Metals

The structure of metals. We're going to have a look to see how the organization of the particles in a metal gives rise to the properties that we see and we use every day. In society, metals are very important materials for construction and for many other purposes. So what we normally see is we see a lump of metal. We recognize metals very easily. We've all seen lumps of iron, lumps of aluminium. We know that car bodies are made of steel. So what does the metal actually look like if we go right the way down to the particular level? Well, we find that if we could look inside our metal, that we have a regular arrangement of ions within the metal structure. And they're arranged in rows, columns and files. It's a lattice. But what holds these ions together? The ions are all positively charged and what has to hold them together is negative electrons. Delocalized electrons. We can see down here, we've got the delocalized electrons are shown as a blue region. We know that the metals appear on the left hand side of the periodic table. This is the side of the periodic table in which the electrons are more easily lost. And so what happens is the electrons are shared over the whole body of the metal structure. We often use the term a sea of electrons. There's our word there, a sea just like the ocean, a sea of delocalized electrons all surrounding the vibrating metal ions in the lattice. And this idea can explain the. It can explain the properties that we find in the most important metals of society. If we increase the temperature of our metal, we could see the particles simply vibrate more. What happens when they vibrate more? Well, they cause each other to be pushed a little bit further apart. And this means that the whole structure of the metal expands. And we're very used to the idea of metals being very sensitive to heat and expanding. Railway lines have to have special sections put in so that in the hot summer they don't bend. Okay, so when we get to a hot enough temperature, of course, the metal will melt. The particles will fall apart from each other. They'll vibrate apart, just like a house vibrating apart in an earthquake. The temperature at which the metal melts will depend upon the force between the electrons and the positive ions in the structure. This force in turn will depend upon the magnitude of the charge of the ions. If we have a look at a simple case here, we know that the force of electrostatic attraction is proportional to Z plus times Z minus. Well, Z minus is the electrons. Z plus is the magnitude of the charge on the positive ions. If we have a Look at two metals in group one. Sorry? The metal in group one with a metal in group two, magnesium and sodium. You see that magnesium has got a double plus charge, so that means it will be surrounded by a C, which contains two electrons per atom, and they will be attracted to the double plus charge. Magnesium has got a melting point of 650 degrees, whereas sodium, which is similar in size, it's next to magnesium in the periodic table has got a melting point of 98 degrees. Far less, because the sodium ion has only got a single positive charge and the electron C has one electron per atom in it. And so consequently, the force of attraction that holds the whole lattice together in sodium metal is much weaker than the force of attraction that holds a whole lattice together in magnesium metal. Sodium has a much lower melting point. We can compare metals in the same group as well. If we have a look at group one, you'll see what happens. As we descend the group, the actual atomic radius increases. Now, this means that the single positive charge on the ion as we go down the group doesn't change. It's still a one plus charge. But that one plus charge is being attracted to the electrons in the sea of electrons much further away from where the nucleus is. And so what happens is the bonding gets a lot weaker. So as we increase the radius, we decrease the charge density and that decreases the bonding. So if we have a look at the melting points, you'll see that lithium, this is in kelvin. Now 453kelvin gradually descends down to cesium, which is only 302kelvin, or in, in real money, that's 29 degrees Celsius. Think about 29 degrees Celsius. In the context of a hot summer's day. There are many, many days in Madrid, for example, when the temperature gets to over 35 degrees. Cesium would be a liquid. So the two factors that affect the bonding strength within a metal are the magnitude of the charge on the metal ion and the number of electrons it provides to the lattice, which will be the same as the magnitude and also the charge density of the metal ion. The bigger the ion, the the weaker the forces. Now, using this model of metallic structure, we can explain the properties of metals. One of the important properties of metals is the fact that they are malleable. They can be bent into shape. We can explain this by the model is that we have the sea of delocalized electrons around some double plus charge ions. Notice the regular lattice arrangement. If we apply a force to this metal, what happens is the sea of electrons moves with the ions, the ions are displaced from their position, but the integrity of the metal itself is not, is not compromised. The metal still retains its strong forces of attraction. And that means that what we can do is we can bend the metal into shape. The conductivity of metals, which is the one property that all metals share, is easily explained with this model as well. Once again, we have the sea of electrons, this time shown in yellow. And we can see there are a few little electrons I've put in here too. Just to give us a bit more visuals on the thing. Delocalize running around over all of the metal ions. Well, they're running around randomly around the metal ions, but if we apply a potential to each side, to either side, we're going to connect the potential difference. Now, electrons can come in from the negative side, can join the sea of electrons, and then can stream out the positive side. In other words, there is free conduction of electrons between all of the ions because of the sea of electrons that is already around the ions. So this explains the conductivity of metals very, very easily. The conductivity will change as we heat up the metals because the vibrations will tend to cause more resistance to the motion as the, as the particles vibrate around the metal ions vibrate, there'll be more resistance, but still the metal will be a good conductor. Another thing to notice is that the metal itself, the conductivity of the metal, will depend upon the number of electrons that, that each metal atom provides to the delocalized C. So something like aluminium, which has got a lattice with triple positive charge ions. Each atom provides three electrons to the lattice. And for this reason, aluminium is an excellent conductor of electricity. So to summarize, a metal structure is a lattice, a regular array of ions surrounded by a sea of delocalized electrons. The strength of the force holding the, of the electrons holding the ions together depends upon the magnitude and the size of the ions themselves. The most important property of, of metals is that they conduct electricity. But they're also usually malleable and hard and with a high melting point. There are exceptions, but they usually have got a high melting point, particularly the transition metals.