There are a number of elements that have been discovered and are listed in the periodic table.   These elements each have their own individual properties. There are an infinite number of compounds that can be made up from these elements each with its own characteristics.   The compounds are formed when elements combine or react with each other in a particular way in certain proportions.  A compound is formed when two or more atoms chemically bond together, the resulting compound is unique both chemically and physically from its parent atoms.

G.N.Lewis (in about 1916) observed that many elements are most stable when they contained eight electrons in their valence shell.   He suggested that atoms with fewer than eight valence electrons bond together to share electrons and complete their valence shells.

There are a number of bonding types including Ionic, Covalent, Metallic, Van der Waals ......etc.etc ref. Molecular Bonds.

Bonds form tending to stabilise a chemical system by releasing energy.   The larger the amount of energy released during the formation of a bond, the more stable the bond will be...... If two atoms release energy (Exothermic reaction ) by forming a bond, then the atoms will be more stable by staying together than they would be as individual atoms.

The bond between two different atoms when one atom (the cation) donates its valence electrons to another atom (the anion).   The resulting electrostatic charge bonds the two atoms together.   ref. Cations -ions

There are simple rules relating to ionic bonds

A compound illustrating ionic bonds is salt (Sodium Chloride).   For this compound a metal bonded to a non-metal: the metal is first in the chemical name: salt dissolves easily in water,: saline solutions are conductive: salt is a crystalline solid with a melting point of 800�C...

This bond most commonly occurs when two non-metals bond together.   This type of atomic bonding occurs when atoms share electrons.   As opposed to ionic bonding in which a complete transfer of electrons occurs, covalent bonding occurs when two (or more) elements share electrons.   Covalent bonding occurs because the atoms in the compound have a similar tendency for electrons (generally to gain electrons).    Because both of the non-metals will want to gain electrons, the elements involved will share electrons in an effort to fill their valence shells.

A good example of a covalent bond is that which occurs between two hydrogen atoms.   Atoms of hydrogen (H) have one valence electron in their first electron shell.  Since the capacity of this shell is two electrons, each hydrogen atom will 'want' to pick up a second electron.   In an effort to pick up a second electron, hydrogen atoms will react with nearby hydrogen (H) atoms to form the compound H₂.    Because the hydrogen compound is a combination of equally matched atoms, the atoms will share each others single electron, forming one covalent bond.   In this way, both atoms share the stability of a full valence shell.

There are two types of covalent bond the polar and the non-polar bond.  A good example of a non-polar bond is the hydrogen atom.  Because both atoms in the bond are the same with similar electrical properties there is no preference for the electron to be close to either of the atoms.   In reality whenever two atoms of the same type form a covalent bond the bond is non-polar.

A good example of a polar covalent bond is water .  This molecule has one large oxygen atom bonded to two small hydrogen atoms.   The oxygen atom shares two valency electrons and each hydrogen atom shares one.  The oxygen atom is a much larger atom than hydrogen atom and has a larger number of electrons.  The oxygen atom has a larger attraction for the elctrons compared to the hydrogen atom and the participating electrons are within the influence of the oxygen atom for a greater proportion of the time.  The hydrogen atoms therefore have a tendency to be positively charged and the oxygen atom tend to be negatively charged. This type of atom is called a dipole.

This bond relates to metals.   Metals in the solid phase consist of a crystal lattice of the metal cations, with the valence electrons being shared among all the cations..  Because the valence electrons are not associated with any particular atom, they are free to move under the influence of external electrical forces, as long as they remain within the bounds of the crystal.

All metals are made up of a vast collection of ions that are held together by metallic bonds.   A metal atom has a positive nucleus with negative electrons outside of it.   In a solid, each atom loses the outermost electron, which takes part in bonding.   They form a lattice of regularly spaced positive ions.   Each cation has no control over its bonding electron.

The positive metal cations are attracted to the negatively charged delocalised electrons.   The negative electrons are in turn attracted towards the positive metal cations.   It is these attractions that hold the structure together forming metallic bonds.

This type of bond is the result of an inter-molecular attractions.   These attractions are between molecules and neighboring molecules. The dipole molecules in liquid water are attracted to each other by electrostatic forces, and these forces have been named Van der Waals forces.

Even though the water molecule as a whole is electrically neutral, there is a dipole moment across the molecule (see above) this is a minute separation of the positive and negative charge centers.   This results in a net attraction between such polar molecules which causes some attraction between water molecules and contributes to viscosity and surface tension.   It is accepted that the Van der Waals forces holds water in the liquid state until thermal agitation becomes violent enough to break the bonds at 100�C.   With cooling, residual electrostatic forces between molecules cause most substances to liquify and eventually solidify.

Non-polar molecules also experience weak Van der Waals bonding because they can be polarised and experience fluctuating dipole moments which result in net attractions between molecules over time.