Electrons in covalent compounds are rarely exactly, equally "shared" by all the atoms. Rather the electrons have a greater relative attraction for some of the atoms compared to others. As a result, the electrons are not equally distributed throughout the molecule. This results in partial negative charges in some regions (those with a greater pull of electrons) and corresponding partial positive charges in the regions that are deficient in electrons. We determine where these charges are by comparing electronegativities.
The electronegativity of an element describes its relative "pull" of electrons in a covalent compound compared to other elements. High electronegativity elements will attract electrons more than low electronegativity elements.
The resulting covalent bonds then have a "polarity." That is to say one end is positive and one end is negative. This polarity is described quantitatively as a dipole moment. This is a measure of two charges separated by a particular distance.
Depending on the geometry of a molecule, the entire molecule may have a dipole moment. This is a quantity that we can measure in the lab as an observable. Molecules with large dipole moments we call "polar." Those with very small or no dipole we call "non-polar."