2.6A Ionic Bonds Part 1

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Do curso por Georgia Institute of Technology

Material Behavior

256 ratings

Georgia Institute of Technology

Material Behavior

256 ratings

Have you ever wondered why ceramics are hard and brittle while metals tend to be ductile? Why some materials conduct heat or electricity while others are insulators? Why adding just a small amount of carbon to iron results in an alloy that is so much stronger than the base metal? In this course, you will learn how a material’s properties are determined by the microstructure of the material, which is in turn determined by composition and the processing that the material has undergone. This is the first of three Coursera courses that mirror the Introduction to Materials Science class that is taken by most engineering undergrads at Georgia Tech. The aim of the course is to help students better understand the engineering materials that are used in the world around them. This first section covers the fundamentals of materials science including atomic structure and bonding, crystal structure, atomic and microscopic defects, and noncrystalline materials such as glasses, rubbers, and polymers.

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Atomic Structure and Bonding [Difficulty: Easy || Student Effort: 2hrs]

In this module, we will discuss the structure of the atom, how atoms interact with each other, and how those interactions affect material properties. We will explore how the types of atoms present in a material determine what kind of bonding occurs, what differentiates the three types of primary bonds - metallic, ionic, and covalent, and the implications of the type of bonding on the material microstructure. You will learn how atoms arrange themselves as a natural result of their size and bonding. This knowledge will provide you with a foundation for understanding the relationship between a material's microstructure and its properties.

2.5 Primary Bonds5:33

2.6A Ionic Bonds Part 17:27

2.6B Ionic Bonds Part 27:55

2.6C Ionic Bonds Part 35:04

Conheça os instrutores

Thomas H. Sanders, Jr.
Regents' Professor
School of Materials Science and Engineering

In this lesson we're going to introduce the ionic bonds and

how those ionic bonds are formed.

If we start out with a sodium ion, and remember that sodium is all the way

to the left of the periodic chart, it has one electron in the s shell.

And, if we take that electron away,

we have satisfied the inner core of electrons.

And as a result of stripping away that electron,

the cation given by the NA+ winds up being smaller because

of the contraction of the cloud associated with the missing of that electron.

The process that is related to the change in the Na to

the Na plus one cation requires a certain amount of energy.

And that energy is referred to as the ionization potential.

So what we're doing here is taking the sodium,

making it a cation, and it requires a certain amount of energy.

Now, we go over to the right-hand side of the periodic chart

where we have the halogens.

And, in this case,

we're looking at chlorine as a typical example of the halogens.

And the neutral chlorine atom will have a specific atomic radius.

If we add an electron, what we wind up doing, then,

is to increase the radius of the electron cloud.

The process of going from the neutral chlorine atom to the chlorine anion

winds up providing some energy to the system, and

we refer to this as the Electron affinity.

So the amount of energy that's released,

as associated with the chlorine neutral atom to chloride anion,

turns out to be on the order of about 4 electronvolts.

So the sum of these two processes,

where we take and make a sodium ion and a chlorine ion,

actually requires an amount of energy that is associated

with the difference in those two energy products.

If we look at the periodic chart, one of the things that we'll see is,

as we go down the periodic chart in a column, so

let's take a look again at the first column where we have sodium, potassium,

rubidium, as we go down that column with increased the atomic number,

what we see is the neutral atom winds up increasing.

And that is given by the symbol of a green atom.

If we ionize, and we strip an electron from any one

of those cations along that column, we find that we reduce the atom.

Going from the atom to the cation, there is a reduction in radius.

And if we look at all of those associated green and reds,

what we see is the corresponding change for the different elements,

where we have the process where we go from the neutral atom to the cation.

On the other side of the periodic chart, where our tendency is to add electrons,

what we then do is to start out with the green atom, and

as a result of creating the anion,

we wind up increasing the size of the anion over the size of the neutral atom.

Now, in addition to seeing a change associated with

the addition of one electron or the removal of one electron,

if we look at an element like iron, for example, we can see that the relative size

of the atom changes as we go through and consider different valence states of iron.

So in the first case we have iron as a neutral atom.

And over to the right, when iron is in a plus 2 configuration,

that is we have lost two electrons, there is an associated reduction.

If we go a little bit further, where we have a second

oxidation state where iron is in the 3+ state,

we see that that progressive change in the number of electrons that we have removed

winds up decreasing the size of the atom into the cations that emerge.

Now, when we look at the periodic chart, what we see is as we go down the columns,

we will see a tendency for the size of the atoms to wind up increasing.

And if we look at the elements in the first column, those are the Alkali metals,

and we look at the second column, those are the Alkaline earths.

These are the elements that are associated with the s electrons.

Now, as we go across the periodic chart,

what we find is there is an increase in the electronegativity.

That is, as we go further to the right,

we being to see the addition of electrons to the associated

neutral atom to wind up forming the anions.

We discussed in this module what happens when we consider

the elements in the group carbon, silicon, and germanium,

those that are associated with the sp3 hybridized orbitals, and

the characteristics that are associated with the diamond cubic,

crystal structures of diamond, as well as silicon and germanium, and

ultimately to their properties that are attributed to the semiconductor behavior.

Now, when we come all the way over to the far right we have that halogen column.

And that halogen column sitting right next to the inner column,

we find that it is very easy for us to add an electron to stabilize that outer orbit.

And then, when we look at the elements that are in that last column,

those are the Noble gases.

Those are the ones that are associated with stable electron configurations.

And if we look at the bottom of the periodic chart, where we have taken

out the rare earths, and now what we're considering with is the behavior of

those rare earths that are controlled by the filling of those outer f orbitals,

going from cerium all the way across for the rare earths.

In the next lesson, we're going to be describing the ionic bond.

And with the ionic bond, there's some very simple models that we can use when

we describe the ionic bond that will lead to an understanding of the development

of bond force curves, and ultimately the development of the bond energy curves.

These two behaviors we will be able to begin to understand

some of the properties and characteristics that are associated with the material

that we'll be dealing with in the remainder of this course.

Thank you.

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