Lesson video

Hello, my name's Dr.

Warren.

I'm so pleased that you can join me today for this lesson on, "Corrosion and its Prevention: Including Half Equations".

It's part of the industrial chemistry unit.

I'm here to work with you throughout the lesson, support you all the way, especially through the tricky parts.

The learning outcome for today's lesson is, I can describe corrosion in detail, including chemical equations and describe what actions can be taken to prevent corrosion.

And we have some key words in today's lesson including: rusting, corrosion, sacrificial protection, electroplating and tarnishing.

Now we're going to look at these keywords in some sentences.

Rusting refers to the corrosion of iron or steel, in the presence of water and oxygen.

Corrosion is the gradual deterioration of a substance when it reacts with substances in the environment, for example, when a metal oxidises in air.

Sacrificial protection refers to the process of using a more reactive metal to corrode preferentially, to protect another metal from corrosion.

Tarnishing is the process by which a thin, discoloured layer forms on the surface of a metal due to reactions with the environment, for example, oxidation.

Electroplating is a process that uses electricity to deposit a thin layer of metal onto the surface of another conductive object.

You may wish to pause a video now and note down these key words and their definitions so that you can refer to them later on during the lesson.

In today's lesson, we have two learning cycles.

The first learning cycle is about understanding corrosion, and the second learning cycle is about preventing corrosion.

So let's get started with our first learning cycle on understanding corrosion.

Corrosion is a gradual deterioration of a substance when it reacts with substances in the environment, for example when a metal oxidises in air.

And you can see in this image, a corroded metal hinge, which was probably silver looking when it was new and now looks a rust colour.

Corrosion weakens the structure leading to failures and economic losses.

Now we might have to be careful when we think about corrosion, not to confuse it with corrosive chemicals as we has symbol, which hopefully you can recognise from previous learning and also some everyday households, chemicals.

The hazard symbol is warning about corrosive substances.

These are substances that can cause damage to living tissues and materials through direct contact.

Corrosive chemicals pose an immediate threat upon contact, whereas corrosion refers to the gradual degradation of materials like metals.

They are not the same thing and should not be confused.

While corrosive chemicals can accelerate this process, they do not refer to the same thing.

So that's something that we need to watch out for, the difference between corrosion and corrosive chemicals.

Quick check for understanding: true or false? Corrosion and the action of corrosive chemicals both refer to the same process.

Well done if you chose false.

Now let's have a look at the reason why is it A, because corrosion is a process that occurs only in the presence of oxygen or B, corrosion refers to the gradual degradation of metals, while corrosive chemicals can cause immediate damage upon contact? Well done if you chose B.

That is the correct answer.

Good work.

So there's different types of corrosion.

We're gonna look at them now.

Galvanic corrosion occurs when two different metals are in contact with a corrosive environment.

And we have an example here that you can see in the image.

And what actually happens is one metal, which is more reactive, becomes an anode and that corrodes and the other metal, which is less reactive, becomes the cathode and that is protected.

So the example that we can actually see is steel and copper.

If they're in contact in a salty environment, for example, on a ship, the steel will corrode faster.

And this is a process that is actually used to help protect some metals.

Pitting corrosion.

Well, this is a localised corrosion that leads to the creation of small holes or pits in the metal.

And if again, if you have a look at the image here, the photograph, you can see clearly those small holes or pits in the metal.

So what actually happens here? Well first of all, it often occurs in metals that have protective oxide layers.

Chloride ions from salt can break down the protective layer and cause the pits.

Stainless steel can develop pits when it's exposed to salt water.

So again, this is a potential issue for ships that are made from stainless steel and obviously are in the sea, which contains salt water.

You'll know from previous learning that metals react with oxygen to form oxides, and this is called oxidation.

When metals react with oxygen, they often form metal oxides.

And you can see in the photograph here a piece of lithium metal.

Now if we were to cut that lithium, inside it would be very shiny, but it is dull on the outside.

Lithium is in Group 1 of the periodic tables and Group 1 metals oxidise very quickly, and we say the surface has tarnished.

Most metals oxidise slowly.

So that dull colour on the outside is a coat of lithium oxide, but we say it has tarnished.

Tarnishing is a surface phenomena where a metal develops a film due to a chemical reaction with substances in the air.

Often, metals like silver, copper, and brass form a thin layer of oxide or sulphate on the surface.

And we say that is tarnished, it's just dull.

It doesn't look as attractive, but it doesn't cause structural damage.

So tarnishing often results in decolorization.

So if you've got that silver picture frame that was once very shiny after maybe several years of just standing on the shelf, it becomes a very dull, unattractive colour.

That is the result of tarnishing.

So another quick check for understanding: true or false? Tarnishing leads to structural damage of metals.

True or false? Well done if you picked false.

Now let's have a look at the reason why.

Is it A, tarnishing doesn't cause damage, but weakens the internal structure of the metal.

Or B, tarnishing can easily be removed without affecting the metal underneath.

Well done if you chose B, we can easily remove the tarnishing.

And if we go back to that example of the silver picture frame, if we get some silver polish, we can just rub it and then clean it over with a cloth, and the tarnish will be removed, and the frame will look a shiny, silver colour again.

Some metals form thin, stable oxide layers preventing further corrosion.

For example, some metals can make the metal appear very dull.

And aluminium is a good example of this.

When aluminium reacts with oxygen, it forms aluminium oxide, AL2O3, and that layer protects it.

So it's a very thin layer over the aluminium metal.

It adheres or sticks strongly to the metal, and it prevents further oxidation and corrosion.

It's effectively transparent compared to some metal oxide layers.

So you can see the aluminium foil in the image.

It is covered or protected in the layer of aluminium oxide.

So it doesn't appear to be quite as reactive as the position in the reactivity series of metals would suggest.

Metals that form protective layers include: aluminium, that goes to aluminium oxide, and it's one of the reasons why it's used for food containers.

Chromium, which goes to chromium oxide, which is essentially used in stainless steel.

And titanium, which goes to titanium oxide, is used in the aerospace industry.

And here is an image of some stainless steel which is used for cutlery, a knife and a fork.

Iron goes to iron oxide, which is rust.

And this is slightly different because it's porous and flakes off leading to more corrosion.

So another quick check for understanding.

Which metals are known to oxidise quickly causing surface tarnish? Is it A, iron and steel? B, Group 1 metals like sodium and potassium? C, stainless steel? Or D, aluminium and titanium? Well done if you chose B, Group 1 metals like sodium and potassium cause surface tarnishing.

Rusting is a specific type of corrosion that only affects iron and steel, and it's really important that we refer to rusting only for iron and steel because in everyday language, it's often used to refer to any type of corrosion which is incorrect.

And you can see a rusty iron chain here with that familiar brown colour.

And you can also see it kind of looks quite flaky and powdery.

That's what rust looks like.

For rusting to occur, we need to have special conditions.

We must have iron, we must have oxygen in the air, and we must have water present.

If we have iron and oxygen or iron and water without that other part, then rusting will not occur.

So true or false? Another quick check for understanding.

Corrosion only occurs when metals are exposed to water.

True or false? That is false.

So well done if you chose that.

Now let's have a look at the reason why is it A, corrosion can occur due to exposure to various environmental factors including oxygen and acidic conditions or B, corrosion occurs because metals react with moisture to form rust? Well done if you picked A.

That is the correct answer.

Alright, I have another question for you here.

Which of the following conditions are necessary for rusting to occur? Oxygen and water? B, any metal, oxygen and water? C, any metal and salt water solution? D iron, oxygen and water? So this is rusting.

Well done if you chose D.

Rusting is that special case, and it must be iron, oxygen, and water.

So overall, rusting, it's quite a complicated reaction.

So our overall equation is iron plus oxygen plus water gives what we call hydrated iron 3 oxide.

And how it actually happens, well we have our iron and then we get our water onto the iron and basically it's a bit like an electrochemical reaction.

We have a cathode, and we have an anode.

And the anode, we can think of is breaking away some irons.

FE2 plus irons are breaking away from the iron and that goes into the water.

So it goes into the solution.

So our oxidation that occurs at the anode is iron goes to iron two plus plus two electrons.

Those electrons are then transferred or attracted towards the cathode and this is where it meets oxygen.

So at the cathode, we have oxygen coming in from the air and a reduction reaction takes place.

So oxygen plus two waters plus four electrons gives 4OH minus irons.

So we have our cathode reaction.

And then the result is we have hydrated iron oxide, and we get that familiar rust colour.

And the formula for that is FE2O3, which is our iron oxide plus X H2O, and X H2O is this hydrated parts, it's our water molecules around it, and there will be various amounts of water molecules depending on how the reaction has taken place.

So rusting overall is quite a complex reaction, and it occurs because we get this oxidation and reduction happening on the iron surface.

So there are many ways rusting occurs.

So we're gonna have a look at one way and summarise what we've just had in that previous slide.

So the two main half equations are O2 plus 2H2O plus 4e minus gives 4OH minus.

Then FE goes to FE2 plus plus 2E minus.

So the iron and the hydroxide irons then react to form iron hydroxide.

FE2 plus plus 2OH minus gives FE bracket OH bracket 2.

That is your iron 2 hydroxide.

This further reacts with oxygen to form that hydrated iron oxide, the rust that we talked about in the previous slide.

So that is 4FE OH2 plus O2 gives 2FE203 dot H2O plus 2H2O.

That is the formula for iron oxide when it is rust; it's got this hydrated part.

So those equations are quite complicated, and you will really need to work through them and understand what is happening at each part of the reaction.

So quick check for understanding.

Which equation represents a reduction reaction in the rusting process? Is it A, B, or C? So remember, reduction is gain of electrons.

Well done if you chose A, oxygen plus two waters plus four electrons goes to 4OH minus.

So that is reduction.

It has gained, the oxygen has gained electrons.

B is oxidation because the iron has lost electrons.

And C is not a reduct reaction.

So very well done if you've got that correct.

So the rusting process speeds up when iron is in the presence of the electrolytes, for example, salty water.

And here we have got a diagram showing the electrolyte and our iron.

So an electrolyte is an aqueous solution containing irons capable of conducting electricity as they are free to move.

The irons produced by the electrolytes carry and transfer electrons, and they facilitate electrochemical reactions in rusting.

So another quick check for understanding.

Why do electrolytes speed up the rusting process? A, they provide a protective coating on the iron surface? B, they facilitate electrochemical reactions by providing irons that can carry and transfer electrons? C, they increase the temperature of the water? D, the absorb moisture from the air? A, B, C, or D? Well done if you chose B.

So this brings us up to our first task.

We have three questions for you.

One, describe what is meant by the terms 'corrosion' and 'rusting', giving details on how they are similar and different.

Two, explain how a metal oxide layer can protect the underlying metal from further corrosion.

Provide an example of a metal that does this.

Three, using the half equations provided, explain what happens during rusting.

Use the terms 'oxidation', 'reduction' and 'electrons' in your answer.

And the two half equations are provided.

Stop the video, have a go at the questions and then when you are ready, we'll have a look at the answers together.

Let's have a look at the answer to question one.

Describe what is meant by the terms 'corrosion' and 'rusting', giving details on how they're similar and different.

First of all, we need to state what they are.

Corrosion is the general process of metal degradation, whereas rusting is specifically the corrosion of iron or steel with water and oxygen.

Similarities: both involve the deterioration of metals, both can weaken metal structures.

Differences: corrosion affects various metals; rusting affects only iron and steel.

Rusting specifically needs oxygen and water.

So very well done if you've got that correct.

And when doing a question like this, it's always good to lay out very clearly the similarities and the differences.

Question two, explain how a metal layer can protect the underlying metal from further corrosion and give an example.

So first of all, the protection mechanism.

The metal oxide layer forms a physical barrier and this prevents oxygen and water from reaching the underlying metal.

An example, well you could have chosen aluminium.

Aluminium forms aluminium oxide.

The oxide layer is thin, adheres strongly and prevents further oxidation.

The thin layer of aluminium oxide is also nearly transparent, so it doesn't significantly alter the appearance of the metal.

So really well done if you've got that right.

Question three, using the half equations provided, explain what happens during rusting and we need to remember to use 'oxidation', 'reduction' and 'electrons' in your answer.

So let's have a look at the answer to this.

First of all, FE2, FE2 plus plus two electrons, the iron loses electrons.

This process is called oxidation.

Iron is oxidised to form iron two irons, which are FE2 plus.

Using the other half equation, the oxygen, O2, gains electrons.

This process is called reduction.

Oxygen is reduced to form hydroxide irons, OH minus.

The electrons move from iron to oxygen causing iron to oxidise and oxygen to reduce.

So very well done if you've got that answer.

It's always good to lay out your work carefully and work through a logical answer.

Excellent work.

This brings us to the end of our first learning cycle on understanding corrosion.

So we'll now move on to our second learning cycle about preventing corrosion.

So desiccants are substances that absorb moisture from the air and help to keep sealed environments dry.

And there's typically two kinds of desiccants that you may have come across.

First of all, silica gel.

These are small transparent beads, and you may have noticed these little silica gel packets if you have bought any electronic devices such as a laptop.

They often are used in packaging for electronics and also food products, and they absorb the moisture from the air.

Another type that's used a lot is calcium chloride, and this looks more like white powder or granules of powder, and they're used in more industrial applications such as dehumidifiers, where we want to have the whole atmosphere dried out.

Desiccants can reduce metal oxidation.

Rust prevention can be achieved by eliminating oxygen or water, and that's what these desiccants do.

They absorb water from the air.

Now we can actually carry out a practical activity experiment to investigate rusting.

So in this example, we have an iron nail placed in four test tubes: A, B, C, and D.

So in the first test tube, we are placing boiled water and oil.

In the second test tube, we are placing some salt water solution.

Then the third one is gonna be left open to the air.

And the fourth one, D, is a sealed container with some calcium chloride in.

So having left these for a period of time, let's have a look at the results.

So that first one A, the boiled water and oil layer.

Well what you'll see is there is no rust because the boiling water removes oxygen, and the oil stops new oxygen from entering.

The salt solution water B, if we look at that nail, it's quite brown.

It does look rusty.

The salt water solution acts as an electrolyte, so it speeds up that rusting reaction.

C, open to the air.

Well, we've got some rust on that.

Not quite as much as B.

There is air and moisture.

Remember moisture contains water and that causes the normal rusting effect.

And in D, it's sealed with calcium chloride.

So this means we can see here, there is no rust on the iron, and the reason is the calcium chloride dries out the air, it removes the moisture from the air.

So still oxygen present, but no moisture.

So this confirms what we learned in learning cycle one, that for corrosion or rusting to occur, there must be iron in the presence of water and oxygen.

So true or false, desiccants reduce metal oxidation by absorbing moisture from the air? Well done if you picked true.

Let's have a look at the reason why.

Well, is it A, desiccants keep the environment dry, preventing water from reactive with the metal? Or B, desiccants coat the metal forming a protective barrier against oxidation? Well done if you chose A, the way desiccants work is they remove the water from the atmosphere.

They're like an air dryer.

That's one of the ways to think of it.

So very well done if you got that correct.

So a simple way of preventing corrosion is to cover up the metal and that basically will provide a physical barrier.

And there's several different ways in which we can do this.

So we can grease something.

We can cover something with a layer of oil or grease.

And if you think about riding a bicycle, the bicycle chain needs to be nicely oiled and that stops it from going rusty over time, stops it from corroding.

We also paint things.

Again, railings outside are often painted.

The paint provides a physical barrier and protects the metal underneath from corrosion.

Or we can put a plastic coating over something.

So for example, this metal test tube holder has been painted to reduce corrosion and it, yeah, it looks a bit better as well.

The coatings can be damaged easily.

So do require regular maintenance.

So for example, if a painting or plastic coating gets a massive scratch into it, that place where the scratch is will get oxygen and water there, and that will part of the metal will start to corrode.

Sacrificial protection is a method where a more reactive metal is used to protect a less reactive metal.

And this is because the more reactive metal corrodes, or we can say it is sacrificed instead of the protective metal.

And this is something that is often used on ships, and you can see an image here of the underneath of a ship where you have got a rudder.

So what actually happens is zinc blocks are placed onto the ship's hull, and the zinc blocks act as an anode.

They will corrode instead of the steel hull.

So basically it will protect the actual hull of the ship, the steel, the zinc will corrode away, but eventually it will need to be replaced.

And if the zinc is not replaced, then the steel will start to corrode, and we don't want that to happen because that will damage the structure of the ship.

Galvanising is a special type of sacrificial protection and what happens here is you coat iron or steel with a layer of zinc.

It's usually done by hot dipping.

The zinc corrodes and prevents rust.

So you can see a galvanised steel handrail here, and if you kind of just look at it, what actually happens is a zinc starts to corrode away, but the steel handrail is protected.

So structurally it's still very strong.

Okay, so let's check our understanding.

True or false: sacrificial protection involves using a less reactive metal to protect a more reactive metal? True or false? Well done if you picked false.

Now let's have a look at the reason why.

Sacrificial protection involves coating the metal with oil to protect corrosion.

So is it A or is it B, sacrificial protection involves a more reactive metal corroding instead? Well done if you chose B.

Well done if you got that correct.

Electroplating is the process of coating a metal object with a thin layer of another metal using an electric current.

And it's used to enhance appearances and improve corrosion resistance.

So there are some basic principles to electroplating.

First of all, the metal to be plated is the cathode.

And you can see this in the diagram.

It's the grey metal that's labelled Me.

The metal salt solution containing irons of the plating metal is the electrolyte.

So in the diagram, we are going to cover this metal in copper.

So our electrolyte must have CU2 plus irons in it.

The electric current causes the metal to deposit on the object.

So when this cell is switched on, the copper 2 irons are attracted towards the cathode and are deposited on the metal.

So before electroplating can occur, the object must be cleaned, and it's often submerged in the metal salt solution as part of the cleaning process.

Why do we do electroplating? Well, there are quite a lot of benefits.

First of all, the end product looks shiny and has a decorative finish.

And in the image that you can see here, we have some silver plated copper lions at a V&A in London that look a lot nicer.

Sometimes jewellery is electroplated.

So to make a cheaper metal look more expensive and attractive such as gold, we electroplate a cheap metal with gold, for example.

But it also protects the base metal from corrosion, which is another really important reason for using electroplating.

The downside of electroplating is it's expensive because it requires specialist equipment and uses electricity in the process.

So let's have a quick check for understanding.

In the process of electroplating, what role does the metal objects to be plated play? Is it the anode, the cathode, the electrolyte, or the solvent? Well done if you chose B, it is the cathode.

Regular checks and maintenance ensure long-term protection.

So you can see here, we've got some divers in the image checking the ship's hull, which is really important to make sure that no corrosion is taking place.

So what the divers will do to start off with, they will carry out some inspections to see what condition the hull is actually in.

After that, there may be some cleaning and some touch-ups.

This prolongs the lifespan of metals, reduces the cost of replacements, repairs and increases safety and reliability.

So it's actually important that those regular checks and maintenance are in place.

Now, some alloys are corrosion resistant.

So stainless steel is an example of an alloy, and you can see an image of some stainless steel cans, so made of iron, nickel and chromium as well as other elements.

The metals like chromium form a metal oxide protective layer making stainless steel very corrosive resistant.

Quick check for understanding: true or false? Alloys are always more corrosion resistant than pure metals.

False.

Let's have a look at the reasons.

Is it A, all alloys are less resistant to corrosion than pure metals? Or B, only certain alloys like stainless steel are more corrosion resistant due to their specific composition? So well done if you chose B.

That is the correct answer.

So that brings us to task B.

We have a table here.

What we'd like you to do is to fill in the table with the information about different methods used to prevent corrosion.

For each method, describe how it works, provide an example and list its advantages and disadvantages.

So pause the video, have a go at the question and then we'll look at the answers together.

Here we have our completed table, and I'm just gonna talk through each of the different methods.

So desiccants, well these absorb moisture from the air, reducing humidity around the metal to prevent rusting.

An example is calcium chloride or silica gel in sealed containers.

The advantage is they're simple and effective in enclosed spaces.

The disadvantages, they're limited to enclosed or sealed environments.

Physical barrier.

Well this physically separates the metal from environmental factors.

Examples include painting, plastic coating and greasing.

The advantages: they're inexpensive, easy to apply and provides immediate protection.

The disadvantages: they can be damaged easily and require regular maintenance.

Alloying.

This is when we have mixing metals with other elements to improve properties.

An example is stainless steel, which of course is iron, nickel, and chromium.

Advantages: we get high resistance to rust, and it's long lasting.

The disadvantage is it can be more expensive than the pure metals.

Sacrificial protection: this is where the more reactive metal corrodes in place of the protective metal.

An example is zinc anodizes the steel on ship hulls.

And advantages: provides a long-term protection.

The disadvantages: requires periodic replacement of the sacrificial metal.

Galvanising: coating iron or steel with a layer of zinc to protect it against rust.

The example is hot dip, galvanised steel handrail.

Advantages: it provides durable protection even if scratched.

The disadvantage is the zinc layer can eventually wear away and need recoating.

And finally, electroplating.

This is when we coat metal objects with a thin layer of another metal using electric current.

An example is silver plated jewellery.

The advantages, it enhances appearances and is corrosion resistance.

The disadvantages, it can be expensive and requires specialist equipment.

So very, very well done if you manage to complete that table and get everything correct because there is quite a lot of different methods that we have covered during this lesson.

This brings us to the end of our lesson today.

We'll just have a look at the key learning points.

Corrosion is a general process of degradation of metals through various causes.

Rusting is aspecific type of corrosion affecting iron in the presence of air and moisture.

A metal oxide layer can protect the underlying metal if the oxide forms a physical barrier.

Some coatings are reactive and contain a more reactive metal to provide sacrificial protection.

For example, zinc to galvanise.

Electroplating can be used to improve the appearance and/or the resistance to corrosion of metal objects.

I hope that you have enjoyed today's lesson.

I look forward to learning with you again very soon.