Water: Structure, Basic physical properties, Molecular Structure, Polarity, Hydrophobic and Hydrophilic Interactions
Water is a one-of-a-kind, all-encompassing substance that is essential to all living things. Some of the water’s distinctive qualities are necessary for life, while others have far-reaching implications for the size and shape of living organisms, how they function, and the physical restrictions or constraints they must operate within.
Basic physical properties
1. Water is a tiny solvent, with each molecule occupying roughly 0.03 nm3 in the liquid state at normal temperature and pressure, but it is extremely cohesive due to strong intermolecular interactions (hydrogen bonds, or H-bonds) between oxygen and hydrogen atoms.
2. Its high boiling point, the enormous amount of heat required to evaporate it, and its high surface tension all reflect this. Replacement of one or both hydrogens significantly lowers these intermolecular interactions, lowering their amplitude.
3. Because of the strong cohesive interactions in water, it has high viscosity (the ability to store a large amount of potential energy for a given increment in kinetic energy)
4. Water’s physical qualities have biological significance. Water occurs mostly in liquid form in the range of settings where life thrives, owing to its high boiling point, while the other two phases, ice and vapour, play an important part in moulding the environment.
5. The high specific heat and heat of vaporisation of water have significant cellular and physiological ramifications for organisms, particularly for the efficiency of activities involving heat transmission, temperature regulation, cooling, and so on.
6. The major property of water that governs how quickly molecules and ions can be transported and diffused in an aqueous solution is viscosity. As a result, it sets a physical upper limit on the rates of numerous molecular-level events, in which organisms must survive and evolve within.
7. These include ion channel conductance rates, substrate interaction with enzymes, binding rates, and macromolecular assembly rates. It also establishes a limit on the length scale beyond which biological processes can take place only by diffusion.
8. Evolution has pushed the components of living systems to the limits established by water’s viscosity in many circumstances, such as enzyme-substrate reactions.
9. Water’s high surface tension is significant on two levels. First, below a length scale of roughly 1 mm, surface tension forces outweigh gravitational and viscous forces, effectively impenetrating the air-water interface. This becomes a crucial element in minuscule insects’, bacteria’s, and other microorganisms’ environments and lifestyles. Second, the surface tension of water plays an important impact in its solvent characteristics at the molecular scale (0.1–100 nm).
10. Water’s high dielectric constant contributes to its solvent properties. Though important, the ecological importance of water expansion when it cools below 48°C and freezes is mostly indirect due to geophysical factors such as ocean and lake freezing, the creation of the polar ice cap, and freeze weathering.
Molecular Structure and Polarity
It is made up of two O–H bonds with a length of 0.096 nm and an angle of 104.58 degrees. Water’s size, shape, and polarity are also important characteristics. If two unbonded atoms are brought close enough to have their electron orbitals overlap, they will resist each other strongly. Two atoms attract each other weakly at longer distances due to an induced dipole-induced dipole (London dispersion) force. The van der Waals interaction is a combination of repulsive and attractive interactions. The diameter of an atom is typically defined as the point where repulsive and attractive forces balance, which for oxygen and hydrogen is 0.32 nm and 0.16 nm, respectively.
As a result, the water molecule resembles a sphere. Water is electrically neutral, but because oxygen’s electronegativity is significantly higher than hydrogen’s, electrons are concentrated more around the former, making it electrically polarized. Assigning a partial charge to each atom to duplicate the molecule’s net charge, dipole moment, and possibly higher-order electrical moments is a handy technique to depict the polarity of a molecule.
The polarity of an atom is measured by the magnitude of its partial charge. In water, each hydrogen has a charge of around 10.5 while the oxygen has a charge of the opposite sign and double the magnitude. The hydrogens of an apolar molecule like methane, on the other hand, have a partial charge of 0.1, and the dipole moment of methane is zero. As a result, water is a highly polar molecule capable of strong electrostatic interactions with itself, other molecules, and ions.
When all of this is considered, a water molecule can be visualised as a slightly sticky sphere with a radius of 0.32 nm and two positive charges of 1 1/2 and a negative charge of 2 1 at the hydrogen and oxygen atomic centres, respectively. This basic molecular model can explain many of the features of liquid water, including its cohesiveness, high heat of vaporisation, dielectric constant, and surface tension. Other properties, such as the density’s temperature dependence, necessitate a more complex model that takes into account water’s flexibility, polarizability, and quantum mechanical effects.
The dielectric constant is a measurement of how quickly an electric field can polarise a material in comparison to a vacuum. The magnitude of dielectric polarisation (dipole moment per unit volume) caused by a unit field is used to determine it. Water has a dielectric constant almost 80 times that of vacuum and is polarizable to an order of magnitude more than most chemical solvents. A polar liquid’s dielectric constant is determined by four primary factors: The molecule’s permanent dipole moment, the density of dipoles, how easy they may reorient in response to a field, and how cooperative this reorientation is.
Water has a high dipole moment, is tiny enough to have a lot of dipoles per unit volume, and can be reoriented quickly (within 10 ps) in the liquid state. Furthermore, because water molecules are heavily H linked, the polarisation reaction is cooperative: water molecules cannot simply reorient themselves without the help of their neighbours. They realign themselves in three-person groupings. Finally, the polarizability and flexibility of water contribute a little amount to the dielectric constant. Water’s extremely high dielectric constant can be explained by all of these characteristics.
The dielectric constant is increased by lowering the temperature because it minimises the randomising thermal fluctuations that resist dipole alignment by an electrostatic field. Water’s static dielectric constant increases through the freezing point, which is unusual. Although the time scale of reorientation is six orders of magnitude longer, the relevance of the cooperative impact of dipole reorientation is demonstrated by the high dielectric constant of ice.
Water can be found in three states: vapour, liquid, and ice, the latter of which has at least nine different forms. The liquid phase is the most significant for biological events. The hydrogen-bonding interaction dominates both structures. In biological compounds, the hydrogen bond (H-bond) is a strong link formed between polar hydrogen and another heavy atom, commonly carbon, nitrogen, oxygen, or sulphur.
An H bond between two waters has a strength of 22.7 kJ mole 2 1 in the gas phase, but its strength in liquids and solids is highly dependent on geometry and the surrounding molecules. Although its energy as a function of length and angle can be quite accurately described by a Coulombic interaction between the partial atomic charges on the hydrogen, the heavy atom it is covalently attached to, and the oxygen, nitrogen, carbon, or sulphur atom with which it is making the H-bond, it is sometimes described as intermediate between ionic and covalent bonds in character. From oxygen to oxygen, the H-bonds are 0.275 nm long and linear. Ice has a symmetrical H-bonding pattern, with each water forming two donor H bonds with its hydrogen atoms and two acceptor H bonds with hydrogen atoms from neighbouring waters. Water has a 2–2 H-bonding symmetry, which is an important property. When combined with an H–O–H bond angle that is very close to the ideal tetrahedral angle of 109.58, and the tendency for the four neighbouring waters to resist each other electrostatically, the result is a unique tetrahedral structure.
Experiments and computer simulations demonstrate that the open structure of liquid water is due to a high degree of angular ordering. The probability distribution of H-bond angles produced by each water to its 4.7 neighbours is depicted. It is bimodal, as opposed to Ice I’s H-bond angle distribution, which has a single peak at 08. In liquid, four of the H bonds are nearly linear (mean angle of around 128), and their length is quite close to that of Ice I. This suggests that most of Ice I’s tetrahedral structure survives in liquid water, though in a deformed form. Because these neighbours must sit in a face of the tetrahedron created by the primary hydration fluids, the H-bond to the additional neighbour (s) is more twisted, with an average angle of around 528. The abnormal temperature dependency of water density is also due to the open tetrahedral structure. When water melts, it contracts and continues to compress until it reaches a temperature of 48 degrees Celsius, at which point it expands like most liquids. The collapse of the open tetrahedral structure due to increasingly twisted H-bonds overrides the natural tendency for materials to expand as molecules grow further apart during the contraction phase.
Hydrophobic and Hydrophilic Interactions:
The most essential biological function of water is as a solvent. It can dissolve a wide range of essential compounds, from simple salts to tiny molecules like sugars and metabolites to huge molecules like proteins and nucleic acids. Chemical reactions, molecule association and binding, diffusion-driven interactions, and ion conduction are all molecular activities that occur at substantial rates only in solution, emphasising the relevance of water’s solvent qualities.
Water’s differential effect as a solvent – the fact that it dissolves some molecules considerably better than others – is just as essential as its ability to dissolve certain compounds. The solubilities are 50 orders of magnitude higher! Ions and charged amino acids like arginine and aspartic acid are found at the high end of the spectrum. These are hydrophilic solutes (water-loving). Asparagine, the peptide backbone of proteins, the phosphate sugar backbone of nucleic acids, sugars, and lipid head groups are all included in this category of neutral amino acids. Aliphatic amino acids like leucine, aromatic amino acids like phenylalanine, and lipid hydrocarbon ‘tails’ are all on the low solubility end of the spectrum. Hydrophobic solutes are those that repel water. Other solutes, such as nucleic acid bases and the amino acid tryptophan, have a range of solubility and can’t be categorised as either hydrophobic or hydrophilic.