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Tetrahedral
If E is 0 -- if there are no lone pairs around the central atom, in other words -- the atom will be at the center of a tetrahedron, with each of the four bonds pointing toward one of the four corners. The angles between any two bonds are equal and close to 109.5 degrees. This is the most common arrangement for carbon atoms in organic compounds. Of the many examples, methane is probably the most popular.
Seesaw
If E is 1 -- if there is one lone pair of electrons, in other words -- the atom has a seesaw atomic configuration. Two of the bonds point straight up and down; the other two form a triangle with the central atom in the horizontal plane. The angle between the vertical bonds and the horizontal ones is ~90 degrees; the angle between the two horizontal-plane bonds is a little below 120 because of the lone pair. Sulfur tetrafluoride is one example.
Square Planar
Imagine that E is 2 and thus the formula is AB4E2 with two lone pairs of electrons and four bonds. In this case, the compound has a square planar atomic configuration, with all four bonds lying in the same plane and an angle of 90 degrees between any two bonds. Xenon tetrafluoride is one example of this somewhat more unusual geometry.
Considerations
You won't encounter any compounds with three lone pairs and four bonds, so zero, one and two lone pairs are the only possibilities you need to consider for an AB4 compound. For the one-electron pair possibility, remember that in VSEPR, the lone pair behaves as if it repels the bonding pairs more strongly than they repel each other. Consequently, the angle between the two bonds is going to be less than 120 degrees, although you can't use VSEPR to calculate the bond angle. You merely get a rough idea of what it will be.