4: Properties of Water - Geosciences

4: Properties of Water - Geosciences

4: Properties of Water

The Water in You: Water and the Human Body

Water is indeed essential for all life on, in, and above the Earth. This is important to you because you are made up mostly of water. Find out what water does for the human body.

The Water in You: Water and the Human Body

​​​​​​​Water serves a number of essential functions to keep us all going

Think of what you need to survive, really just survive. Food? Water? Air? Facebook? Naturally, I'm going to concentrate on water here. Water is of major importance to all living things in some organisms, up to 90% of their body weight comes from water. Up to 60% of the human adult body is water.

According to H.H. Mitchell, Journal of Biological Chemistry 158, the brain and heart are composed of 73% water, and the lungs are about 83% water. The skin contains 64% water, muscles and kidneys are 79%, and even the bones are watery: 31%.

Each day humans must consume a certain amount of water to survive. Of course, this varies according to age and gender, and also by where someone lives. Generally, an adult male needs about 3 liters (3.2 quarts) per day while an adult female needs about 2.2 liters (2.3 quarts) per day. All of the water a person needs does not have to come from drinking liquids, as some of this water is contained in the food we eat.

Water serves a number of essential functions to keep us all going

  • A vital nutrient to the life of every cell, acts first as a building material.
  • It regulates our internal body temperature by sweating and respiration
  • The carbohydrates and proteins that our bodies use as food are metabolized and transported by water in the bloodstream
  • It assists in flushing waste mainly through urination
  • acts as a shock absorber for brain, spinal cord, and fetus
  • forms saliva
  • lubricates joints

According to Dr. Jeffrey Utz, Neuroscience, pediatrics, Allegheny University, different people have different percentages of their bodies made up of water. Babies have the most, being born at about 78%. By one year of age, that amount drops to about 65%. In adult men, about 60% of their bodies are water. However, fat tissue does not have as much water as lean tissue. In adult women, fat makes up more of the body than men, so they have about 55% of their bodies made of water. Thus:

  • Babies and kids have more water (as a percentage) than adults.
  • Women have less water than men (as a percentage).
  • People with more fatty tissue have less water than people with less fatty tissue (as a percentage).

There just wouldn't be any you, me, or Fido the dog without the existence of an ample liquid water supply on Earth. The unique qualities and properties of water are what make it so important and basic to life. The cells in our bodies are full of water. The excellent ability of water to dissolve so many substances allows our cells to use valuable nutrients, minerals, and chemicals in biological processes.

Water's "stickiness" (from surface tension) plays a part in our body's ability to transport these materials all through ourselves. The carbohydrates and proteins that our bodies use as food are metabolized and transported by water in the bloodstream. No less important is the ability of water to transport waste material out of our bodies.

10 properties of water

Water is undoubtedly one of the essential compounds on planet earth. It plays a critical role in the bodies of living things.

Pure water is colourless, odourless, and tasteless. Like any other chemical substances, it has unique properties that distinguish it from the rest. So, what are the properties of water?

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1. Water is polar

In chemical bonding, polarity refers to the distribution of electric charges over the atoms joined by a bond. Water is a molecular structure, which is composed of an oxygen atom and two hydrogen atoms.

Chemistry model of the water molecule - H2O consisting of two hydrogen atoms and an oxygen atom. Photo: azatvaleev
Source: Getty Images

The hydrogen atoms are covalently bonded to the oxygen one to form molecules. These molecules are then attached via hydrogen bonds.

Four pairs of electrons surround the oxygen atom. Two pairs are involved in bonding with the hydrogen atom, while the remaining two are unshared and on the opposite of the oxygen atom.

Because oxygen is more electronegative than hydrogen, it attracts the shared electrons and has partial negative charges. The less electronegative one (hydrogen) becomes partially positively charged.

As a result, the molecule is slightly charged at the ends. These partial charges cause the attraction of the molecules to form hydrogen bonds.

2. It is a universal solvent

Water is a universal solvent because most substances dissolve in it. This is one of the special properties of water, and it is made possible because of its polar characteristics.

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Considering that water is slightly charged at the ends (oxygen is negatively charged and hydrogen is positively charged), it can dissociate ionic compounds. Dissociation refers to the separation of negative and positive ions in an ionic compound.

When an ionic compound is put in water, it dissociates into positive and negative ions. The positive ions are attracted towards oxygen (which is negatively charged), and the negative ions are attracted towards hydrogen (which is positively charged). This is how dissolution occurs.

A woman looking at water in a conical flask. Photo:
Source: UGC

Being a universal solvent, it shouldn't be confused to mean that it dissolves every substance. Some organic compounds, such as oil, wax, and many more, do not dissolve in water.

3. Has high surface tension

Surface tension refers to the property of a liquid that allows it to resist an external force. The property is made possible because of cohesive forces. If you carefully place a needle on the surface of H2O, it floats because of surface tension.

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Usually, most people refer to the layer formed as skin. However, it is not really true that a "skin" forms on the surface.

The stronger cohesion between the water molecules, as opposed to the attraction of its molecules to the air makes it more difficult to move an object through its surface than to move it when it is completely submerged in it. Its surface tension is 72 dynes/cm at 25°C (room temperature).

4. Has high specific heat capacity

The specific heat capacity of any given substance refers to the amount of heat energy required to raise the temperature of one kilogram of that particular substance by one degree Celsius. H2O has a specific heat capacity of 4,200 Joules per kilogram per degree Celsius (J/kg°C). This means that you require 4200 Joules to raise a kilogram of water by one degree Celsius.

This is one of those properties of water that make it essential to life. As a result of its high specific heat capacity, H2O plays a crucial role in regulating temperatures on the Earth surface.

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H2O can absorb large amounts of heat energy before it begins to get hot. It also means that it releases heat energy slowly when situations cause it to cool.

Because water masses form a significant fraction of planet Earth, this property allows for the moderation of the Earth's climate. It also helps organisms living in it, such as amphibians and reptiles, to regulate their body temperature more effectively.

5. H2O is less dense as a solid than as a liquid

Density is the measure of the amount of matter (mass) contained in a given volume. It is calculated by dividing mass by volume (Mass/Volume). The SI unit of density is Kg/m^3.

With most liquids, solidification occurs when the temperature drops, thus lowering kinetic energy between molecules. As a result, the molecules move closer to one another as compared to when they were in liquid form.

Therefore, the density of their solid-state is greater density than their liquid state. But why does ice float on water? In H2O, things are a bit different.

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Water in ice state (solid-state) is less dense than its liquid state. When freezing occurs below -4 degrees Celsius, hydrogen bonds' orientation causes molecules to push farther apart, increasing their volume. This is referred to as the anomalous expansion.

Considering that the mass has remained considered, and the volume has increased, the density will eventually be lowered. This is the reason why ice floats in water.

6. Cohesive and adhesive properties

Cohesive and adhesive forces have a great influence on the properties of water. Cohesive forces refer to attraction forces existing between molecules of the same substance.

Adhesive forces, on the other hand, refer to attraction forces between unlike substances. H2O molecules have strong, cohesive forces due to their ability to form hydrogen bonds with one another.

It also has adhesive properties, which allow it to stick to other substances.

7. Boiling and freezing points

Boiling point refers to the temperatures at which the atmospheric pressure exerted by the surroundings upon a liquid is equivalent to the pressure exerted by the vapour of the liquid. Under this condition, extra heat causes the transformation of the liquid into its vapour form without raising the temperature.

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On the other hand, the freezing point is the point at which a liquid begins to turn into a solid. Considering that H2O can exist in the three states of matter (solid, liquid, and gas), it has specific boiling and freezing points.

Pure H2O boils at 100 degrees Celcius, and temperatures above that converts it into a gaseous state (vapour). It freezes at zero degrees Celcius, and temperatures below that cause it to solidify.

8. Amphoteric properties

Amphoteric substances exhibit both acidic and basic properties. The water molecule is amphoteric because it has both hydrogen and oxygen atoms.

A scientist with equipment holding tools during a scientific experiment. Photo: SARINYAPINNGAM
Source: Getty Images

The hydrogen atom could act as an acid in a reaction, and the oxygen atom (because of the two lone pairs) could react with a (H^+) to form a hydroxide ion (OH^-). The hydroxide ion (OH^-) could act as a base in a reaction.

This property is exhibited in chemical reactions of water with acids and bases. When it reacts with hydrochloric acid (HCl), it accepts the (H^+) from the acid, thus acts as a base (proton acceptor).

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On the other hand, when it reacts with ammonia gas, it donates the (H^+) to the gas to form Ammonium ion. In this particular reaction, it acts as an acid (proton donor).

9. Capillary action

Capillary action is the phenomenon of the ascension of liquids through slim tubes, cylinders or permeable substances. The action is due to adhesive and cohesive forces interacting between the liquid and the surface.

When intermolecular bonding of a liquid itself is substantially inferior to a substances’ surface it is interacting with, then capillary action occurs. H2O is one liquid that experiences this property.

Other liquid substances such as mercury do not experience this behaviour because the cohesive forces between their molecules are greater than the adhesive forces between its molecules and those of the tube they are in.

Therefore, if you insert a tube in a container (beaker) with mercury, you will notice that unlike in water, its level in the tube is slightly below its level in the container.

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Normally, when you insert a tube in a container of water, H2O level in the tube will rise above that of the normal level in the container.

10. Water can exist in solid, liquid, and gaseous state

Of all properties of water, the existence of H2O in different states of matter is an essential characteristic what you should know about. At standard temperature and pressure (STP), water is a liquid. It is one of the rare inorganic compounds that exist in the liquid state at these conditions.

Arrangement of particles in different states of matter. Photo: ttsz
Source: Getty Images

When its temperatures are lowered to its freezing point (zero degrees Celsius), it changes state to solid (ice). Its form can also be changed to gaseous when temperatures are increased beyond its boiling point (100 degrees Celsius).

There you have it. These are some of the most essential physical and chemical properties of water that you should know.

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Water is the chemical substance with chemical formula H
2 O one molecule of water has two hydrogen atoms covalently bonded to a single oxygen atom. [24] Water is a tasteless, odorless liquid at ambient temperature and pressure. Liquid water has weak absorption bands at wavelengths of around 750 nm which cause it to appear to have a blue colour. [3] This can easily be observed in a water-filled bath or wash-basin whose lining is white. Large ice crystals, as in glaciers, also appear blue.

Under standard conditions, water is primarily a liquid, unlike other analogous hydrides of the oxygen family, which are generally gaseous. This unique property of water is due to hydrogen bonding. The molecules of water are constantly moving concerning each other, and the hydrogen bonds are continually breaking and reforming at timescales faster than 200 femtoseconds (2 × 10 −13 seconds). [25] However, these bonds are strong enough to create many of the peculiar properties of water, some of which make it integral to life.

Water, ice, and vapour Edit

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). [26]

Heat capacity and heats of vaporization and fusion Edit

Water has a very high specific heat capacity of 4184 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. [27]

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) [28] and the heat capacity of steam at 100 °C is 2080 J/(kg·K). [29]

Density of water and ice Edit

The density of water is about 1 gram per cubic centimetre (62 lb/cu ft): this relationship was originally used to define the gram. [30] 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 [31] this is unusual. [d] Regular, hexagonal ice is also less dense than liquid water—upon freezing, the density of water decreases by about 9%. [34] [e]

These effects are due to the reduction of thermal motion with cooling, which allows water molecules to form more hydrogen bonds that prevent the molecules from coming close to each other. [31] While below 4 °C the breakage of hydrogen bonds due to heating allows water molecules to pack closer despite the increase in the thermal motion (which tends to expand a liquid), above 4 °C water expands as the temperature increases. [31] Water near the boiling point is about 4% less dense than water at 4 °C (39 °F). [34] [f]

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). [35] [36]

The unusual density curve and lower density of ice than of water is vital to life—if water were most dense at the freezing point, then in winter the very cold water at the surface of lakes and other water bodies would sink, lakes could freeze from the bottom up, and all life in them would be killed. [34] Furthermore, given that water is a good thermal insulator (due to its heat capacity), some frozen lakes might not completely thaw in summer. [34] The layer of ice that floats on top insulates the water below. [37] Water at about 4 °C (39 °F) also sinks to the bottom, thus keeping the temperature of the water at the bottom constant (see diagram). [34]

Density of saltwater and ice Edit

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 [38] (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 [38] 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 [38] 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.

Miscibility and condensation Edit

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 super-saturated and can occur if the air is rapidly cooled, for example, by rising suddenly in an updraft. [g]

Vapor pressure Edit

Compressibility Edit

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). [39]

The bulk modulus of water is about 2.2 GPa. [40] The low compressibility of non-gases, 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. [40]

The bulk modulus of water ice ranges from 11.3 GPa at 0 K up to 8.6 GPa at 273 K. [41] 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.

Triple point Edit

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 [42] [43] but, starting in 2019, the kelvin is now defined using the Boltzmann constant, rather than the triple point of water. [44]

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. [43] 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. [45] [46] [47]

The various triple points of water
Phases in stable equilibrium Pressure Temperature
liquid water, ice Ih, and water vapor 611.657 Pa [48] 273.16 K (0.01 °C)
liquid water, ice Ih, and ice III 209.9 MPa 251 K (−22 °C)
liquid water, ice III, and ice V 350.1 MPa −17.0 °C
liquid water, ice V, and ice VI 632.4 MPa 0.16 °C
ice Ih, Ice II, and ice III 213 MPa −35 °C
ice II, ice III, and ice V 344 MPa −24 °C
ice II, ice V, and ice VI 626 MPa −70 °C

Melting point Edit

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). [49] The melting point of ordinary hexagonal ice falls slightly under moderately high pressures, by 0.0073 °C (0.0131 °F)/atm [h] or about 0.5 °C (0.90 °F)/70 atm [i] [50] 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 [51] ).

Electrical properties Edit

Electrical conductivity Edit

Pure water containing no exogenous ions is an excellent electronic insulator, but not even "deionized" water is completely free of ions. Water undergoes auto-ionization in the liquid state when two water molecules form one hydroxide anion ( OH −
) and one hydronium cation ( H
3 O +
). Because of auto-ionization, 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 [52] .

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. [53] 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. [ citation needed ]

In pure water, sensitive equipment can detect a very slight electrical conductivity of 0.05501 ± 0.0001 μS/cm at 25.00 °C. [53] 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). [54] 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. [31]

Polarity and hydrogen bonding Edit

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°, [55] 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. [56]

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. [57]

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
2 S ), has much weaker hydrogen bonding due to sulfur's lower electronegativity. H
2 S 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.

Cohesion and adhesion Edit

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. [58]

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. [ citation needed ] 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. [59] They are important in biology, particularly when cells are dehydrated by exposure to dry atmospheres or to extracellular freezing. [60]

Surface tension Edit

Water has an unusually high surface tension of 71.99 mN/m at 25 °C [61] which is caused by the strength of the hydrogen bonding between water molecules. [62] This allows insects to walk on water. [62]

Capillary action Edit

Because water has strong cohesive and adhesive forces, it exhibits capillary action. [63] Strong cohesion from hydrogen bonding and adhesion allows trees to transport water more than 100 m upward. [62]

Water as a solvent Edit

Water is an excellent solvent due to its high dielectric constant. [64] 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. [65] 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 non-polar substances such as fats and oils are not. Non-polar 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.

Quantum tunneling Edit

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. [66] 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. [67] Later in the same year, the discovery of the quantum tunneling of water molecules was reported. [68]

Electromagnetic absorption Edit

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] [69]

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. This repeated, constantly reorganizing unit defines a three-dimensional network extending throughout the liquid. This view is based upon neutron scattering studies and computer simulations, and it makes sense in the light of the unambiguously tetrahedral arrangement of water molecules in ice structures.

However, there is an alternative theory for the structure of water. In 2004, a controversial paper from Stockholm University suggested that water molecules in the liquid state typically bind not to four but only two others thus forming chains and rings. The term "string theory of water" (which is not to be confused with the string theory of physics) was coined. These observations were based upon X-ray absorption spectroscopy that probed the local environment of individual oxygen atoms. [70]

Molecular structure Edit

The repulsive effects of the two lone pairs on the oxygen atom cause water to have a bent, not linear, molecular structure, [71] allowing it to be polar. The hydrogen-oxygen-hydrogen angle is 104.45°, which is less than the 109.47° for ideal sp 3 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. [72] 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).

Self-ionization Edit

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

The equilibrium constant for this reaction, known as the ionic product of water, K w = [ H 3 O + ] [ O H − ] >=[< m O^<+>>>][< m >>]> , 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. [73] 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:

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 K e q ≈ K w = [ H 3 O + ] [ O H − ] >approx K_< m >=[< m O^<+>>>][< m >>]>

Geochemistry Edit

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
2 O , naturally found in large quantities on the ocean floor.

Acidity in nature Edit

Rain is generally mildly acidic, with a pH between 5.2 and 5.8 if not having any acid stronger than carbon dioxide. [74] 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.

Several isotopes of both hydrogen and oxygen exist, giving rise to several known isotopologues of water. Vienna Standard Mean Ocean Water is the current international standard for water isotopes. Naturally occurring water is almost completely composed of the neutron-less hydrogen isotope protium. Only 155 ppm include deuterium ( 2
H or D), a hydrogen isotope with one neutron, and fewer than 20 parts per quintillion include tritium ( 3
H or T), which has two neutrons. Oxygen also has three stable isotopes, with 16
O present in 99.76%, 17
O in 0.04%, and 18
O in 0.2% of water molecules. [75]

Deuterium oxide, D
2 O , is also known as heavy water because of its higher density. It is used in nuclear reactors as a neutron moderator. Tritium is radioactive, decaying with a half-life of 4500 days THO exists in nature only in minute quantities, being produced primarily via cosmic ray-induced nuclear reactions in the atmosphere. Water with one protium and one deuterium atom HDO occur naturally in ordinary water in low concentrations (

0.03%) and D
2 O in far lower amounts (0.000003%) and any such molecules are temporary as the atoms recombine.

The most notable physical differences between H
2 O and D
2 O , other than the simple difference in specific mass, involve properties that are affected by hydrogen bonding, such as freezing and boiling, and other kinetic effects. This is because the nucleus of deuterium is twice as heavy as protium, and this causes noticeable differences in bonding energies. The difference in boiling points allows the isotopologues to be separated. The self-diffusion coefficient of H
2 O at 25 °C is 23% higher than the value of D
2 O . [76] Because water molecules exchange hydrogen atoms with one another, hydrogen deuterium oxide (DOH) is much more common in low-purity heavy water than pure dideuterium monoxide D
2 O .

Consumption of pure isolated D
2 O may affect biochemical processes – ingestion of large amounts impairs kidney and central nervous system function. Small quantities can be consumed without any ill-effects humans are generally unaware of taste differences, [77] but sometimes report a burning sensation [78] or sweet flavor. [79] Very large amounts of heavy water must be consumed for any toxicity to become apparent. Rats, however, are able to avoid heavy water by smell, and it is toxic to many animals. [80]

Light water refers to deuterium-depleted water (DDW), water in which the deuterium content has been reduced below the standard 155 ppm level.

Water is the most abundant substance on Earth and also the third most abundant molecule in the universe, after H
2 and CO . [21] 0.23 ppm of the earth's mass is water and 97.39% of the global water volume of 1.38 × 10 9 km 3 is found in the oceans. [81]

Acid-base reactions Edit

Water is amphoteric: it has the ability to act as either an acid or a base in chemical reactions. [82] According to the Brønsted-Lowry definition, an acid is a proton ( H +
) donor and a base is a proton acceptor. [83] When reacting with a stronger acid, water acts as a base when reacting with a stronger base, it acts as an acid. [83] For instance, water receives an H +
ion from HCl when hydrochloric acid is formed:

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

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:

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:

Ligand chemistry Edit

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
2 O) 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
2 O 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. [84]

Organic chemistry Edit

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 SN2 substitution and E2 elimination reactions the latter is then known as a dehydration reaction.

Water in redox reactions Edit

Water contains hydrogen in the oxidation state +1 and oxygen in the oxidation state −2. [85] It oxidizes chemicals such as hydrides, alkali metals, and some alkaline earth metals. [86] [87] One example of an alkali metal reacting with water is: [88]

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. [89] Note that the rusting of iron is a reaction between iron and oxygen [90] 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
2 O . Almost all such reactions require a catalyst. [91] An example of the oxidation of water is:

Electrolysis Edit

Water can be split into its constituent elements, hydrogen, and oxygen, by passing an electric current through it. [92] This process is called electrolysis. The cathode half reaction is:

The anode half reaction is:

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. [92] The operating potential is actually 1.48 V or higher in practical electrolysis.

Henry Cavendish showed that water was composed of oxygen and hydrogen in 1781. [93] The first decomposition of water into hydrogen and oxygen, by electrolysis, was done in 1800 by English chemist William Nicholson and Anthony Carlisle. [93] [94] In 1805, Joseph Louis Gay-Lussac and Alexander von Humboldt showed that water is composed of two parts hydrogen and one part oxygen. [95]

Gilbert Newton Lewis isolated the first sample of pure heavy water in 1933. [96]

The properties of water have historically been used to define various temperature scales. Notably, the Kelvin, Celsius, Rankine, and Fahrenheit scales were, or currently are, defined by the freezing and boiling points of water. The less common scales of Delisle, Newton, Réaumur and Rømer were defined similarly. The triple point of water is a more commonly used standard point today.

The accepted IUPAC name of water is oxidane or simply water, [97] or its equivalent in different languages, although there are other systematic names which can be used to describe the molecule. Oxidane is only intended to be used as the name of the mononuclear parent hydride used for naming derivatives of water by substituent nomenclature. [98] These derivatives commonly have other recommended names. For example, the name hydroxyl is recommended over oxidanyl for the –OH group. The name oxane is explicitly mentioned by the IUPAC as being unsuitable for this purpose, since it is already the name of a cyclic ether also known as tetrahydropyran. [99] [100]

The simplest systematic name of water is hydrogen oxide. This is analogous to related compounds such as hydrogen peroxide, hydrogen sulfide, and deuterium oxide (heavy water). Using chemical nomenclature for type I ionic binary compounds, water would take the name hydrogen monoxide, [101] but this is not among the names published by the International Union of Pure and Applied Chemistry (IUPAC). [97] Another name is dihydrogen monoxide, which is a rarely used name of water, and mostly used in the dihydrogen monoxide parody.

Other systematic names for water include hydroxic acid, hydroxylic acid, and hydrogen hydroxide, using acid and base names. [j] None of these exotic names are used widely. The polarized form of the water molecule, H +
OH −
, is also called hydron hydroxide by IUPAC nomenclature. [102]

Water substance is a term used for hydrogen oxide (H2O) when one does not wish to specify whether one is speaking of liquid water, steam, some form of ice, or a component in a mixture or mineral.

10 Properties Of Water

Polar covalent bonds are a type of covalent bond and means unequal sharing of electrons.

High Heat Capacity - huge amount of heat stored.Water's heat capacity stores a lot of heat which is why it is difficult to boil water.

Density - the amount of thickness in a substance. Water's density is less and more stable as a solid.

Capillary Action - is the movement of liquid to slide through narrow areas by the attraction of molecules of the liquid to the molecules of the solid.

Solute - substance dissolved in a solevent to form a solution.

Solvent - fluid that dissolves solute.

Surface Tension - water is pulled together creating the smallest surface area possible.

Cohesion - water attracted to other water molecules becayse of polar properties.

Adhesion - water attracted to other molecules.

Gas, Liquid, and Solid - are known as the three states of matter or material, but each of solid and liquid states may exist in one or more forms

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Properties of Water

Water’s chemical description is H2O. As the diagram to the left shows, that is one atom of oxygen bound to two atoms of hydrogen. The hydrogen atoms are “attached” to one side of the oxygen atom, resulting in a water molecule having a positive charge on the side where the hydrogen atoms are and a negative charge on the other side, where the oxygen atom is. This uneven distribution of charge is called polarity. Since opposite electrical charges attract, water molecules tend to attract each other, making water kind of “sticky.” As the right-side diagram shows, the side with the hydrogen atoms (positive charge) attracts the oxygen side (negative charge) of a different water molecule. (If the water molecule here looks familiar, remember that everyone’s favorite mouse is mostly water, too). This property of water is known as cohesion.

Water, the liquid commonly used for cleaning, has a property called surface tension. In the body of the water, each molecule is surrounded and attracted by other water molecules. However, at the surface, those molecules are surrounded by other water molecules only on the water side. A tension is created as the water molecules at the surface are pulled into the body of the water. This tension causes water to bead up on surfaces (glass, fabric), which slows wetting of the surface and inhibits the cleaning process. You can see surface tension at work by placing a drop of water onto a counter top. The drop will hold its shape and will not spread.

In the cleaning process, surface tension must be reduced so water can spread and wet surfaces. Chemicals that are able to do this effectively are called surface active agents, or surfactants. They are said to make water “wetter.” Surfactants perform other important functions in cleaning, such as loosening, emulsifying (dispersing in water) and holding soil in suspension until it can be rinsed away. Surfactants can also provide alkalinity, which is useful in removing acidic soils.

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