CHEMICAL BOUNDING

Chemical Bond Types

Ionic Bonds
An ionic bond is formed by the attraction of oppositely charged atoms or groups of atoms. When an atom (or group of atoms) gains or loses one or more electrons, it forms an ion. Ions have either a net positive or net negative charge. Positively charged ions are attracted to the negatively charged 'cathode' in an electric field and are called cations. Anions are negatively charged ions named as a result of their attraction to the positive 'anode' in an electric field.
Every ionic chemical bond is made up of at least one cation and one anion.
Ionic bonding is typically described to students as being the outcome of the transfer of electron(s) between two dissimilar atoms. The Lewis structure below illustrates this concept.

For binary atomic systems, ionic bonding typically occurs between one metallic atom and one nonmetallic atom. The electronegativity difference between the highly electronegative nonmetal atom and the metal atom indicates the potential for electron transfer.
Sodium chloride (NaCl) is the classic example of ionic bonding. Ionic bonding is not isolated to simple binary systems, however. An ionic bond can occur at the center of a large covalently bonded organic molecule such as an enzyme. In this case, a metal atom, like iron, is both covalently bonded to large carbon groups and ionically bonded to other simpler inorganic compounds (like oxygen). Organic functional groups, like the carboxylic acid group depicted below, contain covalent bonding in the carboxyl portion of the group (HCOO) which itself serves as the anion to the acidic hydrogen ion (cation).

Covalent
A covalent chemical bond results from the sharing of electrons between two atoms with similar electronegativities A single covalent bond represent the sharing of two valence electrons (usually from two different atoms). The Lewis structure below represents the covalent bond between two hydrogen atoms in a H₂ molecule.

Dot Structure	Line Structure

Multiple covalent bonds are common for certain atoms depending upon their valence configuration. For example, a double covalent bond, which occurs in ethylene (C₂H₄), results from the sharing of two sets of valence electrons. Atomic nitrogen (N₂) is an example of a triple covalent bond.
Double Covalent Bond

Triple Covalent Bond

The polarity of a covalent bond is defined by any difference in electronegativity the two atoms participating. Bond polarity describes the distribution of electron density around two bonded atoms. For two bonded atoms with similar electronegativities, the electron density of the bond is equally distributed between the two atom is This is a nonpolar covalent bond. The electron density of a covalent bond is shifted towards the atom with the largest electronegativity. This results in a net negative charge within the bond favoring the more electronegative atom and a net positive charge for the least electronegative atom. This is a polar covalent bond.

Coordinate Covalent

A coordinate covalent bond (also called a dative bond) is formed when one atom donates both of the electrons to form a single covalent bond. These electrons originate from the donor atom as an unshared pair.

Both the ammonium ion and hydronium ion contain one coordinate covalent bond each. A lone pair on the oxygen atom in water contributes two electrons to form a coordinate covalent bond with a hydrogen ion to form the hydronium ion. Similarly, a lone pair on nitrogen contributes 2 electrons to form the ammonium ion. All of the bonds in these ions are indistinguishable once formed, however.

Ammonium (NH4+)	Hydronium (H3O+)

Network Covalent

Some elements form very large molecules by forming covalent bonds. When these molecules repeat the same structure over and over in the entire piece of material, the bonding of the substance is called network covalent. Diamond is an example of carbon bonded to itself. Each carbon forms 4 covalent bonds to 4 other carbon atoms forming one large molecule the size of each crystal of diamond.

Silicates, [SiO₂]_x also form these network covalent bonds. Silicates are found in sand, quartz, and many minerals.

Metallic

The valence electrons of pure metals are not strongly associated with particular atoms. This is a function of their low ionization energy. Electrons in metals are said to be delocalized (not found in one specific region, such as between two particular atoms).

Since they are not confined to a specific area, electrons act like a flowing “sea”, moving about the positively charged cores of the metal atoms.

• Delocalization can be used to explain conductivity, malleability, and ductility.
• Because no one atom in a metal sample has a strong hold on its electrons and shares them with its neighbors, we say that they are bonded.
• In general, the greater the number of electrons per atom that participate in metallic bonding, the stronger the metallic bond.

Hydrogen bond
A hydrogen bond is a weak type of force that forms a special type of dipole-dipole attraction which occurs when a hydrogen atom bonded to a strongly electronegative atom exists in the vicinity of another electronegative atom with a lone pair of electrons. These bonds are generally stronger than ordinary dipole-dipole and dispersion forces, but weaker than true covalent and ionic bonds