Lesson video

Hello, my name's Mrs. Navin.

And today, we're going to be talking about the mole as part of our unit on calculations involving masses.

Now, you may have come across this term before in different ways.

For instance, the animal that's known as the mole or a mole that refers to a growth on one's skin.

But the mole that we talk about in today's lesson will not only help us to answer that big question of what are substances made of, but it also helps chemists going forward to talk about particles that are reacting in a more accurate and a manageable manner.

So by the end of today's lesson, you should hopefully feel more comfortable being able to describe what a mole is, and it's linked to the Avogadro's constant.

Now throughout today's lesson, I'll be referring to some key terms, and these include mole, relative formula mass, and Avogadro's constant.

Now the definitions for these key terms are given in sentence form on the next slide, and you may wish to pause the video here, so that you may jot them down so you can refer back to them later in today's lesson or later on in your learning.

So today's lesson is broken into two parts.

The first we'll look at how we can define as chemists, what a mole is, and then we'll move on to look at how it's used with Avogadro's constant.

So let's get started by looking at what we mean by a mole.

Now in our everyday lives, we tend to use words that indicate a defined number.

For instance, if I said I had a pair of socks, you'd know that I'm talking about I have two socks.

If I said I had a dozen of something, you'd know immediately that I'm talking about having 12 of those things.

Going back further in history, there are other terms that were used to indicate a defined number.

For instance, a score.

This is very famous in Abraham Lincoln's Gettysburg Address, four score and 20 years ago.

So he's talking about four times 20 is a score.

You might have come across this idea of a gross, which actually refers to 144 of something, dozens, 12 times 12, 144.

You may have even come across this idea of a ream of paper.

Now a ream simply refers to having 500 pieces of something.

So a ream of paper means you get 500 pieces of paper in that pack.

So we have lots of examples of where a word has been used to indicate a defined number.

Now through their practical work, chemists are constantly measuring substances and that might be the substances volume, its mass or even its concentrations.

Now these measurements help chemists keep track of particles, both those that are reacting and those particles that are produced.

And the mathematical relationships then between the particles in a reaction tends to be shown using a symbol equation like we have here.

So when chemists come to discuss the reactions that they have displayed and represented using these balanced symbol equations, they tend to refer to the particles that are involved.

So if we were to read this balanced symbol equation out, we might read it like this.

Two molecules of hydrogen gas react with one molecule of oxygen gas to produce two molecules of liquid water.

So we're telling a bit of a story with these balanced simply equations, but focusing in on the particles involved in reacting and those that are being produced.

Now the problem we have with this is that particles are so tiny that more often than not, more than simply one particle of any substance is reacting at any time.

So we need to come up with a more accurate way of describing what's going on with this reaction while also keeping track of those mathematical relationships.

Enter the mole.

Now the mole is simply a term that chemists use to represent a defined number, but this number tends to be very large.

In fact, it's this number which is 602 septillion.

Not two, not 12, 602 septillion.

That's a lot of things.

Now if we were to look at that in standard form, it's a little bit easier to get our head wrapped around 6.

02 times 10 to the 23rd, and we refer to a mole with the units mol.

But let's go back to this number.

It's ridiculously large.

So let's try and see if we can come up with some analogies to help us really grasp the enormity of how many things we are representing if we ever refer to a mole.

Let's have a look at this.

If you were to count one apple every second, it would take 11,900 billion years to count a mole of apples.

6.

02 times 10 to the 23rd.

In fact, that's over 2.

6 million years older than our planet.

That's how long it would take to count one mole of an apple, if you counted one apple per second.

Let's look at another example.

If we had one mole that's 6.

02 times 10 to the 23 ring donuts and you put them next to each other all on top of each other, you would cover the surface of the earth to a depth of about five miles.

That's only about half a mile shorter than Mount Everest.

That's about 700 metres shorter than Mount Everest.

If you wanna put it in terms of what we have here in Britain, it would be equivalent to climbing and descending Ben Nevis, the highest mountain on the island three times.

That's how deep a mole of ring donuts would cover the earth in.

That's a lot of donuts.

So where did this value come from? It seems a bit ridiculous when you think about it, but there was actually a lot of work that went into coming up with this number.

In 1811, the Italian scientist, Amadeo Avogadro, suggest that there was a link between the volume of a gas and the number of particles that might be found in it.

Now a lot was done with some of his research.

A lot of people scoffed at him, didn't believe him, but over time, there was a lot of research and a lot of discussion.

And finally then nearly 150 years, 160 years, sorry, years later, in 1971, there was an international group of scientists that came together and agreed that one mole, that 6.

02 times 10 to the 23rd of carbon 12 atoms, exactly equaled 12 grammes of carbon.

And that this value then 6.

02 times 10 to the 23rd would be referred to as Avogadro's number.

Let's look at that definition that was agreed in 1971, a little bit more closely.

They essentially agreed that one mole of carbon 12 atoms, so that's 6.

02 times 10 to the 23 of carbon, 12 atoms, equals 12 grammes of carbon.

Now some of you eagle eyed out there will have noticed there's a link between the relative atomic mass that we find on the periodic table and the mass of carbon atoms that made up one mole.

So really what they were agreeing was that the mass of one mole is equal to the relative atomic mass of that substance measured in grammes.

Okay.

Now using this particular definition meant that the number of particles in a mole was based on the average relative masses of that substance, and it had to be determined experimentally.

Now despite the mass of one mole of particles, needing to be determined experimentally, it did have some very good uses.

So I'm gonna take a step back here and just look at how we can use relative masses going forward.

So if we assume that one tennis ball has a mass of 57.

6 grammes.

If I had 12 tennis balls, so if I had a dozen tennis balls, I could find that the mass was 691.

2 grammes, simply by multiplying the mass of one tennis ball by 12.

Same thing if I had a basketball.

If my basketball was 565 grammes, if I had a dozen basketballs, all I would need to do is take the mass of one and multiply it by 12, representing a dozen and I would get a mass of a dozen basketballs was 6,780 grammes.

So despite there being the same number of balls, tennis balls or basketballs, the mass of a dozen of each is different because each of these sports balls have a different mass.

So similarly, because the atoms of different elements have different relative masses, the mass of one mole of different elements also have different masses.

So if we look at carbon, we said the relative atomic mass was 12, which means a mass of one mole is equal to that relative atomic mass in grammes.

So the mass of one mole of carbon is worth 12 grammes.

If we were to look at a different element, this time we looked at magnesium.

Its relative atomic mass was 24, so a mass of one mole would be 24 grammes.

And then if we looked at argon, again, its relative atomic mass is 40, but the mass of one mole then would be 40 grammes, because it has a different element, it has a different relative mass.

So the mass of one mole of them is going to be different as well.

Now one thing to remember about relative atomic masses is that they are a comparison scale and it doesn't have any units.

But what that means then, is the ratio or the comparison, that relative comparison between the relative atomic masses of two different elements.

It's gonna remain the same if we were then to compare the mass of a mole of these elements.

So if we have here, we have helium and carbon atoms. Now if we look at just one atom, the relative atomic mass for helium is four, and the relative atomic mass for carbon is 12.

The difference between those is that, carbon is approximately three times heavier than an atom of helium.

Now if we then look at a mass of one mole, so this time, we're talking about 6.

02 times 10 to the 23 atoms of helium.

That mass of helium is gonna be four grammes and that for carbon will be 12 grammes.

The ratio, that relative comparison between the two remains the same, whether or not we're talking about one atom or 6.

02 times 10 to the 23 atoms. Let's stop here for a quick check.

Now I have a sample of magnesium and of carbon powder, and I've given the relative atomic masses here.

Magnesium is 24 and carbon is 12.

What I'd like you to do, is to decide which pairs of samples contain the same number of atoms and you may wish to select more than one answer.

Now this may require a little bit of discussion with those nearest you, so I might recommend you pause the video here and come back when you're ready to check your answers.

Well done if you manage to choose B and D.

When we compare the relative atomic masses for magnesium in carbon, we can see that magnesium is twice as large as carbon.

So looking at my options from A to D, I'm looking for values in grammes that have that same ratio where the carbon is half as much as the magnesium.

And when I look at that, I can see that only B and D match that ratio.

So that's why those two are the correct answers.

Very well done, if you've got at least one of those correct answers and extremely well done if you've got both of them.

Great start, guys.

Now the thing about science is that it's dynamic.

It's constantly changing as we gather more information, as more research has done, and more discussions take place.

And in 2018, a community of scientists that represented over 60 countries came together and agreed to change the official definition of a mole.

And what they agreed was that one mole is now going to be equal to 602 septillion particles of something.

And again, in standard four, it remains that one mole is equal to 6.

02 times 10 to the 23 particles.

So rather than being defined by the average relative masses that they were before, a mole now represents a constant number of something.

And this number is known as Avogadro's constant, and it has a unit of moles to the minus one.

And the way we read this is particles per mole.

So the mol to the minus one means per mole.

So rather than being defined on these relative masses that needed to be experimentally determined, a mole now represents a specific number of particles that can count it up if we wanted to.

Now like all things with a new definition, parts of it need defining a little bit more, and that's exactly what's happened here, is if a mole equals 6.

02 times 10 to the 23 particles.

What is a particle then? Well, a particle depends on the structure of the element or the compound.

And remember, if we talk about structures, we can talk about a metallic structure, a simple molecular structure or a giant ionic lattice structure.

Now metallic structures tend to be atoms as they're particles.

A simple molecular structure tends to be a molecule then and a giant ionic lattice.

So things that are normally made of metals and non-metals tend to be known as a formula unit then.

Now, because the particles can be an atom, a molecule, an ion, or any other particle, it's really important that we are clarifying what's being counted when we use this updated term of a mole.

Okay.

So how do you determine particle type? Sometimes, it's difficult to remember what type of structure you're actually talking about.

So what I wanna do is go right back to basics and look how we can build up from our understanding of what our substances to determine what kind of particle type we're talking about.

So if we think about an element, we're talking about the most basic building blocks and that we have in chemistry, and these are ones that you can find on the periodic table.

Now elements can be classified very simply as being either metallic or non-metallic.

Now if they're a metal element, they tend to have then a giant metallic structure.

If they are non-metal elements, well, they can be quite a few different structures.

They can be monoatomic, so made of one atom, they can be giant covalent structures, things like carbon or silicon dioxide.

More often than not, they're gonna be simple molecular.

So these substances that are made of very few atoms or very few elements.

Based on these structure types, then we can determine the particle type.

If it's a giant metallic, a monoatomic or a giant covalent structure, the particle type you're talking about tends to be atoms that you are counting when you're referring to a mole.

If you have a simple molecular structure, so something like water, carbon dioxide, then you're talking about a particle type that is a molecule.

Now the other type of substance that we can get is a compound.

Now remember a compound is a substance that's made of two or more elements that have been chemically bonded together.

Now again, you can have then either two non-metals that have been bonded together or a metal and a non-metal.

Now based on which elements have actually been bonded together in our compound, we can get different structure types.

If it's a non-metal and a non-metal, you could have either simple molecular substances or you could get a giant covalent structure.

Okay.

If you have a metal in a non-metal, which tends to be the one you get the most, you tend to get a giant ionic lattice.

Now again, like before, based on these structure types, you can determine the particle type that you're going to have.

And when you have a giant covalent structure or a giant ionic lattice, what you're talking about is this formula unit.

Now this particular diagram here has a lot of keywords and key ideas, and I would actually recommend, if you have some time to pause the video and maybe copy this out or print out this slide because this screams to me, a perfect reference slide.

And what you can do if you wanted to, is actually add little bits of information around them, maybe draw some pictures or give some examples of formula to represent the different structures, the elements, and the particle types.

Draw some particle diagrams. This is the perfect revision material.

So if you want to take a moment and use this to prepare for the future, it's a great resource.

But hopefully, this will help you going forward now when we're counting up our moles, what particle type we're talking about.

So let's put it into action and we're gonna stop here for a few quick checks.

Which statement describes the number of particles of one mole of magnesium? Well done if you chose A.

If you remember, magnesium is a metal and it's structure type then is a giant metallic structure and we refer to those as atoms. So one mole is 6.

02 times 10 to the 23 magnesium atoms when we're talking about a mole of magnesium.

Well done, guys.

Good job.

Let's try another one.

Which statement or statements describes the number of particles in one mole of oxygen? Well done if you chose B.

Now if you remember, oxygen likes to go around in pairs, it's known as a diatomic molecule and one mole is equal to 6.

02 times 10 to the 23.

So B is correct.

We're talking about the correct number of particles in one mole, which is molecules.

There was another possible answer because if we were to count up all the oxygen atoms in one mole of oxygen, we'd have twice as many oxygen atoms as we had of oxygen molecules and 6.

02 times 10 to 23 times two, is equal to 12.

04.

So well done if you got B and incredibly well done if you managed to choose D as well, guys.

Good job.

Okay, guys, let's try one more.

Which statement describes the number of particles that we would be finding in one mole of sodium chloride? Well done if you chose C.

Now we still have one mole and that's 6.

02 times 10 to 23.

Now because we're talking about an ionic substance, we've got a metal and a non-metal.

ionic substances are formula units.

So we're now down to looking at C and D as our options.

Ions, not atoms, are what make up sodium chloride and that's what makes C the best of our four options here.

So incredibly well done if you manage to choose C as your answer.

Really, really, well done, guys.

It is not an easy idea to get your head wrapped around and I am so proud of you, persevering through this great job.

Okay, guys, time for our first task of today's lesson.

What I'd like you to do is to match each term to the correct description.

You may wish to discuss this with some people nearest you.

So pause the video and come back when you're ready to check your answers.

Okay.

Let's see how you got on.

So atoms are the smallest part of an element that can take part in reactions and we've given some examples there.

So if you weren't sure about the descriptions, hopefully some of the examples helped you.

We could see that they are just the element symbols there.

Molecules then is a particle consisting of two or more atoms that have been joined together.

So again, we've got some formula there to help you guide you to that correct answer if you weren't sure.

And then a formula unit is gonna be the lowest whole number ratio of a compound with an ionic bond.

So the red flag in that description was the fact we were talking about an ionic substance and formula units are how we describe that particular structure of ionic substances.

So very well done if you manage to match all of those terms correctly.

Great start to the task, guys.

For the next part of this task, what I'd like you to do is to write the formula for each of the substances that have been listed below.

Now you may need to just refresh your memory on how to write some of these formulas or work with the people nearest you.

So I recommend you pause the video and come back when you're ready to check your answers.

Okay, let's see how you got on.

So sodium chloride is NaCl, making sure that you have your capital and lowercase letters so it's clear what these element symbols are that you're using.

Oxygen is O2, and remember that two should be subscript.

Magnesium is Mg, carbon dioxide is CO2, helium is He, carbon should be just the letter C, and silicon dioxide is SiO2.

Well done if you manage to get those correct.

For the next part then what I'd like you to do is to classify those same substances into particle types and you may wish to refer to the formulas that you did in the previous part of this task in order to help you identify which particle type.

Now remember particle types we're talking about atoms, molecules or formula units.

So you may wish to pause the video and come back when you're ready to check your answers.

Okay, let's see how you got on.

So sodium chloride is a metal and a non-metal.

So it's an ionic substance and therefore it is a formula unit.

Oxygen had the formula O2 and therefore it's a molecule, 'cause it's two atoms that are joined together.

Magnesium as a metal, is referred to as an atom by particle type.

Carbon dioxide is a simple molecular substance and therefore, a molecule.

Helium is just the element by itself, it's monoatomic and therefore, the particle unit is an atom.

Carbon is also an atom for the same reason as helium.

And silicon dioxide is a formula unit.

This is kind of a tricky one, but silicon dioxide in terms of its structure type is a giant covalent structure.

It looks a little bit like diamond does.

So that was a tricky one, but very, very well done if you managed to get that correct.

Great job, guys.

Okay.

For the last part of this task, I'm gonna ask you to do a few different things and you're going to need probably two different coloured pens, markers or work on using circles or squares.

But essentially what you're gonna do is you're gonna look at this table of masses.

We've got nine samples of different elements and those elements are helium, carbon, and magnesium.

And I've provided you with the relative atomic mass.

You could have got these off the periodic table as well, if you have that handy.

What I'd like you to do is four things.

So the first thing I want you to do, is draw a circle around each sample that contains the same number of atoms as 12 grammes of carbon.

Next, I'd like you to draw square or use a different colour to identify each sample that contains the same number of atoms as 24 grammes of carbon.

With those remaining samples that don't have a circle or a square or a colour around it.

Do they contain the same number of atoms? And I'd like you to explain how you worked out your answers to A, B, and C.

Now this is gonna take a little bit of time and I would highly recommend you work with somebody else on this just to check your work.

It's always helpful to discuss the numbers with somebody else and see if you agree.

So definitely pause the video here and come back when you're ready to check your answers.

Okay, let's see how you got on.

So the first thing you were asked to do was to draw a circle around or highlight with one colour all the samples that contain the same number of atoms of 12 grammes of carbon.

And the way I'm gonna decide which one's to circle is to remember that the mass in grammes is equal to the relative atomic mass of the substance.

So definitely 12 grammes of carbon, but then four grammes of helium and 24 grammes of magnesium should have been the ones that you circled or highlighted here.

So very well done if you managed to get those correct.

So the next thing you were asked to do was to draw square around or use a different colour to identify the samples that contained the same number of atoms as 24 grammes of carbon.

And to decide which ones to do there, I noticed that 24 grammes is equal to twice the relative atomic mass of carbon.

So the masses I'm looking for in the samples have to equal two times the relative atomic mass of those substances.

And when I look at that, then I should have a square or highlight of eight grammes of helium, 48 grammes of magnesium, and 24 grammes of carbon.

So the last thing you were asked to do then was to compare the remaining samples so those that don't have a circle or a square or a highlight around it.

And what can we say then about the number of atoms in those samples? Well, the first thing I can see then, is that the six grammes of carbon and the two grammes of helium are probably gonna contain the same number of atoms because both of those samples are equal to half of the relative atomic mass for those elements.

Magnesium on the other hand, does not contain the same number of atoms as those other samples.

And that's because eight grammes of magnesium is actually knot 0.

3 times the relative atomic mass of magnesium.

And so therefore, it doesn't have the same number of particles.

So that was a really tricky kind of mental mass thing.

But if you have a really systematic approach and just remember that ratio we talked about, the relative atomic mass for one atom versus the mass in grammes of multiple particles, then it becomes a little bit easier.

So remember the main thing I was a asking you to do as well, is to think about how are you going to choose which samples have the same number of atoms? That was your because clause, why have you chosen to match up these samples together? So that's what I had started each of my feedback slides with.

So hopefully that made a little bit of sense to you and why we chose those particular samples.

But really, really a fantastic start to today's lesson, guys, really proud of you.

Great job.

Now that we're feeling a little more confident talking about what a mole is, let's look at how we can use it with Avogadro's constant.

Now sometimes when we're talking about moles or Avogadro's constant, it's useful to think back to those terms that we use to represent specific numbers.

Now, a lot of the products that we come across in everyday life are provided in packages of specific numbers, and we can actually refer to those products as fractions or multiples of the package that they come in.

And it still indicates a specific number.

Now a good example of this is eggs.

Eggs tend to come in packets of a dozen and I could refer to a fraction of that package like half a dozen eggs, in which case, I take a dozen, multiply it by that fraction, and I know that I'm talking about six eggs.

Similarly, I could talk about multiple packages.

Let's say I have two dozen eggs, I take 12 representing a dozen and multiply it by two, and I'm talking about 24 eggs.

So I can talk about individual numbers of things by referring to a fraction or a multiple of the package that they come in.

So chemists are able to do something really similar but using a mole to do it.

So we're using this understanding that Avogadro's constant is 6.

02 times 10 to the 23 particles per mole, and we're using that to package our particles.

So one mole, one package, will tell me I have 6.

02 times 10 to the 23 particles, so I could talk about a mole of sugar.

What I'm really talking about is 6.

02 times 10 to the 23 particles of sugar or I could talk about a mole of salt.

And what I'm talking about is a package of 6.

02 times 10 to the 23 particles of salt.

So if I wanna know how many packages, multiples or fractions I have of a mole, I'm gonna take the mole in my sample, fraction of multiples, is gonna equal then the number of particles that I have present in my sample, divided by Avogadro's constant.

So let's look at how I can use that relationship.

If a recycling centre has collected 2.

3 million pen lids and I now want to know what that is in moles to two significant figures, I'm going to use that relationship, moles equals particles divided by 6.

02 times 10 to the 23.

So because we're talking about moles, I need to clarify what particles I'm talking about.

And in this case, it's the pen lids that I am counting up.

So I'm going to take that value of 2.

3 million, but change it into standard form for ease of using in my calculator and divide it by Avogadro's constant.

And when I do that, I get this value, but I want this to two significant figures.

So my final answer then is going to be 3.

8 times 10 to the minus 18 moles of pen lids.

So I've gone through an example here of how you can use that mathematical relationship.

What I'd like you to do now then, is to have a go yourself at calculating the number of moles of sand that are found on our planet to two significant figures.

So grab yourself a calculator, pause the video, and come back when you're ready to check your answer.

Okay.

Let's see how you got on.

So we're using that same relationship of moles equals particles divided by Avogadro's constant.

The particles here are the grains of sand.

So you're gonna take that value of 7.

5 times 10 to the 18 and divide it by 6.

02 times 10 to the 23.

And when you round that answer to two significant figures, you should have 1.

2 times 10 to the minus five moles of sand grains.

Very well done if you managed to get that correct.

Great job, guys.

Now why do chemists use moles to refer to the number of particles? Well, we know that the number of particles in a sample is gonna be equal to the number of moles times Avogadro's constant.

Okay.

That's 6.

02 times 10 to the 23.

It's a lot simpler to talk about the number of moles and compare those moles in a reaction than it is to refer to this large number of 6.

02 times 10 to the 23 or a fraction or multiple thereof.

So let's look at an example.

If I have one mole of salt, okay.

That's one times 6.

02 times 10 to the 23, means I have 6.

02 times 10 to the 23 formula units of salt in that sample.

If I have one zero, sorry, 0.

1 moles of salt, again, I'm gonna use that calculation, 0.

1 times Avogadro's constant and I get a value then that's 6.

02 times 10 to the 22 formula units of salt.

Now the only difference between these two numbers of particles that are in my samples are those exponents, 10 to the 23 versus 10 to the 22.

But if I refer to the moles, one mole versus knot 0.

1 mole, it's a lot simpler and it's easier to signify that one is significantly less than the other in terms of the number particles in that sample.

Having said that, whilst is really a lot easier for us to refer to and compare the number of moles within a sample.

It's a far more accurate representation of the number of particles if we continue to use that mathematical relationship.

It allows us to calculate the specific number of particles in a sample.

So if I had two moles of sulphur, I could take two times Avogadro's constant and I would know that there are 12.

04 times 10 to the 23 atoms of sulphur in that sample.

And if I actually had a sample of knot 0.

25, so a quarter of a mole of sulphur, again, multiply that by Avogadro's constant and I would see that there's 1.

51 times 10 to the 23 atoms of sulphur.

So this mathematical relationship is allowing me to accurately and more specifically clarify the number of individually particles that are in this sample.

It's not as easy to talk about, but I can clarify the number of particles that are there.

Let's have a go now then to use that relationship, calculating the number of molecules in 5.

47 moles of water to three significant figures.

So as a reminder, the relationship is particles equals moles times 6.

02 times sent to the 23.

Now the particles here are the molecules we've been told about in our instructions.

Then we're gonna take the number of moles we have of water, multiply it by Avogadro's constant, and we get this value here.

But we wanted it to three significant figures, so our final answer then is going to be 3.

29 times 10 to the 24 molecules of water.

What I'd like you to do now is to use this as a guide to do your own calculations of the number of molecules that are found in knot point, knot two five moles of oxygen.

So grab your calculator, pause the video, and come back when you're ready to check your answer.

Okay.

Let's see how you got on.

So we're using the same relationship as before.

Our molecules are our particles again.

We have our moles times Avogadro's constant, gives us this value here, but to three significant figures then our final answer should be 1.

51 times 10 to the 22 molecules or particles of oxygen.

So very well done, if you manage to get that correct.

Remember, it is really important that you are showing you're working out in case you have gone wrong so that we can identify where and you can improve going forward.

But I would always recommend that you're not rounding any of your answers until you get to that final answer.

Great job, guys.

Keep it up.

Now one other thing we can remember is that if the mass of an individual object is known as well as the number of objects present, then we can calculate a collective mass.

So for instance, if I had a sheet of paper and I know that a sheet of paper that's A4 in size, is five grammes.

How might I go about calculating the mass of a ream of paper? Well, the first thing I need to do is remind myself of how large or what a ream represents.

And a ream equals 500 of something.

So what I'm gonna do then is multiply then the mass of one sheet by the number present, 500.

So I'll take five times 500, gives me 2,500, which means then the mass of one ream of paper is going to be 2,500 grammes.

Now chemists then can use a very similar process then to predict the mass of a sample.

And let's look at how that's done.

So if I know that one atom iron has a mass of 9.

28 times 10 to the minus 23 grammes, and I wanna know then what the mass of 1.

5 moles of irons will be to three significant figures.

Firstly, I need to know the number of particles that are 1.

5 moles.

So I'm gonna take that times Avogadro's constant and that tells me how many particles I have in that package of iron.

And then I'm gonna calculate the number of particles by the mass of one particle of iron, and that gives me then a total value of 83.

8 grammes.

So I've given you, guys, an example layout here.

What I'd like you to do now then is to calculate what the mass of knot 0.

750 moles of gold atoms would be to three significant figures.

So grab your calculator, pause the video, and come back when you're ready to check your answer.

Okay, let's see how you got on.

So the first things first, you need to calculate the number of particles that are present in knot 0.

75 moles.

And we get 4.

52 times 10 to the 23, multiplying that by the mass of our gold.

And our final answer to three significant figures then, and we should have a value of 148 grammes is how heavy knot 0.

75 moles of gold atoms would be.

Well done, guys.

You're doing a fantastic job.

Keep it up.

Okay, guys, let's move on to the last task of today's lesson.

For this first part, I'd like you to calculate the number of particles each sample suggested below and to give your answer to three significant figures.

So calculators at the ready, pause the video, come back when you're ready to check your answers.

Okay, let's see how you got on.

Now the main thing here is that you needed to use that equation of moles times Avogadro's constant will tell you the number of particles.

So if we take 3.

5 times 6.

02, times 10 to the 23, and then round that final answer to three significant figures, we should have 2.

11 times 10 to the 24 particles or atoms of iron.

For B then knot 0.

32 moles of sodium chloride, then we should have to three significant figures, 1.

93 times 10 to the 23 particles.

And then for C, we have for knot 0.

845 moles of carbon dioxide then, would equate to 5.

09 times 10 to the 23 particles or molecules of carbon dioxide.

So very well done if you managed to get those correct.

I have put the workings out here in case you went wrong.

So you could just double check that you've put the numbers in the right place and got similar answers on your workings to those three significant figures.

Well done, guys.

For the next part of this task, I'd like you to calculate the number of moles in each of the samples listed below and to give your answer to three significant figures.

So once again, you wanna get your calculators ready.

Don't forget to show your work and then pause your video, come back when you're ready to check your answers.

Okay, let's see who you got on.

So the main thing here then, is we're going to be taking the number of particles and divide by Avogadro's constant to calculate the number of moles in each sample.

So for A to three significant figures, we should have 1.

01 moles of helium.

For B, we should have to three significant figures, 12.

5 moles of CH four, and for C then, we should have to three significant figures, 0.

0880 moles of magnesium oxide.

And again, I've shown you the working out if you wanted to double check your numbers against mine.

But the main thing here is that we're dividing by Avogadro's constant to find the number of moles.

Well done, if you manage to get those correct.

You're on a great roll, guys, keep it up.

Now for the final part of today's last task, I want you to calculate the mass in each sample to three significant figures.

So don't forget to show you're working out, get your calculators at the ready, pause the video, and come back when you're ready to check your answers.

Okay, let's see how you guys got on.

So the main thing here is, it was a two step process.

You firstly needed to find out how many particles there were in that many moles by multiplying the number of moles by Avogadro's constant, and then you needed to multiply the number of particles by the mass of one particle.

So for A, when we do that, this is the number of particles present in my sample.

And the mass then to three significant figures should be 43.

3 grammes of gold atoms. For B then again, these are the number of particles that I have, and when I multiply them by the mass of one particle of magnesium, I get a final answer of 35.

9 grammes of magnesium.

And then for C, I've got magnesium oxide here, this is how many particles there are in my sample.

When I multiply it by the mass of one formula unit then, I get a final answer of 30.

6 grammes of magnesium oxide.

So give yourself one mark for correctly calculating the number of particles in each sample.

And a second mark for calculating the correct mass processing in that sample.

And give yourself a third mark for being able to correctly round your answer to three significant figures.

Incredibly well done, guys.

I'm really proud of you.

Great job.

Okay, we have done a lot of learning today in our lesson.

And let's take a moment now to summarise what we've learned.

So we've learned that chemicals are measured in moles, and moles are currently defined as Avogadro's constant number of particles, which means then that one mole is equal to 6.

02 times 10 to the 23 particles.

We also learned that the number of particles, so that could be atoms, ions, molecules in one mole of a substance does not change.

Okay.

And Avogadro's number then, can be used to calculate the number of particles in a mole or how many moles there are in so many particles of a substance.

We've also learned that the mass of one mole of an element is equal to the relative atopic mass of that element measured in grammes.

So I've had a really good time learning with you today.

I hope you've had a good time learning with me, and I hope to see you again soon.

Bye for now.