Polarity and Hydrogen BondingSee also: Chemical polarity
An important feature of water is its polar nature. The water molecule forms an angle, with hydrogen atoms at the tips and oxygen at the vertex. This angle formed is 104.3 degrees as opposed to the typical tetrahedral angle of 109 degrees. Because oxygen has a higher electronegativity than hydrogen, the side of the molecule with the oxygen atom has a partial negative charge. Also the presence of the lone pairs tend to push the oxygen away. An object with such a charge difference is called a dipole meaning two poles. The oxygen end is partially negative and the hydrogen end is partially positive, because of this the direction of the dipole moment points from the oxygen towards the center of the hydrogens. The charge differences cause water molecules to be attracted to each other (the relatively positive areas being attracted to the relatively negative areas) and to other polar molecules. This attraction contributes to hydrogen bonding, and explains many of the properties of water, such as solvent action.
A water molecule can form a maximum of four hydrogen bonds because it can accept two and donate two hydrogen atoms. Other molecules like hydrogen fluoride, ammonia, methanol form hydrogen bonds but they do not show anomalous behavior of thermodynamic, kinetic or structural properties like those observed in water. The answer to the apparent difference between water and other hydrogen bonding liquids lies in the fact that apart from water none of the hydrogen bonding molecules can form four hydrogen bonds, either due to an inability to donate/accept hydrogens or due to steric effects in bulky residues. In water, local tetrahedral order due to the four hydrogen bonds gives rise to an open structure and a 3-dimensional bonding network, resulting in the anomalous decrease of density when cooled below 4 °C.
Although hydrogen bonding is a relatively weak attraction compared to the covalent bonds within the water molecule itself, it is responsible for a number of water's physical properties. One such property is its relatively high melting and boiling point temperatures; more energy is required to break the hydrogen bonds between molecules. The similar compound hydrogen sulfide (H
2S), which has much weaker hydrogen bonding, is a gas at room temperature even though it has twice the molecular 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.
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