Video Transcript
In this video, we will learn about
the different ways that atoms and ions can bond to form chemical compounds. We’ll discover how the outer
electrons or valence electrons and atoms are responsible for chemical bonding. And we’ll learn how to depict this
using Lewis structures.
There are 118 different elements on
the periodic table, meaning that there are 118 different types of atoms that make up
all the matter in the world around us. Well, not quite, since some of the
elements are man made, but the point remains. But when we do look at the chemical
composition of the substances that make up the world around us, there’s another
layer of complexity. We rarely see these different
elements existing as individual atoms. Rather, we’ll find that most
substances are composed of atoms or ions of an element that are attached to each
other by some kind of chemical bond.
Chemical bonds between atoms or
ions generally occur because the formation of a bond lowers the potential energy of
the atoms or ions, which makes everything more stable. This is just like how if you
release a ball at the top of a hill, it will roll down the hill and settle in a
valley where its potential energy is minimized as well. There are three main types of
chemical bonds that atoms or ions can form that generally depend on the types of
elements we’re dealing with.
Covalent bonding occurs mainly
between nonmetals. This type of bonding occurs when
atoms share their electrons to form what we call molecules. Covalent bonding is a very common
type of bonding you’ll encounter in chemistry. For example, all the atoms in water
are covalently bonded, as are all the molecules that make up the gases in the air,
oxygen, nitrogen, and carbon dioxide, as well as most of the compounds found in
living beings like proteins, fats, and carbohydrates.
The next kind of bonding is ionic
bonding. This type of bonding generally
occurs between metals and nonmetals. Ionic bonding is caused by an
electrostatic attraction between oppositely charged ions. At the atomic level, these
electrostatic attractions cause the ions to be arranged in a very orderly and
repeating pattern that we call a lattice, which you can visualize like fruit that’s
nicely stacked up at the store. At the macroscopic level, compounds
with ionic bonds usually form solids that are crystalline, like we see with NaCl,
which you have at your house as table salt.
Finally, we have metallic bonding,
which, as the name suggests, occurs when we only have metal atoms. In metallic bonding, electrons are
also shared, but in a very different way than what we saw in covalent bonding. In metals, the electrons flow
around the metal atoms, which is often referred to as a sea of electrons. Unlike what we saw in covalent
bonding where the electrons are localized in a bond between the atoms, the electrons
in a metallic bond are spread out throughout the entire metal, which we call
delocalization. The fact that electrons are free to
move around the metal atoms is why metals are generally great conductors of
electricity.
You noticed here we classified type
of bonding by the types of elements that it occurs between, metals or nonmetals,
which is generally enough to get us by. However, there are some exceptions
to this, for example, BeCl2, or beryllium chloride, because a metal, beryllium, and
a nonmetal, chlorine. But even though the bonding and
beryllium chloride is between metal and nonmetal atoms, the type of bonding between
the atoms is best described as covalent bonding, not ionic bonding like we’d
expect.
But exceptions like this are not
generally the norm. You’ll also notice that no matter
the type of bonding, bonding always has to do with electrons. In covalent bonding, the bond is
formed by sharing electrons between atoms. In ionic bonding, atoms gain or
lose electrons to form ions, which are then attracted to each other. And in metallic bonding, the
electrons become delocalized throughout the metal. Specifically, it’s the outermost
electrons that participate in chemical bonding, which we call valence electrons. The inner or core electrons
typically don’t participate in bonds because they’re too close to the nucleus and
they’re shielded by the outer electrons.
For main group elements, that is,
elements in the s- and p-blocks, or elements in groups one and two and 13 to 18, the
number of valence electrons is consistent on a group, so we can use the periodic
table to determine how many valence electrons an element has. Magnesium is located in group two
of the periodic table. So atoms of magnesium have two
valence electrons, just like atoms of all other elements that are found in group
two. Chlorine is located in group 17 of
the periodic table, so atoms of chlorine have seven valence electrons.
And since neon is located in group
18 of the periodic table, atoms of neon have eight valence electrons. The number of valence electrons an
atom has largely determines how many bonds it can form in the case of covalently
bonded atoms, or what types of ions an atom can form in the case of ionic
bonding. The scientist Gilbert Lewis
pioneered many studies on the bonding behavior of atoms. He found that, in general, when
atoms bond, they tend to lose, gain, or share electrons to attain a filled outer
shell. Since for many atoms that are
main-group elements a full outer shell is eight electrons, this behavior is often
referred to as the octet rule.
Now, this octet rule is just a rule
of thumb, and it’s quite easy to find elements that don’t obey this rule. For instance, the first electron
shell can only hold two electrons, so hydrogen and helium will have a full outer
shell when they have two electrons. Boron also commonly breaks the
octet rule. Even though it can accommodate
eight electrons in its outer shell, we commonly find it with less. And things get even less concrete
once we get into the fourth period. Nevertheless, this is still a good
rule of thumb for us to use when considering how atoms bond and form ions.
With the octet rule in mind, let’s
take a closer look at some of these electron shell diagrams that we were looking at
earlier. As we said, magnesium has two
electrons in its outer shell. So if magnesium lost those two
electrons, it could have a full outer shell, which is why magnesium and other
elements in group two of the periodic table commonly form two-plus ions. Chlorine has seven valence
electrons, just like all elements in group 17 of the periodic table, which means
that you’ll commonly see chlorine gaining one electron, since it’s only one electron
away from a full outer shell of eight electrons, which would form the Cl− ion.
Unlike magnesium, chlorine can also
bond covalently with other atoms to complete its outer shell. Here, chlorine is bonding with
hydrogen. I’ve left out the inner electrons
in chlorine just for simplicity in this diagram. This would form the molecule
HCl. Each atom participating in the bond
shares one of its electrons, which completes the outer shell for both chlorine and
hydrogen. Finally, neon has eight valence
electrons in its outer shell, which means that its outer shell is already full. This means that neon and other
elements in group 18 of the periodic table won’t tend to form ions or bonds. In other words, neon and other
noble gases are unreactive.
To depict how these valence
electrons participate in a bond, we can use something called Lewis structures. Lewis structures, named after the
scientist Gilbert Lewis, have quite a few names, like Lewis dot diagrams or electron
dot structures. But we’ll refer to them as Lewis
structures in this video. Now, we could certainly continue to
use the electron shell diagrams that we’ve been referring to throughout this
video. But drawing the inner or core
electrons over and over again when they don’t really participate in bonding can be a
hassle and can clutter the way of seeing what’s really going on.
So to draw a Lewis structure, we’ll
first write the element symbol of the atom or ion, which will represent both the
nucleus and the inner electrons. Then we’ll indicate the number of
outer or valence electrons with dots. We’ll place one dot at a time on
each side of the element symbol until we run out of valence electrons. If the atom has a lot of valence
electrons, once there is one electron on each side of the element symbol, we’ll pair
up the electrons and continue around until we run out of valence electrons. The dot placement here isn’t
important, just that we place one on each side before pairing them up.
Looking at a Lewis structure can
tell us a lot about the ions and atom we’ll tend to form or the number of bonds the
atom will tend to form. When we look at chlorine Lewis
structure, we can see that one of the electrons are unpaired. This means that it can share that
unpaired electron with another atom to form a bond, which is the conclusion we came
with earlier when we were looking at the molecule HCl. And just like we were looking at
earlier with the electron shell diagrams, the empty spot in chlorine’s Lewis
structure indicates that it can gain one electron to fill its outer shell and form
the Cl− ion.
And if we want to draw the Lewis
structure for the Cl− ion, we would want to put brackets around the Lewis structure
and put the charge of the ion outside the brackets. We can do the same thing for
magnesium. From looking at its Lewis
structure, we can tell that it has two electrons in its valence shell, which it will
lose to form the magnesium two-plus ion. So determining the Lewis structure
for an atom can tell us a lot of information. It can tell us the charge of the
ions that it’s most likely to form by seeing how close the atom is to having a full
outer shell. And the number of unpaired
electrons in the Lewis structure indicates the number of covalent bonds that that
atom can form.
And just as an aside here before we
move on to some example problems, throughout this video we’ve depicted electrons in
our diagram as dots. But electrons aren’t actually
dot-centered around the nucleus. Their behavior is quite a bit more
complicated than just simple dots. In reality, electrons are spread
out around the nucleus in what’s called an electron cloud. This electron cloud can take on
rather funny shapes, but these dot diagrams are ultimately a useful depiction to
help us keep track of bonding behavior. So now, let’s test our new
knowledge on the types of chemical bonds with some example problems.
Which of the following is not a
type of chemical bonding? (A) Covalent, (B) nuclear, (C)
metallic, or (D) ionic.
There are three main types of
chemical bonding that we’ll encounter in chemistry. The first is covalent bonding. This type of bonding is typically
found between nonmetal elements. It involves the sharing of outer
electrons to create a chemical bond. When atoms bond covalently, they
can form simple molecules like water or they can form large repeating network
covalent structures, like graphite. Then there’s ionic bonding. Ionic bonding primarily occurs
between nonmetal and metal elements. This type of bonding is caused by
an electrostatic attraction between oppositely charged ions. This is the type of bonding you see
in table salt, which has the chemical formula NaCl or sodium chloride.
The final kind of bonding is
metallic bonding, which, as the name suggests, occurs between metals. This type of bonding features
electrons that are shared between the metal nuclei that flow around the nuclei like
a sea of electrons. So as we can see, answer choices
(A), (C), and (D) all refer to types of chemical bonding. Nuclear does not. Instead, it refers to a type of
chemical process that generally results in a change to the nucleus of an atom.
The valence shell of oxygen is the
second electron shell and contains six electrons. How many covalent bonds can oxygen
form?
The valence shell is the outermost
electron shell. The electrons in the valence shell
are the electrons that are involved in the formation of bonds and ions. This is because atoms will
typically gain, lose, or share electrons and bonds so that they have a full outer
shell, which is typically eight valence electrons. The problem tells us that oxygen
has six valence electrons. But if the problem didn’t give us
this information, we would be able to determine the number of valence electrons
using the periodic table, since for main group elements, the number of valence
electrons is consistent down a group.
Either way, since oxygen has six
valence electrons, it will gain two electrons according to the octet rule so that it
can have a full valence shell. It can do this by either forming an
ion or by forming bonds. But this question specifically
asked us about how many bonds oxygen can form. When atoms bond covalently,
electrons are shared between atoms with each atom that’s participating in the bond,
giving one of its electrons to the bond. So each covalent bond that an atom
forms will effectively make it gain one electron. So oxygen can form two covalent
bonds, which is what we see in molecules that oxygen forms when it bonds covalently,
like H2O where oxygen is bonded to two hydrogens.
Now let’s finish things up with the
key points for this lesson. Chemical bonding joins atoms
together to form compounds. There are three main types of
chemical bonds: covalent bonds, which typically form between nonmetal elements,
ionic bonds, which typically form between a metal and a nonmetal, and metallic bonds
which form between metals. The valence or outer electrons are
the electrons that are responsible for creating these different types of chemical
bonds.
When atoms bond, they typically
lose, gain, or share electrons so that they could attain a full valence shell of
eight electrons. We can use Lewis structures to
depict the valence electrons for an atom or ion by drawing the element symbol to
represent the nucleus and core electrons and dots around the element symbol to
represent the valence electrons.
