In this video we will discuss the relative polarity of bonds and characterize bonds as
nonpolar covalent bonds, polar covalent bonds, and ionic bonds. There is a common misconception
that bonds are either all covalent or all ionic. In actuality there is a continuum from
covalent bonds, a bond that shares electrons, to ionic bonding, transfer of electron from
one electron to another. The continuum between covalent and ionic is based on the properties
of the elements. One of these properties is electronegativity which will help us determine
whether a bond is nonpolar covalent, polar covalent, or ionic. Let's define these three
types of bonds. A nonpolar covalent bond is a bond in which the bonding electrons are
shared equally between two atoms, resulting in no charges or partial charges on the atoms
in the bond. A polar covalent bond is a bond in which the bonding electrons are shared
unequally between the two atoms resulting in partial charges on the atoms. An ionic
bond is a bond in which there is a complete transfer of one or more valence electrons
from one atom to another resulting in full charges on the resulting ions. To determine
whether something is nonpolar covalent, polar covalent, or ionic bonding we will consider
the electronegativity of the atoms in the bond. Electronegativity defined as the relative
ability of an atoms to attract electrons to itself in a chemical compound. Elements with
high electronegativities tend to have higher electron densities in chemical compounds and
are found in the upper right corner of the periodic table. Fluorine is considered the
most electronegative atom. Let's look at a few examples. Consider the diatomic molecule
chlorine. In this case we're dealing with a bond between two atoms of the same element.
In this case we would expect it to be fully nonpolar since we were dealing with the same
electronegativity for the element in the bond. Therefore the bond between the chlorines will
be considered a nonpolar covalent bond. Now let's consider hydrogen chloride HCl. In this
case we see that hydrogen is on the far left of the periodic table and chlorine is on the
far right of the periodic table. The molecule still contains a covalent bond but there is
a large electronegativity difference between the elements. This large electronegativity
difference causes there to be polar covalent bond or unequal sharing of electron density
in the bond. The chlorine in this case has a high electronegativity and attracts more
electron density to itself. This will result in a partial negative charge on the chlorine
atom. The hydrogen has a lower electronegativity. The hydrogen gives up electron density to
the chlorine resulting in a partial positive charge. Therefore the bond between the hydrogen
and the chlorine would be a polar covalent bond. Let us consider sodium chloride for
the third example. In this case the sodium has a very low electronegativity an electronegativity
value of 0.9 on most electronegativity charge. The sodium atom would give up an electron
to become a sodium ion. The chlorine has a high electronegative value of 3.0 and would
gain an electron from the sodium to become a chlorine anion. Therefore bonding in sodium
chloride is considered an ionic bond. Now notice that in this example I use the electronegativity
values of the elements that are often provided in text to determine the bonding nature between
the two elements. The difference between the electronegativity values for elements in a
bond can be used to determine whether the bond is nonpolar, polar, or ionic. As a general
rule a bond between two elements that has electronegativity difference of 0.5 or less
is generally considered a nonpolar covalent bond. Whereas a bond between elements with
a electronegativity difference in the range of 0.5 to 2 is generally considered a polar
covalent bond. A bond between elements with an electronegativity difference of 2.0 or
more is generally considered ionic. It is important to understand that this is a general
rule to determine nonpolar covalent, polar covalent, and ionic bonding nature. Let's
consider three more examples: the C-H bond in methane, the O-H bond in water, and the
sodium oxygen attraction in sodium hydroxide. For the C-H bond in methane we consider this
a nonpolar covalent bond since the electronegativity difference between the carbon and the hydrogen
is low. The carbon has an electronegativity of 2.5 and the hydrogen has an electronegativity
of 2.1. The 0.4 difference in electronegativity is less than 0.5 and therefore the bond will
be considered a nonpolar covalent bond. For the O-H bond in water the electronegativity
difference between the oxygen's electronegativity of 3.5 and the hydrogen's electronegativity
of 2.1 is larger than .5 and less than 2. Therefore the O-H bond would be considered
a polar bond resulting in a partial negative on the oxygen and a partial positive on hydrogen.
For the sodium and oxygen in sodium hydroxide there is a substantial difference between
the sodium's electronegativity of 0.9 and the oxygen's electronegativity of 3.5. In
this case the sodium atom would give up its valence electron to become a sodium cation
and the oxygen atom would gain an electron to become an anion, the hydroxide anion therefore
the electrostatic attraction between the positively charged sodium and the negatively charged
oxygen would be considered an ionic bond. We can use the electronegativity differences
to determine the characteristics of bonds either nonpolar covalent, polar covalent,
or ionic bonds.