Ionic structure

Ionic bonding happens when a metal bonds to a non metal. Usually, although there are exceptions in this particular case, we'll just consider the bonding between a metal and a non metal. Metals can lose their outer electrons very easily. That's what makes them metallic. In a metal structure, the outer electrons act as a sea of delocalized electrons that hold the metal ions together. But the fact that they have fairly low ionization energies means that they can lose electrons to something to another particle. That has a requirement for electrons, such as a non metal atom, for example, group 17 or group 16 element, which has a high electron affinity. In other words, it releases energy when it gains that electron. So the energy released by accepting an electron by group 17 compensates for the energy needed that was to remove the electron from the metal. However, that's not the most important energy change. The most important energy change is what happens after the ions are formed. They are then attracted to one another by a strong electrostatic attraction. And this is a non directional force. And this causes the ions to build up virtually infinite lattice, or as long as there are ions to join the lattice to form a regular array. And we see that in the macroscopic world as a crystal. On the screen you can see the salt crystal. The salt crystal looks just like a cube. But were we able to look inside our salt crystal, we would see that there is a regular of array of positive and negative ions all vibrating about one another at room temperature. Obviously they need some kinetic energy. And this structure is highly ordered. There are many, many strong forces of electrostatic attraction holding the ions together. Now, now each individual ion, each negative ion, as we can see over here, let me just take the microscope over here. Each negative ion is surrounded usually by six positive ions. That means it's held in place by six strong forces of electrostatic attraction. And likewise, each positive ion is surrounded by at least six negative ions. So once again, there are six strong electrostatic forces on holding the positive ions in place. The whole structure together then is one enormous mess of positive and negative attractions going on. And that's what makes the macroscopic properties of ionic substances so unique. Now we can have a look at the macroscopic properties. Let me just drop the microscope over there, or the magnifying glass rather. We have this regular arrangement which we call a lattice. I'm just moving this out of the way because it's a nuisance. The properties that we see in the macroscopic world, they have high melting points. This is because it requires a huge amount of energy to actually separate the positive ions from the negative Ions. If we have a look at this picture here, we can see us. Take thy magnifying glass. We'll look into our structure again here. This represents sodium chloride. Sodium chloride has got a melting point of 801 degrees C. If you think about a sodium atom and a chlorine atom, obviously not the same as sodium chloride ionic structure. But were it to be a molecule, then it would have a relative mass of 23 for sodium, 35.5 for chlorine, which would make 58.5, rather something about the same size relative mass wise as an ethanol molecule. And ethanol, of course, is a liquid at room temperature, whereas to melt sodium chloride, we require a temperature of about 801 degrees. So by putting energy in, we increase the temperature. All that happens is the vibrations of the ions within the lattice get greater and greater and greater until eventually we reach up to the melting point. If I can get my thermometer up to the melting point. There we go. At this stage, the particles have shaken themselves apart, rather like a large building being shaken apart by an earthquake. In this case, the earthquake is the energy that I've put in, the kinetic energy of the temperature. If the temperature cools down again below the melting point, the particles reform and find their ways back into the crystal lattice. Very slow. Doing this. There we go. And there they are back in the crystal lattice at 789 degrees Celsius. Okay. The actual melting point, of course, will depend upon the force of attraction between the negative and positive charges. I'm just going to put the temperature down because this vibration is making me crazy. There we go. Let's get it nice and calm. There we are. Much calmer down there. 86 degrees. Let's take it to room temperature. Oh, colder. Minus 22. It's still vibrating, of course, because it's still got some thermal energy. Okay. The actual melting point will depend upon the force of electrostatic attraction. And this depends on two things. The most important of these two things is the magnitude of the charge. Now, if we have a look at, say, a sodium chloride organization, we've got single positive charges attracted to single negative charges. Remember, our force of electrostatic attraction is proportional to Z plus times Z minus over R squared. Well, the plus and the minus are both one in magnitude and R squared is the distance between them or the internuclear distance. So if we replace our sodium chloride by magnesium oxide, we've now got a 2 plus times 2 minus in our equation. And so the numerator of the equation becomes four times larger. And this Means there's four times as much force required. And so, correspondingly, the melting point of Magnesium oxide is 3,100 degrees very, very hot indeed. So this is the most important factor is the magnitude of the charge. The, the other factor, which we'll go back to here, is the radius of the ions. If the particles can get closer together, well, the force of electrostatic attraction is greater. So we can compare sodium chloride with a melting point of 800:1 with potassium chloride, where potassium is a larger ion than sodium. And this means that the r, the distance, the internuclear distance, is going to be correspondingly larger and therefore the denominator will be larger and the actual force is less. And so the melting point of potassium chloride is a little less than that of sodium chloride. Notice that this factor is much less important than the charge on the ions. What other properties do ionic crystals have? Well, the fact that the ions themselves are charged, the means that they're very attractive for ionic solvents or solvents which are highly polar, such as water. And this is why many ionic compounds dissolve in water. The solubility in water is fairly straightforward. The water molecules surround the ions. The hydrogen of water is partially positive. It will get as close as it can to the negative charged particles, whereas the lone pairs on the oxygen will get as close as it can to the positively charged ions. And this way, the water molecules can carry the ions off into solution. It's not such a straightforward thing because it depends upon several factors whether this actually happens. There are three factors in the lattice enthalpy, the hydration enthalpy of the ions, and the entropy change on dissolution. We can assume that the entropy change on dissolution is pretty much the same for most ionic compounds, because we're going from a lattice which is highly ordered to a solution which is much less ordered. So really it's a balance between the lattice enthalpy, which is the force of attraction of the ions for each other, and the hydration enthalpy, which is the force of attraction or the energy released when the ions get caged up by water molecules. We're going more into this when we study energetics. As I said, the prediction of solubility is difficult because these two factors tend to be influenced by the same. The same factors, the same things. Let's have a look at another property of ionic crystals. They are brittle. Brittle is the opposite of malleable. It means they tend to break rather than bend. And if we have a look at a two dimensional picture of our ionic crystal, you've got the sodium and the chloride ions attracted together in two dimensions. If I produce stress on this structure by pushing the arrow, what happens is the similar charged ions come into contact, and that causes the whole structure to break along a line. We call this line a cleavage plane. Have a look again. So as the arrow pushes the ions on the top row to the right, you'll see the negatives get close to negatives at the same time as positives get close to positives, and that's when the structure breaks. So this is the cleavage of an ionic crystal. And so they tend to snap or to crack or to split apart rather than bend. We say they're brittle. And the final property of ionic materials is the fact that they're good electrical conductors when they are molten, but they do not conduct when they're in the solid form. For something to be able to conduct electricity, we have to have charged particles that are free to move. Clearly, there are plenty of charged particles in the ionic crystal, but they're not free to move. They're all held strongly in fixed positions or vibrating about the same position in the solid. But when we melt our ionic material, even if it melts at a high temperature, the ions become free to move, and they can then carry charge. Similarly, if we dissolve our ionic material in water, then the ions once again become free to move. And so a solution of an ionic crystal or ionic compound in water is also an electrolyte. The word electrolyte means a liquid that can carry electricity using ions. So let's go back to the original picture, the original proposal, if you like, that our ionic crystal is a regular, giant array of oppositely charged ions arranged in a lattice or network with many, many strong forces of attraction. And this is what produces the macroscopic properties that we measure in the big, wide.