Welcome to our first lesson in our introduction to material structure Ansys Innovation Course.
Today we'll be focusing on atomic bonds,
the foundation of solid materials.
We'll be discussing three types of primary bonds: covalent,
ionic, and metallic. And we'll also be discussing secondary bonds like Van der Waals and Hydrogen.
But before we start talking about the bonds,
we should really be talking about the atoms themselves.
The elements of our Periodic table are our building blocks for materials, so to speak,
as materials scientists and engineers. Now all of the elements circled here
can be found in metallic materials. That's almost the entire Periodic
Table! And this impacts the types of bonds we will see these elements involved in.
Now that we've talked a little bit about the building blocks, let's talk about how they bond
together, because almost all of our atoms want to stick together. Let's think of two atoms, coming
closer and closer together. When the atoms are close together we have two forces at play here,
repulsive and attractive. These forces are in balance and equal zero if we have atomic bonds.
This gives us our interatomic distance of that bond. Now
I think it's often easier to think of this in terms of energy,
which is closely related to force. Energy is the derivative of force with respect
to atomic separation so we can remake our graph here in terms of energy and interatomic distance.
When our forces are in balance, we have an energy minimum. This makes sense that our bonds allow the
atoms to have some decrease in energy-- nature is always trying to exist in the lowest energy
state. So what are these forces that are at play as our atoms bond together. How do they change
over different bonds? Let's start with ionic bonds as they're often the easiest to visualize.
These are composed of one non-metallic and one metallic element. For an example we'll
use sodium chloride, which is the composition of regular table salt. Sodium, our metallic atom,
can more readily give up its valence electrons compared to our non-metallic chlorine atom.
When the sodium transfers its electrons towards the chlorine, both atoms have a complete valence
shell-- a more energetically favorable state. The atoms are also both ionic in character
Because we have one positive and one negative ion, large attractive Coulumbic forces are at work
But our ions can't get too close to one another.
If the electron shells overlap too much, large repulsive forces will arise.
Ionic bonds are generally non-directional,
meaning we have a lot of flexibility in how we can arrange our atoms in our bulk material
so long as our net charge equals zero.
What about covalent bonds? Covalent bonds occur in many molecules and materials.
The atoms involved often have small differences in electronegativity.
For covalent bonds, we have an overlap of joined electron shells to be formed
via the sharing of electrons. Each atom will contribute at least one electron to the bond.
This hybridization, especially in carbon materials, is complex and alas we don't have
time to discuss the details here but this phenomenon does lead to a directional bond
that you would have seen as sticks in molecular models at school. This impacts how the atoms can
arrange themselves in crystals. Covalent bonds can be found in pure materials like diamond,
silicon, and germanium. They dominate in silicate ceramics and glasses and contribute
to high melting temperature metallic bonds. They're can also be found in the backbones of
polymeric chains, which we will see in a later lesson.
The last primary bond type we'll discuss is the metallic bond. Now we saw earlier that
most elements can participate in these bonds. Metallic bonds are unique in that instead of
electrons belonging to one or two atoms through transfer or sharing, they delocalize and are free
to move throughout the material. This cloud or "sea" of electrons helps shield the positive cores
of the atoms from one another and are responsible for the electric and thermal conductivity found
in metals. These delocalized electrons move very freely within the material and have no direction,
acting like a glue allowing metals to have a highly compact crystal structure
which we'll see in later lessons. Now we can move on to secondary bonds. Some of you may be
familiar with Van der Waals bonds. Most of you may be familiar with hydrogen bonds, which are
a form of Van der Waals bonds. Van der Waals are considered secondary bonds due to their
significantly weaker bond energies compared to our primary bonds. Van der Waals bonds occur between
dipoles within uncharged atoms or molecules. A dipole exists when there is a separation
of the positive and negative portions within a compound or molecule like shown here. As we
saw in ionic bonds, a negative and positive charge lead to Coulombic attraction. Hydrogen bonds occur
within these dipoles in molecules that contain hydrogen, like water or some polymeric materials.
More on that in lesson three. So we've defined our primary and secondary bonds--
we'll see how these impact a material's microstructure in later lessons
But let's look at some examples of how our bonds impact
a bulk material's properties.
The first you'll already be familiar with if you watched our introduction to material elasticity
course and that's how a bond's stiffness impacts a material’s elastic modulus.
When looking at a stress strain curve, we can see that different material families
have different elastic responses. This boils down to the fact that our different
material families have different bonds and these bonds have different stiffnesses.
If we compare ceramics and polymers together-- Ceramics often have ionic bonds, which are much
stiffer than the Van der Waals bonds that hold polymeric chains together. And we can see this
reflected in the macroscopic elastic behavior of our material families.
Another example of an impact that bonding can have on bulk properties is melting temperature.
When we're melting a material, the bonds must break. Since our bonds are allowing our atoms
to exist at some lower energy state, it makes sense that we need to provide energy to the
system in order to break those bonds and overcome that energy barrier. That's why
you heat up materials when you're melting them. This heat provides kinetic energy in the system,
allowing the atoms to move farther and farther apart, overcoming those attractive forces.
Now, covalent bonds are the strongest bonds that we've talked about today--it takes a lot of energy
to break them. That's why materials like diamond and carbon fiber simply evaporate instead of melt
and silicon and germanium have such high melting temperatures; they're made out of covalent bonds.
We had also mentioned that some high temperature metallic materials have covalent bond-like
characteristics. This high melting temeprature is because of these characteristics. And with that,
we've reached the end of our lesson. We've layed the foundations, defining our atomic bonds based
on general characteristics and seeing some of the impacts they can have on bulk material properties.
Next, we move on to crystalline materials. Specifically, we'll be comparing crystalline
and amorphous materials together, based on their atomic arrangement and defining what a unit cell
in a crystalline material. So join us, as we learn more about material structure.