Electron pairs are very important when it comes to understanding the shapes of molecules. This is explained by something called VSEPR, which stands for Valence Shell Electron Pair Repulsion.

Here's how it works: Electron pairs around a central atom push away from each other. This pushing helps decide how the molecule will be shaped. The goal is to arrange the electron pairs in a way that reduces this push.

Common Shapes of Molecules:

• Linear: This shape has a bond angle of 180°. An example is BeCl₂.

• Trigonal Planar: In this shape, the bond angle is 120°. A good example is BF₃.

• Tetrahedral: This shape has a bond angle of 109.5°. You can see this shape in CH₄.

• Trigonal Bipyramidal: This shape includes bond angles of 90° and 120°. An example is PCl₅.

• Octahedral: This shape has a bond angle of 90°. A molecule that fits this shape is SF₆.

It’s important to note that lone electron pairs take up more space compared to bonded pairs. This influences angles and the shape of the molecule.

For instance, NH₃ (ammonia) has a trigonal pyramidal shape because it has one lone pair.

Understanding these shapes helps scientists predict how molecules will behave!