Video Transcript
In this video, we will learn how to
explain the chemical property of electronegativity. When atoms of nonmetals react with
each other and bond together, covalent bonds are normally formed. These covalent bonds are widely
seen in simple molecules and they hold their atoms together.
So what causes the atoms to be
attracted to one another? Quite simply, it’s the
electrostatic attraction between the shared pair of electrons and the positively
charged nucleus of each atom in the covalent bond. The more electrostatic attraction
we have, the stronger this covalent bond will be. The electrostatic attraction is
largely governed by the size of the positive charges in each nucleus. Since the positive charge in each
nucleus is caused by positively charged protons, its size is determined by how many
protons there are in each nucleus. It’s not just this simple, though,
as the size of each atom is important, too. If the two atoms are large, the
distance between the nucleus of each atom is greater. This makes the electrostatic
attraction weaker. And we get a weaker covalent
bond.
Remember that a shared pair of
electrons is just a shared pair of electrons, and the negative charge of these two
electrons does not change. If the atoms involved in the
covalent bond are identical, the attraction for the shared pair of electrons will
also be identical. We could imagine this situation to
arise in a molecule of chlorine. The single covalent bond formed in
a chlorine molecule contains a shared pair of electrons. One electron comes from each
chlorine atom, respectively, shown by a dot and a cross here. So in this chlorine molecule, we
can imagine that there is perfect sharing of the pair of electrons in this covalent
bond, since the number of protons in each nucleus is identical and the size of each
chlorine atom is the same.
This perfect sharing scenario will
only arise in covalent bonds formed between two identical atoms in a diatomic
molecule. Think of H O N Cl Br I F to
remember the most commonly encountered elements that exist as diatomic
molecules. We should find perfect sharing of
the electron pair in the covalent bonds of hydrogen, oxygen, nitrogen, chlorine,
bromine, iodine, and fluorine.
We can imagine the attraction
between the shared pair of electrons and the nucleus of each atom in a covalent bond
as being similar to a tug of war. In a tug of war, two teams pull on
opposite ends of a strong rope with a marker tied in the middle. The team with the biggest pulling
power or traction wins by pulling the rope marker over a line on the ground
first. Of course, covalent bonds do not
involve pieces of rope and teams of people. But imagine that the nucleus of
each atom is like the team of people and the rope marker is the shared pair of
electrons. It is reasonable to assume that a
larger nucleus with more protons will behave like a larger or stronger team of
people in a tug of war. A larger nucleus will attract the
shared pair of electrons in a covalent bond more strongly than a smaller nucleus
with less protons.
When two different atoms are bonded
together in a covalent bond, we may see uneven sharing or less-than-perfect sharing
of the electron pair in the covalent bond. This could occur if a hydrogen atom
is covalently bonded to a chlorine atom. Chlorine has a much larger nucleus
than hydrogen. Chlorine has 17 protons or positive
charges in its nucleus versus hydrogen’s one. The total amount of positive charge
found in a nucleus is often referred to as the nuclear charge. Since the chlorine atom has a
bigger nuclear charge than the hydrogen atom, the sharing of the electron pair in
the covalent bond is not equal. The electron pair is attracted more
strongly to the chlorine atom. The concept of how well electrons
are shared in covalent bonds is called electronegativity.
Electronegativity is now defined as
the relative attraction that an atom has for the shared pair of electrons in a
covalent bond. This concept was developed using a
scale by the American scientist Linus Pauling in 1932. Linus Pauling called
electronegativity the power of an atom in a molecule to attract electrons to
itself. Linus Pauling introduced the scale
for electronegativity. He calculated relative values for
each element, the calculation involving bond dissociation energies. Bond dissociation energies give an
idea of how much energy is needed to break a specific covalent bond. The Pauling electronegativity scale
has no units, and the numbers range from 0.79 for cesium and francium to 4.0 for
fluorine.
There are definite patterns for the
electronegativity values for each element as we move around the periodic table. Electronegativity displays periodic
trends. As we move across a period from
left to right in the periodic table, electronegativity values generally
increase. If we exclude the transition
elements where there’s not always a clear trend, electronegativity values increase
across each period. Across period two, we see values
range from 1.0 for lithium to 4.0 for fluorine, which is in fact the most
electronegative element of all. A similar pattern emerges for
period three elements sodium to chlorine. Here, the values range from 0.9 for
sodium to 3.2 for chlorine.
In any given period, the
corresponding halogen in group 17 or seven emerges as the most electronegative
element. This is because the number of
protons in the nucleus, or nuclear charge, is increasing as we move across each
period. There are no new shells being
occupied, so no additional screening or shielding of the valence electrons
occurs. Screening or shielding is when
occupied inner shells reduce the effect of the nuclear charge on the valence shell
electrons. A greater nuclear charge attracts
bonding pairs of electrons more strongly. The increasing nuclear charge
across each period also reduces the size of the atoms, so atomic radii decrease. This also adds to the attraction of
valence shell electrons as they’re closer to the nucleus.
All of these factors explain why
electronegativity values increase as we move across a period in the periodic
table. As we move down each group in the
periodic table, electronegativity values decrease. This trend is seen in group one,
where values range from 2.2 for hydrogen to 0.7 for francium. This trend is repeated as we move
down group 17, where values range from 4.0 for fluorine to 2.2 for astatine. So as we move down each group in
the periodic table, electronegativity values decrease. This is because we are adding
occupied shells and the atomic radii or the size of the atoms are getting
larger.
Although the nuclear charge is
increasing as we move down the group, its effect is shielded by the core
electrons. These factors combined explain why
electronegativity decreases as you move down a group. So on a map of the periodic table,
the least electronegative elements are found on the lower-left corner and the most
electronegative elements on the upper-right corner. It is worth noting here that since
noble gases do not generally form compounds, they are not assigned electronegativity
values.
It’s not really too important to
learn electronegativity values for each element. We can use them, however, to
determine whether specified covalent bonds will exhibit perfect sharing or not. For this purpose, it’s more
important to look at differences in electronegativity values. If two elements with similar
electronegativity values are covalently bonded together, we can see that there is
little to no difference in the electronegativity values. An example of this is a
carbon-to-hydrogen covalent bond. Carbon has an electronegativity
value of 2.6. Hydrogen has an electronegativity
value of 2.2. The difference in electronegativity
is 2.6 minus 2.2, which equals 0.4. This difference is considered to be
very small and the bond is said to be a nonpolar covalent bond. There is no difference in the
sharing of the electron pair between the nucleus of each atom in this covalent
bond.
Notice the difference of less than
0.5 is considered to give rise to a nonpolar covalent bond. In contrast, if we examine the
hydrogen-to-fluorine covalent bond, we find that there is a difference in
electronegativity of 4.0 minus 2.2, which equals 1.8. This difference is considered to be
large and will give rise to a highly polar covalent bond. The sharing of the electron pair in
this covalent bond is very uneven. Using electronegativity value
differences, covalent bonds can be considered to lie on a spectrum from nonpolar,
perfect sharing, to highly polar, very uneven sharing and in fact almost ionic.
The differences in
electronegativity values are summed up in this table with descriptions for the type
of bond expected. We can see that if the difference
in electronegativity of the two atoms bonded together lies at the extreme end, for
example, rubidium bonded to fluorine, then the bonding is not expected to be
covalent at all. In this case, there is electron
transfer between rubidium and fluorine, resulting in a fully ionic bond. So it’s more useful to consider
differences in electronegativity when determining the nature of the sharing of the
electron pair in a covalent bond and, therefore, how polarized that covalent bond
would be expected to be. We will now look a question to test
your understanding of the concept of electronegativity.
Why is the element potassium less
electronegative than the element lithium? (A) Potassium has a greater number
of protons and greater nuclear charge to attract the bonding electrons. (B) Lithium has less electrons than
potassium. (C) Potassium does not follow the
general trend of electronegativity. (D) Bonding electrons are further
from the nucleus in atom of potassium causing less attraction. (E) Electronegativity increases as
you descend a group in the periodic table.
In this question, we’re being asked
to compare the electronegativity of two elements. Remember that electronegativity is
the attraction that an atom has for the shared pair of electrons in a covalent
bond. To answer this question, we’ll
check the validity of each statement against the main factors that influence
electronegativity. Electronegativity depends on the
number of protons in the nucleus of an atom, sometimes referred to as the nuclear
charge. It also depends on the number of
shells occupied, which influences the size of an atom and also the amount of
shielding or screening taking place in that atom.
As we move across a period in the
periodic table, electronegativity values generally increase. This is due to increasing nuclear
charge, the same number of shells being occupied, and therefore no additional
screening or shielding. As we move down a group,
electronegativity values decrease. Although there are more protons in
the nucleus of these atoms, the number of shells being occupied is increasing. This increases the amount of
shielding or screening. And this factor outweighs the
increasing nuclear charge.
We’re comparing potassium, symbol
K, with the element lithium, symbol Li. Both metals are located in group
one of the periodic table. And since potassium is further down
the group than lithium, the element potassium is certainly less electronegative than
lithium. Statement (A) suggests that
potassium has a greater number of protons and hence nuclear charge than lithium. This part of the statement is
true. Lithium has three protons in its
nucleus; potassium has 19. This statement could be a good
argument to explain why potassium is more electronegative than lithium. This argument would be tempting
because electronegativity does depend on the size of the nuclear charge. This statement does not provide any
support, though, as to why potassium is less electronegative than lithium. It’s therefore not a correct
answer.
Despite the greater nuclear charge
in potassium, we find there’s more shielding in this atom, and this outweighs the
factor of the increasing nuclear charge as far as electronegativity is
concerned. Statement (B) suggests that lithium
has less electrons than potassium. On the face of it, this statement
is true. Lithium contains three electrons,
and potassium contains 19. Simply stating how many electrons
each atom has does not attempt to explain how many shells are occupied or the degree
of shielding in each atom. There’s also no mention of the
nuclear charge in each atom in this statement. And therefore, it doesn’t provide
an explanation at all. Statement (B) is not the correct
answer.
Statement (C) suggests that
potassium doesn’t follow the general trend of electronegativity patterns in the
periodic table. In fact, the elements in group one
do follow the pattern for electronegativity. As we descend group one,
electronegativity values steadily decrease. This pattern is not always observed
throughout the transition elements, but that’s not what we’re dealing with here. Statement (C) is therefore not a
correct answer.
Statement (D) suggests that the
bonding electrons are further from the nucleus in an atom of potassium. This is a true statement so far as
a potassium atom having more shells occupied is a larger atom. Since the valence electrons are
further from the nucleus, there’ll be less attraction for them too. This will indeed reduce the
electronegativity of the atom and it does form part of an explanation. Statement (D) could be the correct
answer. Statement (E) offers a full
statement. Electronegativity does not increase
as you descend a group in the periodic table. Statement (E) does not offer an
explanation either. So it’s not the correct answer. Statement (D) is the correct
answer.
Now let’s review the key
points. Electronegativity is the relative
attraction an atom has for the shared pair of electrons in a covalent bond. Electronegativity increases across
a period as the nuclear charge increases with no extra shielding. Electronegativity decreases down a
group as there are more shells and more shielding. Large differences in
electronegativity of covalently bonded atoms leads to highly polar bonds.
