Why is a molecule of #CH_4# nonpolar?

1 Answer
Mar 11, 2017

Answer:

The dipoles cancel out each other, depicted in a VSEPR diagram.

Explanation:

When determining the polarity of a molecule, we use 2 methods:

#1: Electrongativity,

#2: Symmetry of molecule.


  1. Electronegativity:

Electronegativity of a compound can be calculated by subtracting the element with the highest electronegativity with the lowest. Electronegativity can be found on the periodic table:

chemistry-reference.com

Compounds with an electronegativity of #0.0->0.6# are non-polar molecules. #0.7->1.7# are polar molecules, and #1.7+# are ionic compounds.

In methane's case,

#2.5-2.1=0.4#

Disclaimer: electronegativity only determines the intramolecular bond within a compound. Not the overall compound's polarity. To do this, we use symmetry.


  1. Molecular Symmetry:

Molecular symmetry is the use of geometrical arrangement that can help predict the shape of molecules. These arrangements are called VSEPR: Valence Shell Electron Pair Repulsion. Lewis diagrams work too, but lacks additional/specific information that is necessary for a more "academic" standard, that a VSEPR diagram provides.

When a molecule is "symmetrical", it means the dipoles cancel. This only occurs when the outside atoms are arranged in such a way that surrounds the central atom, where the dipoles cancel - no side of the molecule has more charge (positive or negative} than the other(s).

In methane's case,

wikibooks

All the outer atoms are the same - the same dipoles, and that the dipole moments are in the same direction - towards the carbon atom, the overall molecule becomes non-polar.

Therefore, methane has non-polar bonds, and is non-polar overall.

Keep in mind that molecules can have polar bonds, but be non-polar overall.

Hope this helps :)