine rejection, the salt that is “frozen out” of this ice and “floats on the surface” increases the salinity and density of the saltwater immediately underneath it. The replacement seawater goes through the same procedure as the sinking denser saltwater thanks to convection. At 1.9 °C, this practically creates freshwater ice on the surface. The growing ice sinks towards the bottom due to the seawater’s increasing density. Ocean currents are produced on a big scale by the process of brine rejection and sinking cold salty water.

Compressibility of Water

Pressure and temperature have an impact on how compressible water is. The compressibility at zero degrees Celsius and zero atmosphere pressure is 5.1 1010 Pa. The compressibility decreases with increasing temperature until it reaches a minimum of 4.41010 Pa1 at the zero-pressure limit, which is about 45 °C. The compressibility, which is 3.91010 Pa1 at 0 °C and 100 megapascals (1,000 bar), diminishes as pressure increases.

Water has a bulk modulus of roughly 2.2 GA. Water in particular has a poor compressibility, which makes it a common misconception that non-gasses are incompressible. Due to water’s poor compressibility, even at pressures of 40 MP and depths of 4 km, there is only a 1.8% reduction in volume.

Triple point

A triple point of water is the combination of temperature and pressure when ordinary solid, liquid, and gaseous water coexist in equilibrium. The triple point of water had been used since 1954 to define the kelvin, the fundamental unit of temperature, but as of 2019, the kelvin is now defined using the Boltzmann constant.

Water has various triple points that include either three polymorphs of ice or two polymorphs of ice and liquid in equilibrium since there are multiple polymorphs (forms) of ice. Early in the 20th century, Gottingen’s Gustav Heinrich Johann Apollon Taman published data on a number of additional triple points. In the 1960s, Kamba and others reported more triple points.

Melting point

Melting point

Pure liquid water may be supercooled much below that temperature without freezing if the liquid is not physically disturbed. Ice has a melting point of 0 °C (32 °F; 273 K) under standard pressure. Up to its homogeneous nucleation point, which is at 231 K (42 °C; 44 °F), it can continue to exist in a fluid condition. Under moderately high pressures, the melting point of common hexagonal ice decreases by 0.0073 °C (0.0131 °F)/atm[h] or roughly 0.5 °C (0.90 °F)/70 atm[i].[52] However, as ice changes into various polymorphs (see crystalline phases of ice) at 209.9 MPa (2,072 atm), the melting point increases noticeably with pressure, reaching 355 K (82 °C) at 2.216 atm. This is because the stabilisation energy of hydrogen bonding is exceeded by intermolecular repulsion.

Electrical Properties

Electrical conductivity

Even “deionized” water does not include all exogenous ions, making pure water an effective electrical insulator. When two water molecules combine to generate one hydroxide anion (OH) and one hydronium cation (H 3O+), water becomes autoionized.

Due to autoionization, pure liquid water at room temperature has an intrinsic charge carrier concentration that is three orders of magnitude higher than that of semiconductor silicon and similar to that of semiconductor germanium. As a result, water cannot be considered to be a completely electrical insulator or dielectric material, but rather a limited conductor of ionic charge. It is well known that water has a maximum theoretical electrical resistance of about 18.2 Mcm (182 km) at 25 °C. This result is consistent with what is frequently observed on reverse osmosis, ultra-filtered, and deionized ultra-pure water systems, such as those used in semiconductor production facilities. In otherwise ultra-pure water, a salt or acid contamination level of even 100 parts per trillion (ppt) starts to noticeably reduce its resistivity by up to several km. Sensitive equipment may pick up a very faint electrical conductivity of 0.05501 0.0001 S/cm at 25 °C in pure water. While water can also electrolyze into oxygen and hydrogen gases, this process moves very slowly and conducts very little current in the absence of dissolved ions. Protons are the major charge carriers in ice (see proton conductor). Ice was once supposed to have a conductivity of 11010 S/cm, which was thought to be modest but observable. However, it is now believed that this conductivity is virtually completely due to surface flaws, and that without them, ice would be an insulator with an impossibly small conductivity.

Polarity and Hydrogen Bonding

The polar nature of water is a significant characteristic. The two hydrogen atoms from the oxygen vertex have a bent molecular geometry in the structure. Two electron lone pairs are also present in the oxygen atom. The H-O-H gas-phase bend angle is 104.48°, which is less than the average tetrahedral angle of 109.47°. This is one consequence that is typically attributed to the lone pairs. The lone pairs need more room because they are physically closer to the oxygen atom than the electrons that are sigma-bonded to the hydrogens. The O-H bonds are drawn closer to one another by the enhanced repulsion of the lone pairs. Water’s polar molecular status is another effect of its structure. A bond dipole moment points from each H to the O as a result of the difference in electronegativity, making the oxygen partially negative and each hydrogen partially positive. Between the two hydrogen atoms, a significant molecular dipole extends towards the oxygen atom. Water molecules group together as a result of the charge imbalances (the relatively positive areas are drawn to the comparatively negative areas). Many of the characteristics of water, including its solvent properties, are explained by this attraction, known as hydrogen bonding.

Several of the physical characteristics of water are due to hydrogen bonding, despite the attraction being relatively weak compared to the covalent bonds within the water molecule. One of these characteristics is that water has relatively high melting and boiling points, which means it takes more energy to break the hydrogen bonds that connect the molecules of water. As a result of sulfur’s poorer electronegativity, hydrogen sulphide (H 2S) possesses substantially weaker hydrogen bonds. Despite having nearly twice the molar mass of water, hydrogen sulphide is a gas at room temperature. Liquid water has a high specific heat capacity due to the additional bonds between water molecules. Water is a suitable heat storage medium (coolant) and heat barrier due to its large heat capacity.

Cohesion and Adhesion

Hydrogen bonds between water molecules work together to keep water molecules close to one another (cohesion). While a significant portion of the molecules in a sample of liquid water are held together by hydrogen bonds at any given time, these bonds are constantly breaking and new ones are being formed with various water molecules.

Cohesion and Adhesion

Water’s polar nature lends it great adhesion abilities as well. Because the molecular forces (adhesive forces) between glass and water molecules are stronger than the cohesive forces, water may form a thin film on clean, smooth glass.[Reference required] Water comes into touch with membrane and protein surfaces in biological cells and organelles. Hydrophilic surfaces are those that have a significant affinity to water. A potent repelling force between hydrophilic surfaces was discovered by Irving Langmuir. Dehydrating hydrophilic surfaces necessitates exerting significant effort against these so-called hydration forces in order to remove the firmly adhered layers of water of hydration. Over a nanometer or less, these forces rapidly diminish from their extremely high levels.

Electromagnetic Absorption

Water absorbs the majority of microwaves, near ultraviolet light, and far-red light, although it is reasonably transparent to visible light, near ultraviolet light, and far-red light. The majority of photoreceptors and pigments used for photosynthetic processes rely on the part of the light spectrum that is well-transmitted through water. In order to heat the water inside the food, microwave ovens take use of water’s opacity to microwave radiation. Weak absorption in the visible red spectrum is what gives water its pale blue color.

A single water molecule can participate in a maximum of four hydrogen bonds because it can accept two bonds using the lone pairs on oxygen and donate two hydrogen atoms. Other molecules like hydrogen fluoride, ammonia, and methanol can also form hydrogen bonds. However, they do not show anomalous thermodynamic, kinetic, or structural properties like those observed in water because none of them can form four hydrogen bonds: either they cannot donate or accept hydrogen atoms, or there are steric effects in bulky residues. In water, intermolecular tetrahedral structures form due to the four hydrogen bonds, thereby forming an open structure and a three-dimensional bonding network, resulting in the anomalous decrease in density when cooled below 4 °C.

Molecular Structure

Due to the oxygen atom’s attraction to the two lone pairs, water has a bent molecular shape rather than a linear one, which makes it polar. The optimal sp3 hybridization angle is 109.47°, however the hydrogen-oxygen-hydrogen angle is lower at 104.45°. According to the valence bond hypothesis, the oxygen atom’s lone pairs take up more space than its bonds with hydrogen atoms because they are physically bigger. According to the molecular orbital theory explanation (Bent’s rule), increasing the energy of the oxygen atom’s hybrid orbitals bonded to the hydrogen atoms while decreasing the energy of the oxygen atom’s nonbonding hybrid orbitals has the overall effect of lowering the energy of the occupied molecular orbitals. This is because the energy of the oxygen atom’s nonboning hybrid orbitals is lower than that of the hydrogen atom’s hybrid orbitals.

Chemical Properties


Self-ionization occurs in liquid water, producing hydronium and hydroxide ions.

2 H 2O ⇌ H

3O+ + OH−
The ionic product of water is the reaction’s equilibrium constant, and its formula is w = [H 3 O + ].
[ O H − ]
K_rm w=[rm H_3O+][{{\rm {{OH^{-}}}}}], which is approximately 1014 at 25 °C. With a value close to 107 mol L1 at 25 °C, the concentration of the hydroxide ion (OH) at neutral pH is equal to that of the (solvated) hydrogen ion (H+). For values at various temperatures, see the data page.

The thermodynamic equilibrium constant is a ratio of all products’ and reactants’ thermodynamic activity, including water:

� e

q = � H 3 O + ⋅ � O H − � H 2 O 2 K_”rm eq” is equal to “frac a_a_rm a_a_rm a_a_rm a_a_rm a_a_rm a_a_rm a_a_rm a_a_rm a_a_rm a_{a_{{{\rm {{H_{2}O}}}}}^{2}}}
But in diluted solutions, the activity of a solute like H3O+ or OH is roughly correlated with its concentration, and the activity of the solvent H2O is roughly correlated with 1, resulting in the simple ionic product.
� e
q ≈
� w = [ H 3 O + ]
[ O H − ]
Displaystyle K_rm eq estimate K_rm w = [rm H_3O + K_rm][{\rm {OH^{-}}}]}


In some instances, chemical reactions with water result in mineral hydration, also known as metasomatism, which is a type of chemical alteration of a rock that produces clay minerals. Weathering and water erosion are physical processes that physically transform solid rocks and minerals into soil and sediment. It also happens as Portland cement dries out.

The wide-open crystal lattice of water ice may accommodate a variety of tiny molecules that can form clathrate compounds, sometimes known as clathrate hydrates. Methane clathrate, 4 CH 423H, is the most prominent of them.

2O, a substance that is abundantly present in nature on ocean floors.

Acidity in nature

pH of water

If there is no acid stronger than carbon dioxide, rain typically has a pH between 5.2 and 5.8. Acid rain will result from the high levels of nitrogen and Sulphur oxides in the atmosphere dissolving into the clouds and droplets.


There are numerous known isotopologues of water that result from the existence of various hydrogen and oxygen isotopes. The current global standard for water isotopes is Vienna Standard Mean Ocean Water. The neutron-less hydrogen isotope protium makes up practically all naturally occurring water. Less than 20 parts per quintillion of tritium (3 H or T), a hydrogen isotope with two neutrons, and only 155 ppm of deuterium (2 H or D), a hydrogen isotope with one neutron, are present throughout the universe. There are three stable isotopes of oxygen as well, with 16 O constituting 99.76%, 17 O 0.04%, and 18 O 0.2% of water molecules.

Due to its greater density, deuterium oxide, also known as D 2O, is also referred to as heavy water. It serves as a neutron moderator in nuclear reactors. Tritium is radioactive and has a half-life of 4500 days. THO is only present in trace amounts in nature and is largely created in the atmosphere as a result of nuclear processes brought on by cosmic rays. Natural quantities of HDO, which consists of one protium and one deuterium atom, as well as D2O, which is far less common (0.000003%), may be found in regular water. Any such molecules are transient since the atoms recombine to form other molecules.

Other than the obvious difference in specific mass, freezing and boiling points, as well as other kinetic effects, are the most noticeable physical differences between H 2O and D 2O. This is due to the fact that the bonding energies of deuterium and protium differ noticeably because deuterium’s nucleus is twice as heavy as protium’s. The isotopologues can be distinguished by their various boiling points. Indicator of self-diffusion for H

2O is 23% more potent than D 2O at 25 °C. Low-purity heavy water contains hydrogen deuterium oxide (DOH) significantly more often than pure dideuterium monoxide (D 2O), since water molecules exchange hydrogen atoms with one another. Pure isolated D 2O may have an impact on metabolic processes; excessive doses have negative effects on the kidney and central nervous system. Humans are often oblivious to different tastes but occasionally report a scorching sensation or sweet flavour; little amounts can be ingested without any negative consequences. Any toxicity must be eaten in very high quantities in order to be felt. Rats, however, can detect heavy water by scent, despite the fact that many animals find it hazardous.

Importance of water

hold the temperature steady. Cushion and lubricate joints. Your spinal cord and other delicate structures need to be protected. Utilise urine, sweating, and bowel motions to eliminate wastes.

Uses of water in our everyday life

  • For drinking purpose.
  • For dish cleaning.
  • For cooking purpose.
  • For feeding plants.
  • For clothes washing.
  • To take bath.
  • For hydro-electricity generation.
  • For the car wash.

How much water should you drink a day?

The four to six cup guideline is for persons who are typically healthy. If you have certain medical conditions, such as thyroid disease or kidney, liver, or heart issues, or if you’re taking medications that cause you to retain water, like non-steroidal anti-inflammatory drugs (NSAIDs), opiate pain relievers, and some antidepressants, you could consume too much water.

Sources of Water

The sources of water include rivers, lakes, oceans, seas, rainwater, underground water, wells, dams, boreholes, and ponds.

Sources of Water


Leave a Reply

Your email address will not be published. Required fields are marked *