Chemical bonds are atoms that are held together within molecules or crystals. They form as the result of atoms sharing electrons. They also form from the complete transfer of electrons. The three main types of bonds are ionic, covalent, and polar covalent. Hydrogen bonds are often included under the heading of chemical bonds as well.
Ionic bonds are chemical bonds that are developed from reactions of low electronegativity elements with those that are high electronegativity. In such cases, a complete transfer of electrons takes place. An example of an ionic bond is regular table salt, referred to scientifically as sodium chloride. In forming this chemical bond, sodium transfers its outer shell electron to chlorine. Sodium has just one outer shell electron and chlorine requires just one for its shell.
Covalent chemical bonds are a bit different and usually form between atoms that lack completely filled outer shells. In covalent bonds, electrons are shared fully. These bonds are formed between atoms that are similar in electronegativity. Most organic compounds contain covalent bonds.
Polar covalent chemical bonds are somewhat in between ionic and covalent bonds; the atoms in these chemical bonds share electrons. However, the atoms focus most of their time on one particular atom, rather than others in the same compound. This type of chemical bond forms when atoms are very different in electronegativity. Water is an example of a polar covalent bond; oxygen beats hydrogen in terms of electronegativity. As such, the electrons in water spend more time around oxygen than hydrogen.
In water, the oxygen end of the water molecule features a negative charge while the hydrogen end has a positive charge. This opposite charge results in a powerful electrostatic attraction that is considered a hydrogen bond. This type of chemical bond has an important role in determining critical characteristics of water, vital for living things. However, hydrogen bonds are not limited to water alone. They can occur within other molecules as well.
