Corrosion

Transcript

0:00Introduction

Hello everyone! Let us start with a small puzzle. Take two iron sheets. Coat one with zinc, and coat the other with tin. Both coatings are there to protect the iron, right? Now scratch both sheets, right down to the iron. The zinc coated sheet does not rust. But the tin coated sheet rusts quickly at the scratch! Same scratch, two different results? Strange, isn't it? Today we learn how iron really rusts, and how we protect it. Corrosion is a part of JEE electrochemistry. So, let's begin.

0:36Rusting: a tiny cell on the iron itself

So, what really happens when iron rusts? Shiny iron metal slowly turns into rust, a compound of iron. Tell me, is the iron oxidised, or reduced? Oxidised, yes, exactly! Iron loses electrons. And here is the key idea of this whole lesson: corrosion is an electrochemical process. It works like a small cell, running right on the surface of the metal. Okay? At one spot on the iron, called the anodic spot, iron is oxidised. F e gives F e two plus, plus two e minus. F e is iron, the arrow means gives, the two plus written up top is the charge on the iron ion, and e minus is an electron. Clear? Now, these electrons do not just sit there. They travel through the metal itself to another spot, the cathodic spot. See, both spots are on the same piece of iron. Now think. At the cathodic spot, who takes these electrons? Oxygen, correct! Oxygen is reduced, in the presence of H plus, the hydrogen ion. O two, plus four H plus, plus four e minus, gives two H two O, okay? The small two written below a symbol means two atoms of that element, so O two has two oxygen atoms, and H two O has two hydrogen atoms. And where does the H plus come from? From carbon dioxide, C O two, dissolved in the water. Got it? Now, a small trap when you write both halves together. One oxygen molecule takes four electrons. But each iron atom gives only two, right? So how many iron atoms must be oxidised for one oxygen molecule? Two, yes! So double the iron half: two F e gives two F e two plus, plus four e minus. Now the electrons lost equal the electrons gained. Students often forget to match them, so watch it! With me? And the story does not end there. F e two plus is oxidised further, to hydrated iron oxide: F e two O three, dot, x H two O, as you see on screen. The dot means water molecules attached to it, and x means the number of water molecules can vary. This is rust.

2:54What makes rusting faster?

Right? Alright, what makes rusting faster? Remember, the reaction at the cathodic spot needs H plus ions. So in acidic water, what do we have more of? More H plus, yes! More hydrogen ions for the cathodic reaction, so corrosion is faster in acidic water. Okay? Next, salty water. Salt water conducts electricity better. It is a better electrolyte, so our little corrosion cell works faster. That is why iron near the sea rusts faster. Clear? And third, corrosion is faster where two different metals are in contact. Here, it is the more easily oxidised metal of the two that corrodes faster. Keep this one in mind. It is the key to our puzzle from the start. Right?

3:43Protecting iron

Now, how do we protect iron? The simplest way is a barrier coating: paint, oil or grease. It simply covers the iron. Got it? But think. What happens when the paint gets a scratch? The bare iron at the scratch is open again, and the barrier cannot help there. So we use something smarter. Galvanising: coating the iron with zinc, Z n. And sacrificial anodes: blocks of magnesium, M g, or zinc, fixed to the iron. With me? These two protect the iron even when scratched. Why? Because zinc and magnesium are more easily oxidised than iron. They give up their own metal, and the iron is saved. That is why we call them sacrificial. Okay? Another way is alloying: mixing iron with other metals, as in stainless steel. Right? And finally, cathodic protection with an impressed current. Impressed means pushed in from an outside source. The name tells you the idea: the iron is made the cathode, so it is not the one being oxidised. Clear?

4:52Example 1: why magnesium blocks protect ship hulls

Now let's try one. Why are magnesium blocks fixed to the hull of a ship, or to buried iron pipes? The hull is the body of the ship, the part that sits in the water. Pause the video, and try it first. Done? Okay, let's check. Magnesium is more easily oxidised than iron, remember? So the magnesium block acts as a sacrificial anode. Got it? Magnesium is oxidised in place of the iron. The iron is left as the cathode, so it is protected. This is cathodic protection. Oh no, careful though! This works only because magnesium is more easily oxidised than iron. Use a less reactive metal, and the trick fails. In the next example, we will see exactly what happens then. Okay?

5:41Example 2: scratched galvanised iron vs tinned iron

And now, our puzzle from the start. Why does a scratched galvanised iron sheet, coated with zinc, not rust, while scratched tinned iron rusts quickly? Pause the video and try it. Hint: compare the E naught values. Done? Okay, let's check. We need the standard reduction potentials. E naught: that is E with a small circle on top, which means standard. These are measured at two hundred and ninety eight kelvin, with ions at one molar and gases at one bar, and always written as reductions. Clear? Zinc, Z n, minus zero point seven six volts. Iron, F e, minus zero point four four volts. And tin, S n, minus zero point one four volts. The capital V stands for volt. Now, which of the three is most easily oxidised? Zinc, perfect! The lower, more negative, the E naught, the more easily the metal is oxidised. Minus zero point seven six is lower than minus zero point four four. So zinc is more easily oxidised than iron, and zinc corrodes sacrificially, even at a scratch. The iron stays safe. Got it? Now tin. Tin is also a coating on the iron. So it should protect the iron the same way, right? No, no, careful! This is the trap. Tin is minus zero point one four volts. That is higher than iron's minus zero point four four. So tin is less reactive than iron. Now it is the iron that is more easily oxidised. With me? So at a scratch, tin and iron are two different metals in contact, remember? And iron becomes the anode. Iron is oxidised, and it corrodes faster. Puzzle solved! Zinc gives itself up for the iron; tin makes the iron the anode. See how simple it is?

7:30Quick revision: rusting cell, faster rusting, Zn and Mg

So, quick revision. Rusting: iron is oxidised at anodic spots, and oxygen is reduced with H plus at cathodic spots. Acidic water, salt water, and two metals in contact make it faster. And zinc and magnesium protect iron because they are more easily oxidised than iron. Now go and try the practice problems on this page. You will find them easy. All the best!