Water (H2O) is an essential inorganic chemical compound that covers approximately 71% of the Earth's surface and is critical for all known forms of life. It exists in solid, liquid, and gas states, exhibiting unique properties such as high specific heat, polarity, and hydrogen bonding, which contribute to its behavior as a solvent and its role in environmental processes like the water cycle. The document details the characteristics, phases, and biological importance of water, as well as its chemical behaviors, such as self-ionization and its role in geochemistry.

1. Water
May 1, 2023admin
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.

2. 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.[25] At room temperature and pressure, water has no taste or odour.
Liquid water appears to be blue because of modest absorption bands at
wavelengths of around 750 nm.[3] 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.

3. 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

4. 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.
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.

5. 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
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

6. 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

7. 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.
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,

8. 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

9. energy of the oxygen atom’s nonboning hybrid orbitals is lower than that of the
hydrogen atom’s hybrid orbitals.
Chemical Properties
Self-ionization
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^{-}}}]}
Geochemistry
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

10. 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
If there is no acid stronger than carbon dioxide, rain typically has a pH
between 5.2 and 5.8.[76] Acid rain will result from the high levels of nitrogen
and sulphur oxides in the atmosphere dissolving into the clouds and droplets.
Isotopologues
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.

11. 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.

12. Sources of Water
The sources of water include rivers, lakes, oceans, seas, rainwater,
underground water, wells, dams, boreholes, and ponds.
Reference
 https://en.wikipedia.org/wiki/Water
 https://www.worldbank.org/en/topic/water
 https://www.worldwildlife.org/threats/water-scarcity