Properties of metals

0:00- [Instructor] Why does a metal spoon heat up so fast when you put it inside a hot soup? Why can't we hammer a piece of aluminum into new shape without shattering it? Why are the parts that carry electricity in our devices are always almost made of metal? What is it about metals that creates all these familiar properties?

0:18Well, that's what we're gonna find out in this video, and we're gonna do that at the level of electrons and atoms, but how? Well, here's our strategy. We'll first figure out how the electrons are arranged or shared between the atoms. This will help us understand the type of bonds that are formed, and that will help us understand their physical or chemical properties, which will help us answer all these questions.

0:45So let's begin. You probably know that metals mostly live on the left side of the periodic table. And in general, these atoms have only a few electrons in their outermost energy level. For example, if you take sodium, it has just one electron in it's valence shell.

1:02So if they lose them, then look, their valence shell ends up with eight electrons giving them octet structure, which makes them stable. And that's why metals like sodium or aluminum or magnesium, whatever you take, they usually tend to lose electrons. And so whenever they react, they tend to form positive ions because they have lost electrons, and we call them cations.

1:31And so what do you think will happen when we have a pure metal, like, for example, a metal of aluminum? Well, again, metal atoms tend to lose their valence electrons. Over here, each aluminum atom has three valence electrons, so, they tend to lose them, forming cations, but these electrons are still part of the metal.

1:52They're just not tied to one specific cation anymore. So instead of thinking of each electron belonging to one atom, this is the picture. You have positive metal ions packed together, and around them is a sea of free-moving electrons. Because these electrons are not stuck to a specific location, they're mobile, we call these electrons delocalized.

2:19And so this model of positive ions in a sea of delocalized mobile electrons is what we call metallic bonding. So look, just by analyzing how the electrons are arranged in the metal, we were able to reason how they should form metallic bonding. That's pretty cool, all right?

2:38Okay, now let's see if this model can help us explain some of its physical properties. Let's start with heat and electricity conductivity. We know metals are good conductors of heat and electricity. Why? Well, because these electrons are free to move. That means when we attach a metal to a battery, for example, the battery can put a force on these electrons, and these electrons can easily move.

2:59When the electrons are moving, the charge is moving, conducting electricity. So they're good conductors of electricity. What about heat? Well, when these electrons move, they will collide and transfer kinetic energy efficiently. That's what conducting heat means, and that's why they can also conduct heat very nicely.

3:19This is why metals are usually used in wires and cookware because they're such good conductors of heat and electricity. This is also the reason why when you insert a spoon in a hot soup, it gets hot pretty quickly, good conductors of heat and electricity. Okay, now let's think about its structure.

3:38We know metals are generally solids at room temperature. Why? Why not liquids or gases? Well, to answer that question, think about the force of attraction of this metal ion. It is attracted to all the sea of electrons around it, right? This attraction packs these metal ions very strongly.

4:02As a result, if you wanted to separate these ions like this, that requires a lot of energy. And guess what? This translates to metals having a very high melting point. For example, if you look at iron, it melts at 1,540 degrees Celsius. And tungsten, for example, melts at 3,400 degrees Celsius.

4:24So because of their incredibly high melting point, metals usually tend to be solids at room temperature. Of course, there are exceptions, like, for example, if you take mercury or gallium, they're also metals, but they have low melting points, so low that they tend to be liquid at room temperature.

4:39But bearing those exceptions, we can now explain why most metals are solid at room temperature. Okay, now let's talk about why we can hammer these metals into thin sheets. The key idea is that the attraction of the metal ions with the sea of electrons is non-directional, meaning the attraction along this direction is pretty much the same along this direction or this direction.

5:01There's nothing special in one specific direction to have the force to be stronger. So think about this, what if we were to move a layer of metal ions or shift them relative to the rest? Would the metallic bonding change? Would the force of attraction on any specific metal ion change?

5:17No, it would pretty much still be the same. So even though the ions have changed their location, the metal has been deformed, the metallic bonding pretty much stays intact. This is why you can easily deform metals without breaking them, and that's why you can easily hammer a metal and deform it.

5:38And that's the reason why metals are malleable, meaning you can hammer them into thin sheets, or even ductile, you can draw them into wires. Okay, another question we could have is, why are metals so shiny? Again, this has something to do with those mobile electrons.

5:56When shine light, these electrons which are free to move can absorb energy and then reflect that light energy back, and that's what makes them so shiny. Okay, so these were the physical properties, but can we explain their chemical properties? Yeah. For that, remember, metals tend to react by losing their electrons in any chemical reaction by forming cations.

6:22Okay, but would that reactivity increase or decrease down the group? Would they get more reactive or less reactive down the group? What do you think? Well, remember down the group, atomic size increases because a new shell gets added. This means valence electrons are farther away from the nucleus and more shielded from the inner electrons.

6:44So the electrostatic attraction between them decreases down the group, and so it makes it easier for the electrons to be removed. This means the reactivity would pretty much increase down the group. That's why we would expect potassium to be more reactive, say with water, compared to sodium, which would be more reactive with water compared to lithium.

7:07Okay, now let's talk about the reaction of metals with oxygen. Why oxygen? Because it's found everywhere in the air. We know metals tend to lose electrons, and oxygen is a non-metal. And we will see in a future video that non-metals tend to gain electrons. So that means when a metal reacts with oxygen, metals lose electrons, forming positive ions, oxygen gain electrons, forming negative ions, and the positive metal ion sticks to the negative metal ion, forming what we call a metal oxide.

7:40So wait, this means if you just take a metal and just keep it in the air, it should form metal oxide all the time, right? But we usually don't notice that around us. Why not? Well, for one, because this only happens on the surface, the surface of a metal would tend to lose electrons to the air, to the oxygen in the air, forming what we call the metal oxide.

8:03But this thin layer of metal oxide formed actually protects the metal underneath from further oxidization. Now this all happens very slowly, and in some cases, this metal oxide looks very similar to the original metal. For example, aluminum oxide looks almost the same as the aluminum beneath it.

8:22That's why even though a new substance has formed on the surface, it's not always easy to notice just by looking, and that's okay. This particular layer actually helps prevent further damage, which is why aluminum foil is safe to use. But in some cases, we do notice it.

8:37For example, iron forms iron oxide, and this thin layer of ion oxide is what we call rust. But again, that's just one example of oxidization. If you look at copper, it goes from being shiny brown to blue-green patina. This is the reason why the Statue of Liberty has turned green over time.

8:55Pretty awesome, right? Okay, finally, let's considered metals reacting with acids. Remember that acid contain H+ cations, and we know metals tend to lose electrons. And so now look. The H+ cation can easily gain that negative electron, forming hydrogen gas. So when we take a metal and dip it in acid, we would expect to get a lot of hydrogen gas, and that's exactly what we see over here.

9:23We're seeing bubbling, fizzing, and of course, a lot of heat being released. And as that happens, the metal starts to dissolve. So putting it all together, metals usually have a small number of valence electrons, and so losing those electrons allows them to reach a much more stable octet configuration.

9:42That's why metals usually tend to lose electrons. That's where everything starts, okay? This means in pure metals, those valence electrons become delocalized and form a sea of mobile electrons. We call this metallic bonding, and this metallic bonding explains their conductivity, their ductility, their malleability, why they're so shiny, and you know, why they didn't have high melting points and you know, why most of them are solids at room temperature.

10:11Of course, the tendency to lose electrons also explains some of the chemical reactions, how they react with oxygen or acids, or in general, how they may react with non-metals, for that matter. I find it so awesome that just by understanding how metals tend to lose their outmost electrons, we can explain so much about them.