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Simple Explanation:
Concept of Electron Densities
Electron densities, including bonds and lone pairs, repel each other due to electron repulsion. To minimize this repulsion, they position themselves as far apart from each other as possible around the central atom.
VSEPR theory uses this concept to predict the molecular shapes of compounds.
Molecular Geometry Chart
This chart displays the molecular geometries and bond angles of various molecules based on electron densities.
A detail you might've noticed is that lone pairs actually have slightly more repulsion than bond pairs. Therefore, it can affect non-planar molecules by slightly decreasing the bond angles adjacent to the lone pairs.
While less emphasized in introductory-level general chemistry, double and triple bonds exert more repulsion than single bonds. This is due to the increased electron density from overlapping orbitals, which can slightly decrease adjacent bond angles.
You would need to memorize these geometries and angles. Although it seems like a lot at first, further grasping on concepts and patterns would do the trick. Here are some tips:
The key bond angles to remember are 90°, 109.5°, 120°, and 180°. If a bond angle is affected by a lone pair on the same plane, simply prefix the angle with a < symbol.
The names for molecular geometry can indicate what a molecule looks like.
Molecular Geometry vs. Electron Geometry
Molecular geometry refers to the shape formed by the atoms in a molecule, while electron geometry describes the shape formed by all electron densities, including both bonds and lone pairs.
A molecule’s molecular geometry is determined by the specific number of bonds and lone pairs it has, whereas its electron geometry is the same for any molecule with the same total number of electron densities.
For example, a molecule with 4 bonds and 2 lone pairs would have an octahedral electron geometry, but it would appear as a square planar.
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