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Electrons are like two negatively charged magnets, they repulse when pushed together. Therefore, atoms keep electrons in separate orbitals. Understanding these orbitals and how they bond is the last piece of general chemistry that is needed to be reviewed before getting to Lewis structures. Watch this video from Crash Course to get a refresher on orbitals.
In order to draw Lewis structures and line structures, you must know the bond angles between atoms to accurately represent the molecule's shape.
In addition to those shapes, atoms may also have lone pairs that affect the shape of the atom. Like Hank Green mentioned in the Crash Course video, H2O is a bent molecule. Since there's only 2 hydrogen attached to the Oxygen, it looks like the molecule should just be linear, with a 180 degree bond angle. However, since Oxygen has a lone pair of electrons, it forces the two hydrogen into a pseudo-trigonal planar arrangement, except one of the electron dense areas is just filed by a lone pair. This pushing around of atoms by lone pairs is known as the VSEPR Theory. (Valence-Shell Electron-Pair Repulsion) This can cause some pretty unique shapes as we get into higher and higher sublevels
To get a better understanding of VSEPR and molecular geometry, try this simulation by PhET.
To find your atom's orbital hybridization, simply add the number of lone pairs and number of atoms connected to it. 1, like hydrogen, would mean its simply s. 2 electron dense regions would be sp hybridized, with a bond angle of 180. Refer to the first chart on this page to see the full list.
If you need to further review this section, watch this video. If you think you are good, take the quiz to test your hybridization skills.
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