Polar and Non-Polar Molecules

By Curtis Bustos

Let’s start by examining what a polar and non-polar molecule is. To do this, we need to recognize the difference between polar and non-polar covalent bonds. Remember, covalent bonds consist of shared electrons between atoms that are bonded together, and sometimes this sharing is not necessarily distributed equally. So, equally shared electrons are called non-polar covalent bonds and unequally shared electrons are called polar covalent bonds.

Note – ionic bonds occur when there is a large difference in electronegativity between two atoms such as with sodium chloride (NaCl).  

However, this is only part of the story. These polar covalent bonds must also contribute toward a dipole-moment or more specifically a net-dipole to result in a polar molecule. Net-dipoles will be explained in further detail, but for now, we will consider the following statements:

A polar molecule has polar covalent bonds (electrons are shared unequally) that result in a dipole-moment and does not cancel out.

A non-polar molecule has non-polar covalent bonds (electrons are shared equally) that DO NOT result in a dipole-moment.  

What causes a molecule to be polar vs. non-polar? Below are four very important factors that contribute toward a molecule’s polarity (DEEM):

-       Dipole moment

-       Electronegativity

-       Electron density

-       Molecular geometry

Dipole moment(s):

Key terms:

Dipole moment – the result of unequal electron-sharing in a covalent bond.

Net dipole moment – the overall dipole in a molecule after adding and subtracting ALL dipoles as there may be more than one.

Note – “pull” will be used throughout this learning tool to describe attractive forces between an atom’s nucleus and nearby electrons.

The arrows on the molecules illustrated above represent dipoles. These arrows point to the more electronegative atoms. These arrows also imply that there is a partial negative-charge on the more electronegative atom while there is a partial positive-charge on the less electronegative atom (hence the positive sign on the opposite end of the arrow). Electronegative atoms are attracted to electrons and pull the electrons toward its nucleus. The greater the electronegative force, the greater the pull.  

Therefore, the polarity and non-polarity of a molecule is dependent on its net-dipole. Below is a list of polar and non-polar molecules that have one or more dipole(s). BF₃ and CCl4 have three and four dipoles respectively but have a net-dipole of zero due to an equal distribution of charge or electron-pull from the fluorine and chlorine atoms. Since there is no net-dipole in these molecules, they are non-polar. In contrast, HCl, NH₃ and H₃CCl are polar because there is an unequal distribution of charge or electron-pull from the more electronegative atoms in these molecules.

 Furthermore, two important factors must be taken into consideration when evaluating a molecule’s net-dipole:

(1) The sharing and non-sharing of electrons within a molecule caused by electronegativity that results in a concentrated electron density within a specific region of a molecule, and

(2) The molecule’s symmetry or geometry as there may be multiple dipoles that net = 0 or cancel-out due to symmetry of charge.

Electronegativity:

Electronegativity describes the ability of an atom to attract electrons to its nucleus. Some atoms have a greater pull on electrons based on their electronegative values. Think about a strong magnet relative to a weak magnet. Therefore, molecules that have atoms with high electronegative values commonly demonstrate a relatively high concentration of electrons around this region within the molecule. Below is a trend in electronegative values with <1.0 being lowest and 4.0 being the highest in electronegativity:

Electron density:

Electron density is a measured probability of finding electrons in a specific location near an atom or region of a molecule. For example:

In the cartoon depicted above, the electron density is primarily located around the oxygen atom. Electrons are being pulled from the hydrogen atoms to the more electronegative oxygen atom.

Additional examples of electron density:

Notice how electron-density is balanced and apparent around H-H and F-F as these atoms equally share their electrons. H-H and F-F equally pull on their electrons as they have the same electronegative values. However, electron-density is not balanced in H-F because the bulk of electron-density is attracted to Fluorine (F). This makes sense because Fluorine is the most electronegative atom on the periodic table of elements. Fluorine pulls the electrons away from hydrogen resulting in an unequal distribution of electron-density.

Electron-density can also be equally spread or distributed over an entire region of a molecule that has a conjugated system such as benzene:

                                                                               

Molecular geometry:

Lewis structures describe atoms and molecules in 2-D while VSEPR, the valence bond model, and molecular orbital theory describe atoms and molecules in 3-D. These depictions help to predict and understand a molecule’s geometry.

Why is a molecule’s geometry important when determining polarity?

Molecular geometry is important because it contributes toward and determines the direction of a vector or a dipole-moment. This can also be visualized through symmetry or asymmetry of charge distribution.

Furthermore, there are two helpful models that are used to describe the shape and or geometry of a molecule:

-       Lewis Dot Model (structures)

-       VSEPR

Lewis Dot Model - uses electron-dot pictures that are determined by the valence electrons in the s and p electron-configuration of an atom.  This helps to visualize a 2-D molecular geometry. For example:

Oxygen (O)  = 1s² 2s² 2p⁴  = 2 + 4 = 6 valence electrons   

Boron (B) = 1s² 2s² 2p¹ = 2 + 1 = 3 valence electrons

Flourine (F) = 1s² 2s² 2p⁵ = 2 + 5 = 7 valence electrons

For example, F₂ or F-F (a homonuclear molecule) would be seen as:

                                                                       

VSEPR (Valence Shell Electron Pair Repulsion) – provides a 3-D shape and explanation for electron geometry and molecular geometry. The valence electrons in an atom’s shell are repulsive toward nearby bonds and electron lone-pairs. This repulsion keeps bonds and electrons at a maximal distance or as far apart as possible for stabilization. For example:

The five basic shapes:

-            1.    Linear geometry

-              2.  Trigonal planar geometry

-               3.  Tetrahedral geometry

-            4.   Trigonal bipyramidal geometry

-               5.  Octahedral geometry

Effects of electron pairs (lone pairs) and bonds:

A classic yet sometimes confusing example of a non-polar molecule with electronegative-charge difference is CO_2. This molecule consists of one carbon bonded to two oxygen atoms. While the oxygen atoms are much more electronegative than the carbon atom, the charge distribution is “leveled” or canceled out due to the linear geometry of CO₂ - linear cancels the two opposing dipoles out! 

Putting it all together:

-       Dipole moment

-       Electronegativity

-       Electron density

-       Molecular geometry

·      Dipole moments determine whether or not a molecule is polar or non-polar. Dipole moments can be caused by atoms that unequally share their electrons due to differences in Electronegativity.

·      This unequal sharing of electrons also causes Electron density to be concentrated and predictable in certain regions of a molecule. Electron density also contributes toward dipole-moments.

·      Finally, dipole-moments can be predicted by vectors and Molecular geometry due to symmetry or asymmetry of charge.

In conclusion, molecular polarity IS VERY IMPORTANT! It influences biological systems and how molecules interact with one another. Put simply, polar molecules tend to not interact with non-polar molecules and non-polar molecules tend to not interact with polar molecules; polar usually interacts with polar and non-polar usually interacts with non-polar. There is a commonly used phrase in organic chemistry: "like dissolves like." This statement adheres to molecular polarity.

Molecular polarity is the driving force for protein folding, the hydrophobic effect and membrane permeability to mention a few interactions with substantial biochemical significance.