Valence bond theory (VB) and molecular orbital theory (MO) help explain why and how electrons are shared between atoms although their approaches are quite different.  Valence bond theory uses the individual atoms coming together with their atomic orbitals overlapping to pair electrons into covalent bonds.  Molecular orbital theory uses the idea of a positive nucleus surrounded by molecular orbitals.  Again, the two theories can be broken down into the “visual” and “predicting” realms.  I will explore each in different posts.
There are three main “variations” using overlapping orbitals (valence bond theory):  Lewis dot, Hybrid orbital theory, and Resonance theory.  It is important to understand that these variations are not unique from each other.  It is more like looking at a puzzle from different angles.  They all use the basic premise that orbitals must overlap to allow sharing of two electrons to create a covalent bond, they merely approach that concept very differently.
In an earlier post, I discussed the use of Lewis dot.  Figure 12.11 (see above) in the physical chemistry textbook, shows how bonds are formed by overlapping of atomic orbitals.
Hybrid orbital theory is a more complex bonding tool that uses the orbital overlap to explain both shape and properties like bond length, bond strength, etc.  This approach requires knowledge of which atomic orbitals will be participating in the bonding process.
In order to bond, the atomic orbitals must “hybridize” into a new orbital.  For example, the carbon atom’s four valence electrons are in the 2s and 2p orbitals, providing it with three possible hybrids.  Because the s orbital is lower in energy than the p orbitals, it must always be included in any type of bond.  Figure 12.13 (refer to physical chemistry textbook) will show you a visual representation of the new “hybrid” orbitals.

There are two very important results from the creation of these hybrid orbitals.  The first is the creation of new orbitals that are identical to each other.  The second is that in the first two cases (sp and sp²), there are left over atomic p orbitals.
This leads to a more in depth look at what a bond really is.  In a single bond there are two electrons shared between two atoms in a very strong covalent bond.  Both valence bond theory and molecular orbital theory have identified this arrangement and it has been named a sigma bond (σ) bond.
Multiple bonds, while overall stronger as a whole, contain a weaker “second” bond, called a pi (π) bond, that can be broken and then reformed repeatedly with rotation.  A double bond contains one pi bond while a triple bond contains two.
One of the main purposes of the hybrid orbital theory was to come up with a system that would explain the actual observations/characteristics of different molecules.  In organic chemistry, we like to use the example of methane, CH₄.  Each bond inside methane is exactly identical, the same length, energy, etc.  How is that possible if it is the result of overlapping atomic orbitals of different energies?
 The hybridization of those orbitals results in four identical hybrid orbitals which would result in four identical bonds.  Only two “things” can “live” in a hybrid orbital, sigma bonds and lone pairs.
The pi bonds are formed not by a complete overlap of the hybrid orbitals, but by a less effective overlap (I call this an interact) of the atomic p orbitals.

The line/bond is the true overlap of the hybrid orbitals to create the sigma bond while the dumbbell shapes are two adjacent atomic p orbitals.  If electrons are available, they will interact and form a pi bond if they are oriented in the same plane.  Notice that because of their arrangement, the sigma bond will allow free rotation but the pi bond would break only to be reformed at increments of 180 degrees.  A triple bond would have two more p orbitals in “z” axis and the pi bonds would reform at 90-degree increments!
See interactive multimedia of ethene (HC=CH) and ethyne/acetylene (HC≡CH) in section 12.4 of the physical chemistry textbook.
Hybrid orbital theory not only gives a very good visual representation of molecules, but it also provides insight into the formation and properties of the resulting molecules.