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Physical Structure
Main Concepts
The Water Molecule

Oxygen is electronegative

A water molecule consists of two atoms of hydrogen linked by covalent bonds to the same atom of oxygen. Atoms of oxygen are electronegative and attract the shared electrons in their covalent bonds. Consequently the electrons in the water molecule spend slightly more time around the oxygen atomic center and less time around the hydrogen atomic centers. The covalent bonds are therefore polar, and the oxygen atoms have a slight negative charge (from the presence extra electron share), while the hydrogens are slightly positive (from the extra un-neutralized protons).

Hydrogen Bonds

Opposite charges attract one another. The slight positive charges on the hydrogen atoms in a water molecule attract the slight negative charges on the oxygen atoms of other water molecules. This tiny force of attraction is called a hydrogen bond.

This bond is very weak. Hydrogen bonds are formed easily when two water molecules come close together, but are easily broken when the water molecules move apart again. They are only a small fraction of the strength of a covalent bond, but, there are a lot of them and they impart some very special properties to the substance we call water.

Water is a Liquid at Room Temperature

Over three-quarters of the planet earth is covered with water. Life probably started in such a liquid environment and water is the major component of living things (humans are over 60 percent water). At room temperature (anywhere from zero degree centigrade to 100 degrees centigrade), water is found in a liquid state. This is because of the tiny, weak hydrogen bonds which, in their billions, hold water molecules together for small fractions of a second.

Water molecules are constantly on the move. If they are moving fast enough they become a gas. A gas is a physical state of matter where the molecules are far apart and moving very quickly. But, because of the hydrogen bonds, as water molecules come together they stick to one another for a small, but significant amount of time. This slows them down, and holds them closer to one another. They become a liquid; a different state of matter where the molecules are closer and slower than in a gas.

Molecular water, therefore is a liquid at room temperature, a fact that is profoundly significant for all living things on this planet.

Water is a Universal Solvent

Everything dissolves in water. Stone, iron, pots, pans, plates, sugar, salt, and coffee beans all dissolve in water. Things which dissolve are called solutes and the liquid in which they dissolve is called a solvent.

Strongly polar substances (things with positive and/or negative charges) easily attract water molecules. The water molecules surround the charged solute; positive hydrogens close to negative charges and negative oxygens close to positive charges on the solute molecule. All this interaction suspends the solute molecule in a sea of water molecules; it disperses and dissolves easily.

Unequal Sharing

Electrons in the bonds between identical atoms (H-H) are shared uniformly, so the electrons spend equal amounts of time around each atomic center. These covalent bonds are non-polar. Electrons shared between unlike atoms are not shared equally, one atom gets more of the common electrons and is thus slightly negatively charged. The other atoms gets less than a full share of the electrons and is thus slightly positively charged.

Substances which dissolve easily and readily in water (sugar, salt, etc.) are called water-loving, or hydrophilic substances.

On the other hand, some solutes are non-polar and do not have any positive or negative charges. Water molecules are not attracted to these types of molecules (and, in fact, are sometimes repelled by them). Although tiny amounts of these substances (plastic, oil, etc.) will and do dissolve in water, most of their molecules simply form a boundary when they come in contact with water, and remain separate entities.

Substances which do not dissolve readily in water are called water-fearing, or hydrophobic substances.

© 2003, Professor John Blamire