AQA Atomic Structure Revision Notes

Hello my name is Chris Harris and I'm from AllergyTutors.com and this video is basically we're just going to go through AQA, Atomic Structure, so this is specifically for AQA and basically we're going to go through the key points just as revision it's a good overview of the topic to make sure that you've covered most of the things that you need to know and just before I start the powerpoints that I'm using here that I've made they are available to purchase, if you just click on the link in the comments box just below, then you should be able to... click on that and you'll be able to buy them there. If you're interested, it'd be great for things that you can print them off, or you can look at them in your own time, you know, use them as revisionals, whatever. So they're quite colorful, so they should be reasonably attractive to use. Okay, right, so let's make a start. Like I say, all of these things, they are tailored to the specification, so to make sure that they are absolutely watertight, well, nearly, so you can look at the specification, basically it just matches that, okay? So we're gonna start with the Atom. So the Atom is, can see it's mainly made up of protons and neutrons they are very small and they contain that are contained in the middle there you also have electrons that's whizzing around in shells so these orbits the middle of the atom as you can see and you need to know these charges as well so protons have a positive charge as you can see and there it is there and neutrons have a zero charge neutron neutral electrons are negatively charged so they have a minus one charge They're relative masses. You need to know as 1, 1, and 1 over 2000 is the relative mass of an electron. You can't put no mass. An electron does have a mass. It's just very, very small. And just to familiarise yourself with the... the elements in particular. So here's lithium for example. The top number tells us the mass number, that's the number of protons and neutrons in the nucleus and the bottom number tells us are just the number of protons. We call this the atomic number or proton number. So yes all atoms remember are neutral because the number of protons equals the number of electrons. Okay right so let's look at some ions and isotopes. So You've got to know the difference between these. We'll start with ions first. So ions have a different number of electrons and protons. So they basically don't have equal amounts like we looked at in an atom. So for example if we look at negative ion, O2 minus has gained two electrons to get a full stable shell of electrons. So you can see here that the protons in this particular one here, that we have eight protons which give a charge of plus eight. Eight neutrons, remember they have no charge. count electrons we've got 10 electrons in O2 minus that means it gives it a minus 2 charge so the total charge on this is minus 2 and that's how we work out them charges okay so this basically allows them to form these stable ionic compounds they gain stability by being attracted to positive ions okay so let's look at a positive ion this is sodium sodium is the opposite it has one electron it's out of shell it's in group one it loses an electron to form a Na plus ion and you can see the same calculation I've done. Protons, 11 protons it's got, 12 neutrons, which have no charge, and electrons, because it's lost an electron, it now only has 10 electrons, so the charge is at minus 10. Now if you do plus 11, 0, and minus 10, you get an overall charge of plus 1. That's why sodium has got a positive charge. Okay, isotopes, right. These are elements with the same number of protons, but a different number of neutrons. they have a different mass effectively. So these are slightly different. So let's look at these three examples here. These are three examples of carbon. Carbon 12, carbon 13, carbon 14. And let's just look at the different numbers. You can see we've got the protons, neutrons and electrons here. Now if we look here, they've got the same number of protons, each all of these isotopes too, but crucially the number of neutrons is different. So carbon 12 has six neutrons, Carbon-13 is 7 and carbon-14 has 8 neutrons. So these are isotopes of each other. Okay, chemically they react the same because they have the same number of electrons, but they have a slightly different mass. Okay, right, we need to know a little bit about the history. This is the fun bit. Okay, back in 1803, John Dalton, he came up with this idea that atoms were all spheres. A very simple idea. And it wasn't until a good while later, nearly 100 years later, that J.J. Thompson had a good go at trying to change the model that Dalton came up with. He discovered the electron. And the atom, he said basically the atom wasn't solid and it was made up of other particles and he named it the plum pudding model. So basically he said that you had the positive pudding bit, which is this bit here, and inside the pudding you had negative electrons, which are the yellow circles there. That's what Thomson came up with. Very shortly after that, Ernest Rutherford caught the atomic bug, I suppose, and he discovered the nucleus. and also discovered that the nucleus was really really really small and Actually the nucleus contained positive charges So he basically said that most of the atoms empty space and because the nucleus was so small And he said we had a cloud of electrons that was surrounding this nucleus So you can see now we're starting to resemble more like the atom that we know today And he proved this using the gold leaf experiment. Basically, what he did is he fired alpha particles at a very thin bit of gold leaf. And most of the alpha particles went through. Now, it tells him that most of it was empty space. So that's where he got his idea from. But some did deflect back. Very, very small number. And so what they must have done is they've hit a small positive nucleus. Because an alpha particle is positively charged. So it's actually hit the nucleus. bounced back some of it was deflected some of it bounced straight back right back at the um at the the the uh the source of the alpha okay so 1913 again look not many years later um neil's ball had a go of it so they really got the atomic bug here and he discovered a problem with rutherford's uh model and basically he said well the electron cloud could collapse because it would obviously fold into the positive charge nucleus um And so he said, well, actually, you couldn't just have a cloud. You had fixed energy levels. And this is where the shell came up with. Now, basically, he proved this because actually when he shone electromagnetic radiation, it was absorbed by the atom and the electrons move between the shells. That's what he noticed. And when they do this, they emit radiation when the electrons move back down to a lower energy level. Now, this could only be explained using a shell model. You couldn't explain this using the cloud model. So this basically kind of cemented Neil's ball idea of the atom. And obviously the model today, we know that actually yes, there are shells, but we now know of the existence of subshells. And basically we can use the ionisation trends to explain this, which we'll look at later. The time-of-flight mass spectrometer. So this is basically mass spectrometry, but using time-of-flight. The first bit, when you add your sample, it's vaporised. So it can travel through the time of flight mass spectrometer. So we turn it into a gas basically. Then we ionise it. Now we can ionise this using a, we call it electrospray ionisation. It basically works when we spray a sample through a high pressure jet. It's like putting your thumb over a hosepipe, it's a really high pressure jet. And basically what they do is they pass a really high voltage through this jet. And this causes the loss of an electron. And what we get is that... gaseous positively charged sample is made and this is really important because we need ionization for the next stage which is acceleration. So you can see there there's the blue particle look they're moving through and so these are accelerated by negatively charged plates on electric fields and basically the particles with a lower mass to charge ratio or mz ratio will accelerate quicker so they'll move through a little bit quicker. And the next bit is the, basically it's the iron drift. This is a bit weird. Now you can see the red one zipped through there and the blue one's a little bit slower. But the particles travel through with a constant speed and kinetic energy. So if I just go back, you can see that both are travelling at the same speed and the blue one, they're travelling at a constant speed, sorry, not the same speed, travelling at a constant speed, but they do have, between them, they have a different speed each. but their speed is the same in terms of the atom. So in other words, the red atom constantly travels through. It doesn't speed up or slow down as it goes through the ion drift. It travels through at a constant speed, but the blue one goes slower. And then the final stage is detection. So once it's drifted through, obviously the blue's going a bit slower than the red, then basically an electrical current is made when the particle hits the plate at the back. And basically ones with a lower MZ and the lighter particles will reach the detector first as they travel the fastest. So we have effectively separated our different parts out and we're detecting them at the other side. Okay, right. Your definitions, right? You need to know these. I can't emphasize this enough. Right, relative atomic mass. You can probably read them there, pretty straightforward. Relative atomic mass is the average mass of an atom of an element when measured on a scale on which the mass of an atom of carbon-12 is exactly 12. Basically, we're measuring... everything relative to carbon-12. The relative molecular mass, so this is very similar if you look the average mass of a molecule when measured on a scale which the mass of an atom is carbon-12 is exactly 12. So this is a molecule instead of an atom and relative isotopic mass is basically the mass of an atom of an isotope with an element measured on a scale in which the mass of an atom of carbon-12 is exactly 12. So there's a lot of carbon-12 being mentioned here you've just got to know these. Ready? Okay. This is a mass spectra. Okay, so here we are. We've got a this one. We look at isotopes So we've got an element here and it's made up of isotopes now the first thing we need to look at really Is the the axes so you can see here that we've got a mass to charge ratio at the bottom And this is basically the mass of the isotope divided by the charge And most do have a plus one charge and so this makes it the same as the isotopic mass Which makes it? relatively straightforward if it had two electrons knocked off which would be quite rare and then the obviously the mass the charge ratio will be half as much um so that's quite that's quite important um because obviously the z bit um stands for charge so if you've got a double charge it's just the mass of the isotope divided by two okay um the bit on the side read this really carefully i mean this is the abundance um it's always shown on the left but this one it can be written as a percentage or a nominal value so it can be just a relative you Abundance this one's percentage abundance So this means that all your isotopes must give 100% if it's a percentage abundance because obviously you can't get bigger than 100% Okay, um so you can see here if we have the 75 and 25 and get 100% from there Okay, so this spectra shows two isotopes of one element So we've got one element going through here made up of two isotopes and we know this because we've only got two peaks So we've got one isotope that has a mass of 35 and one has a mass of 37 This is assuming obviously they have a 1 plus charge. So you can see here that the most abundant, which means the most common isotope, is isotope 35, whatever this is. 37 isn't as common. Right, so, and from all this information we can work out the relative atomic mass, which we're going to look at now, which is pretty useful. Okay, so let's work out the relative atomic mass of these. You need to know this formula. Relative atomic mass is the abundance of... isotope A times by the MZ of A, so that's the mass to charge ratio of A, plus the abundance of B times by the MZ of B. And if you had more than two, you would just literally keep adding up the abundance of A, abundance of B, and abundance of C, et cetera. You just keep adding them up. So basically, you keep adding loads of these brackets, repeating it for each isotope, divided by the total abundance. Now, because this one's percentage abundance, our total abundance is going to be 100, but it might not be percentage, so kind of look out for that. Okay, so let's look at this. Relative atomic mass is going to be 75 times 35, because the abundance of A is 75. The mass of A is 35, mass to charge, sorry, is 35. So we do 75 times 35 plus 25 times 37, so that's them two, divided by 100 equals 35.5. And if you're smart enough, you can look in the periodic table and you can identify that as chlorine. So you've effectively identified your element from your mass spectrometer. So that's pretty nice and straightforward. You can also do it through tables as well. You don't have to give it through spectra. Similar thing, this one's got more isotopes, as you can see. We don't know what the element is of the relative atomic mass, but we're going to try and work it out. So here's that equation again. Look, all we're doing is taking the isotope, multiplying it by the abundance, 20.5. There it is, 70 times by 20.5 plus, and then here's your other isotopes I was talking about. So you just basically add them up. It is a percentage of abundance so we divide it by 100. If it's not then we just add these numbers up and divide by the total of their numbers. Okay so the answer is in this case is 72.6. This means that this element is germanium and you can have a look in the periodic table and check that out. Okay right molecules. You don't need to know a lot about molecules at AS thankfully. So not for the first year anyway. you need to know a little bit more for the second year but not for the first year just to show you look I've changed the axes on the left this is this is obviously relative abundance so I've changed it slightly it's not percentage abundance anymore so we just add up all these masses if we want to work out the total amounts okay molecules are different so we've got our mz there um Molecules are different because when we send them through the mass spectrometer, they actually break into little bits. We call them fragments. Don't worry too much about fragments at AS. But the fragments have a mass. And basically, instead of isotopes that we have in an element, we have these bits. And these fragments obviously form the spectra that we can see here. The most significant thing that we need to know is the M plus 1 peak. Basically, this is... this peak shows the fragments of the that hasn't been broken up so basically if we had say if we had ethane for example this is unfragmented ethane the whole thing is being ionized and the whole molecule has gone through because some molecules aren't fragmented and basically we get what we call a molecular ion peak which is always the last significant peak on the spectra so in this case the last one in here is 50 so we can know that this molecule has a MZ or mass of 50 and so that's pretty much all you need to know about that Okay, just looking at electron configuration, we need to know that electrons are split into four subshells. We have the S, the P, the D, and the F, okay? And we need to know how we write these electron configurations as well. So S is only one orbital, it's spherical, it can only hold two electrons. Your P are like in a figure of eight, and they hold two electrons each orbital, but we have P. We have three P orbitals. So in total... we can fit six electrons in the P subshell. The d orbitals are the, again, you only fit two in each orbital, but we have five of them. So in total, we could fit 10 electrons in the D subshell. And the F block is that funny block that's kind of detached away from the periodic table right at the bottom. These basically have seven orbitals, and you could fit 14 electrons in total here. So let's look at the... The... The shell number first so if we look at an atom the first shell only has one s orbital That's all it has and the maximum we can hold is two so we've done that in green to match In shell number two in an atom and we have the 2s orbital you can see now We've still got an s orbital but now it's the 2s orbital and now we start to get into p orbitals electrons and p orbitals Again, we can fit a maximum of eight electrons in the second shell and because we can have six in the p in the third shell We can have we now have a d orbital and do orbitals member can hold ten electrons So if we have two in the s member, this is three s three p and three d because we have three in the s and Then and so because we have two electrons that can fit in the s we have a two there we have two lots of Three p orbitals or three lots of three p orbitals with two electrons each So and you can see here that we've got five 5d orbitals here which is 5 times 2 and that's 18. Okay so this just basically shows you how it's all structured. Okay so let's look at the electron configuration of an atom. So basically we need to know that they're written like this 1s2. The first number here tells you the shell number. The letter bit here tells you the subshell that we've just looked at before and the number bit tells you the number of electrons in that subshell. So let's look at the electron configuration of iron. Okay, so you can see here that we've got 26 protons. This is elemental iron, so we've also got 26 electrons. So basically we just need to look and see what the electric configuration is. What I've done is I've drawn an energy level diagram to show you how these orbitals correspond to each other in terms of energy. So we have 1s2, basically we've got to get 26 electrons. 2s2, so you can see here that we've got the second shell now. S orbital, we've got two electrons in the s orbital. And we represent them with a box. The boxes, the little arrows show electrons spinning. So spinning in opposite directions. Okay, 2p6, 3s2. And you can see we're filling them up. Remember the p orbital? We can have three orbitals in there. p subshell, so we can have three orbitals. 3p6 and then 4s2. Now it's a bit weird. If you look here, your 4s2 is lower in energy than your 3d. It doesn't have to be in numerical order. But yeah, your 4S is lower than your 3D. So we fill that one first. And then we have 3D6. And so obviously this tells us the electrical configuration of iron. And you can see that configuration there. Okay, so check. They must give or they should add to give 26. So just check. Check them answers there. Okay. We always fill from the lowest energy upwards. So we start from the 1S first. We can't start from 2P or anything. You have to start from the 1S. and we fill orbitals singly as well first then we pair them up so if you look in that top row there we're pointing to you can see that we've got these electrons that are single in other words they don't put that electron in that orbital they prefer to sit separately in their own orbital unless they have to pair up because there's no other orbitals left okay so and this is because of course you've got electrons off the same charge and we've got a bit of repulsion going on okay let's just look at some ions you So with ions you just have to remove the electrons from the highest energy level first. So transition metals behave a little bit differently but we'll have a look at them later. So let's look at the electric configuration for calcium 2+. All this does is this loses two electrons and you lose the two from the 4s. So let's just have a look. There's the 4s2 and there it goes it disappears and what we're left with is your 3p6. because we've taken the electrons from the 4s2. Basically, we need to check the small numbers. So they should give, these numbers here, if you add them up, it should give 20 minus 2, because we've taken 2 electrons away, is 18. If we add all of them up, it should give 18. Okay, so we lose from the 4s. Okay, there it goes and it disappears. Okay transition metals are a little bit different Um, you've got to be careful with these chromium and copper in particular. Okay, so they behave differently So an electron from the 4s orbital actually moves into the 3d orbital to create a more stable half full or full three dubs 3d subshell respectively. Okay, so if we look at chromium okay so the electron configuration of chromium is 1s2 2s2 2p6 3s2 3p6 3d5 4s1 so what we've done is we've removed an electron from the 4s orbital or the 3d orbital should I say so we've removed an electron from there however an electron from the 4s orbital has moved into the 3d subshell so we've created a half full subshell so what we don't do is we don't take from the um basically we don't have this situation here where we take from the d orbital and we still have two in the s orbital so this is a bit unusual really so um so these things behave a little bit differently as well uh so your metal irons these behave a bit strange so if we look at the electron configuration for fe3 plus um So basically what it does is it loses three electrons and two from the 4s and one from the 3d Okay, which is a bit strange So normally you would think you would remove from this one first and because of the sign of energy But when you're removing from a transition metal because these things are so close in energy. We actually remove these ones first Um, it's more stable less energy to do that. Then we start picking away from the 3d. So let's just have a look Um, there it is there. This is the electron configuration for uh iron and what we're going to do is we're going to remove there you go the electrons from there and we've got 3d5 let's just go back look there so we take three electrons the 4s2 goes 3d6 turns into 3d5 let's look at it uh in terms of the numbers if you add them up it should give 26 minus 3 which is 23 so just check the little numbers there look on this diagram there it is loosen the 4s there you go then from the 3d and there's the configuration Okay, so you're going to lose some 4s first. Right, let's look at ionisation. So, this is the minimum amount of energy required to remove one mole of electrons from one mole of atoms in the gaseous state. You must, all these bits which are underlined, you've got to remember, it's always one mole. One mole of atoms, one mole of electrons, etc. Always got to be in the gaseous state as well. I know it looks a bit weird when you're ionising sodium and sodium's in the gas state or gaseous state, but it's always got to be like that, okay? So let's look at sodium. Sodium is Na, forms Na plus plus one electron. This is the first ionization energy of sodium and it's given the value there. You don't obviously have to remember these values. Always include your state symbols, like I say. And ionization always requires energy. So these things are always endothermic. So they always have a positive value. That's always with all ionizations. We need to know about the effects of shielding. This is quite important. Basically, the more shells or electron shells that we have between the positive nucleus and the outer electron, the less energy is required. We've got a weaker attraction here. So there's your positive nucleus. Look at this atom. It's got loads of shells here. There's the outer electron. So trying to take an electron from there is going to be a little bit easier than taking an electron from there because you've got a stronger attraction between the positive and the outer electron compared with this one. This one has more shielding. Okay. Atomic size that's obviously going to play a role as well So the bigger the atom the further away the electrons are from the nucleus the attractive force is weaker So therefore it takes less energy to remove that outer electron and the nuclear charge is pretty important as well So the more protons in the nucleus the bigger the attraction is between the nucleus and the outer electron So this means more energy is required to remove the electron This is particularly useful if you're looking at trends going across a period which we'll look at later on Okay, right. So let's look at successive ionization energy. So this is basically removing an electron from an atom. So we're constantly taking the first electron, then we take the second, then we take the third, then we take the fourth, etc. So that's what we're going to do here. So the removal of more than one electron from the same atom is called a successive ionization. So here's magnesium, we're going to remove an electron from something that's already positively charged. This is called the second ionization energy. it's a little bit bigger compared to the one which would be for the first ionization. So this is the second one here. So it takes a little bit more energy. We're trying to remove something from something that's already positively charged. That's going to take some energy to do. Okay so you can see we've got some distinctive jumps here. So we've got jump here and we've got a jump here. Now these jumps are because we're removing an electron from a shell that's increasingly closer to the nucleus. Because it's close to the nucleus remember this is the nucleus that holds these electrons in So this is going to take a lot more energy to do you can see there's a general increase in energy Moving an electron from an increasingly more positive ion like I said um so yeah ok so let's have a look at some of these now you can see we're removing and these electrons here now these electrons are sitting in the 3s orbital remember these are the ones furthest away from the nucleus so we're starting here these are these two electrons here they sit in the 3s if we look along the next lot are the ones in the second shell now the second shell is much closer to the nucleus so we have obviously we've got six in the p orbital and two in the s orbital But this is what these electrons represent remember what this generally increasing shell Then if we want to remove the ones from the first shell the one closest to nucleus We're gonna need to put a significantly more energy in now for the exam You need to know these jumps you need to be able to explain them So it's all about trying to take an electron for something that's closer to the nucleus Okay, so we know if we look at this element here. So we know this is magnesium because this element has 12 electrons and you can see the 12 down there there is there okay right first ionization trends so we need to know about this in terms of the groups so this is going down group 2 in particular now just to kind of quickly show you the graph first ionization energy this is the energy required to remove one electron from each one of these elements so stop beryllium magnesium calcium strontium barium etc so Basically, the ionization decreases as we go down the group. And this is the reason why. So we've got the atomic radius. As we go down the group, it gets bigger. And the electrons become further away from the nucleus. So this means the attractive force between the outer electrons and the nucleus is weaker. And this means the energy needed to remove that electron is obviously going to be less. Because it's obviously got a weaker force. Also, shielding. How many times has this come up? Shielding is so important. Shielding increases as we go down the group. More shells between the nucleus and the outer shell, so the attractive force is going to be weaker. Really important, that one. And the energy required to remove an electron decreases. So we've got two things here, the ionic radius and shielding. Shielding plays a big role in here. And if we just go back to our history of the atom, this data provides strong evidence for shells. Remember, Niels Bohr said there were shells. And so this is the evidence that would back that up. But, however, it didn't explain data showing going across a period. So this is how we know Niels Bohr's model isn't quite, not quite there. It's not quite the finished article. So let's have a look at the one when we're going across a period. So that's going down a group. So this is going across a period. Now, generally... The ionization energy increases as we go across a period. So you can see from this graph here, there it is. Okay, you can see it's generally going up. I've picked the elements going from sodium all the way to argon. So it's going across a period. Remember, periods are going along the periodic table. And all we're doing is just taking one electron from each one of these elements and measuring how much energy it takes. Now, this is a bit trickier. There's your general increase. Okay, the reason why we have a general increase is because as we go across the period, We have one more proton compared to the previous element. So this increases that nuclear attraction that we were talking about. Shielding is similar. So actually, it's still an important point. We need to say that it's similar in your answers. But it doesn't have any effect really. So yeah, it actually marginally decreases. So because the distance from the nucleus is effectively getting a little bit smaller. But it's not significant enough. to cause too many problems in terms of energy but the shielding is similar. More energy is required to remove that electron so the ionisation energy increases. Now we do have some exceptions here and the examiners are going to pick up on these. So you can see here this one and this one. So we're just going to look at what these exceptions are. So you can see that we've got a decrease in aluminium. That's the first one that we pointed to. This is evidence. For having sub shells okay remember this is beyond Niels Bohr model So the outermost electron in aluminium sits in a higher energy sub shell Slightly further away from the nucleus than the outer electron in magnesium, so if you see here There's aluminium look 3p1 you can see magnesium doesn't have an electron in the 3p orbital But aluminium does now because it's a little bit further away from the nucleus and it's slightly shielded From the 3s orbital this is going to mean that we don't need much energy Basically to to remove it and it drops slightly so Magnesium is out electrons in that 3s orbital so the atomic model neil's ball came up with didn't explain this theory. That's quite important Okay, right, let's look at the next one along. This one's a little bit trickier. This is sulfur. Now, there's a decrease at sulfur. This is evidence for electron repulsion this time in the orbital. So if we look at the element before, which is phosphorus, phosphorus and sulfur, they both actually have electrons in the 3p orbital. So the shielding is the same. So we're not talking about shielding now in this one. The shielding is actually the same. However, This is the energy diagram for sulfur as you can see here. There's the that's the setup there So what the 3p even phosphorus phosphorus would just have three electrons in there And so they've got the same shielding but if we're removing electron from sulfur it involves taking it from an orbital with two Electrons already in it. Okay, so there's the two electrons there now remember what we said last time electrons in the same Orbital repel each other that's not very good in terms of energy. So the electrons repel so less energy is needed to remove an electron from an orbital with two in compared to with one with phosphorus this is the configuration for phosphorus doesn't have that paired electron therefore it takes a little bit more energy this one is paired a bit of repulsion here won't put up much of a fight in terms of trying to take this electron so the energy drops and we don't need as much energy okay so um that is basically it um Like I say, this is just a very brief overview of atomic structure. These PowerPoints, like I say, they can be purchased. So if you just click on the link below in the description box, you'll be able to access them from there if you would like them. But I hope that helps. Bye-bye.

Like I say, all of these things, they are tailored to the specification, so to make sure that they are absolutely watertight, well, nearly, so you can look at the specification, basically it just matches that, okay? So we're gonna start with the Atom. So the Atom is, can see it's mainly made up of protons and neutrons they are very small and they contain that are contained in the middle there you also have electrons that's whizzing around in shells so these orbits the middle of the atom as you can see and you need to know these charges as well so protons have a positive charge as you can see and there it is there and neutrons have a zero charge neutron neutral electrons are negatively charged so they have a minus one charge They're relative masses. You need to know as 1, 1, and 1 over 2000 is the relative mass of an electron.

You can't put no mass. An electron does have a mass. It's just very, very small.

And just to familiarise yourself with the... the elements in particular. So here's lithium for example. The top number tells us the mass number, that's the number of protons and neutrons in the nucleus and the bottom number tells us are just the number of protons.

We call this the atomic number or proton number. So yes all atoms remember are neutral because the number of protons equals the number of electrons. Okay right so let's look at some ions and isotopes.

So You've got to know the difference between these. We'll start with ions first. So ions have a different number of electrons and protons.

So they basically don't have equal amounts like we looked at in an atom. So for example if we look at negative ion, O2 minus has gained two electrons to get a full stable shell of electrons. So you can see here that the protons in this particular one here, that we have eight protons which give a charge of plus eight. Eight neutrons, remember they have no charge. count electrons we've got 10 electrons in O2 minus that means it gives it a minus 2 charge so the total charge on this is minus 2 and that's how we work out them charges okay so this basically allows them to form these stable ionic compounds they gain stability by being attracted to positive ions okay so let's look at a positive ion this is sodium sodium is the opposite it has one electron it's out of shell it's in group one it loses an electron to form a Na plus ion and you can see the same calculation I've done.

Protons, 11 protons it's got, 12 neutrons, which have no charge, and electrons, because it's lost an electron, it now only has 10 electrons, so the charge is at minus 10. Now if you do plus 11, 0, and minus 10, you get an overall charge of plus 1. That's why sodium has got a positive charge. Okay, isotopes, right. These are elements with the same number of protons, but a different number of neutrons. they have a different mass effectively.

So these are slightly different. So let's look at these three examples here. These are three examples of carbon. Carbon 12, carbon 13, carbon 14. And let's just look at the different numbers.

You can see we've got the protons, neutrons and electrons here. Now if we look here, they've got the same number of protons, each all of these isotopes too, but crucially the number of neutrons is different. So carbon 12 has six neutrons, Carbon-13 is 7 and carbon-14 has 8 neutrons. So these are isotopes of each other. Okay, chemically they react the same because they have the same number of electrons, but they have a slightly different mass.

Okay, right, we need to know a little bit about the history. This is the fun bit. Okay, back in 1803, John Dalton, he came up with this idea that atoms were all spheres. A very simple idea. And it wasn't until a good while later, nearly 100 years later, that J.J.

Thompson had a good go at trying to change the model that Dalton came up with. He discovered the electron. And the atom, he said basically the atom wasn't solid and it was made up of other particles and he named it the plum pudding model.

So basically he said that you had the positive pudding bit, which is this bit here, and inside the pudding you had negative electrons, which are the yellow circles there. That's what Thomson came up with. Very shortly after that, Ernest Rutherford caught the atomic bug, I suppose, and he discovered the nucleus.

and also discovered that the nucleus was really really really small and Actually the nucleus contained positive charges So he basically said that most of the atoms empty space and because the nucleus was so small And he said we had a cloud of electrons that was surrounding this nucleus So you can see now we're starting to resemble more like the atom that we know today And he proved this using the gold leaf experiment. Basically, what he did is he fired alpha particles at a very thin bit of gold leaf. And most of the alpha particles went through.

Now, it tells him that most of it was empty space. So that's where he got his idea from. But some did deflect back. Very, very small number.

And so what they must have done is they've hit a small positive nucleus. Because an alpha particle is positively charged. So it's actually hit the nucleus.

bounced back some of it was deflected some of it bounced straight back right back at the um at the the the uh the source of the alpha okay so 1913 again look not many years later um neil's ball had a go of it so they really got the atomic bug here and he discovered a problem with rutherford's uh model and basically he said well the electron cloud could collapse because it would obviously fold into the positive charge nucleus um And so he said, well, actually, you couldn't just have a cloud. You had fixed energy levels. And this is where the shell came up with.

Now, basically, he proved this because actually when he shone electromagnetic radiation, it was absorbed by the atom and the electrons move between the shells. That's what he noticed. And when they do this, they emit radiation when the electrons move back down to a lower energy level. Now, this could only be explained using a shell model. You couldn't explain this using the cloud model.

So this basically kind of cemented Neil's ball idea of the atom. And obviously the model today, we know that actually yes, there are shells, but we now know of the existence of subshells. And basically we can use the ionisation trends to explain this, which we'll look at later. The time-of-flight mass spectrometer.

So this is basically mass spectrometry, but using time-of-flight. The first bit, when you add your sample, it's vaporised. So it can travel through the time of flight mass spectrometer. So we turn it into a gas basically. Then we ionise it.

Now we can ionise this using a, we call it electrospray ionisation. It basically works when we spray a sample through a high pressure jet. It's like putting your thumb over a hosepipe, it's a really high pressure jet. And basically what they do is they pass a really high voltage through this jet.

And this causes the loss of an electron. And what we get is that... gaseous positively charged sample is made and this is really important because we need ionization for the next stage which is acceleration. So you can see there there's the blue particle look they're moving through and so these are accelerated by negatively charged plates on electric fields and basically the particles with a lower mass to charge ratio or mz ratio will accelerate quicker so they'll move through a little bit quicker.

And the next bit is the, basically it's the iron drift. This is a bit weird. Now you can see the red one zipped through there and the blue one's a little bit slower.

But the particles travel through with a constant speed and kinetic energy. So if I just go back, you can see that both are travelling at the same speed and the blue one, they're travelling at a constant speed, sorry, not the same speed, travelling at a constant speed, but they do have, between them, they have a different speed each. but their speed is the same in terms of the atom. So in other words, the red atom constantly travels through.

It doesn't speed up or slow down as it goes through the ion drift. It travels through at a constant speed, but the blue one goes slower. And then the final stage is detection. So once it's drifted through, obviously the blue's going a bit slower than the red, then basically an electrical current is made when the particle hits the plate at the back. And basically ones with a lower MZ and the lighter particles will reach the detector first as they travel the fastest.

So we have effectively separated our different parts out and we're detecting them at the other side. Okay, right. Your definitions, right?

You need to know these. I can't emphasize this enough. Right, relative atomic mass. You can probably read them there, pretty straightforward.

Relative atomic mass is the average mass of an atom of an element when measured on a scale on which the mass of an atom of carbon-12 is exactly 12. Basically, we're measuring... everything relative to carbon-12. The relative molecular mass, so this is very similar if you look the average mass of a molecule when measured on a scale which the mass of an atom is carbon-12 is exactly 12. So this is a molecule instead of an atom and relative isotopic mass is basically the mass of an atom of an isotope with an element measured on a scale in which the mass of an atom of carbon-12 is exactly 12. So there's a lot of carbon-12 being mentioned here you've just got to know these. Ready? Okay.

This is a mass spectra. Okay, so here we are. We've got a this one. We look at isotopes So we've got an element here and it's made up of isotopes now the first thing we need to look at really Is the the axes so you can see here that we've got a mass to charge ratio at the bottom And this is basically the mass of the isotope divided by the charge And most do have a plus one charge and so this makes it the same as the isotopic mass Which makes it? relatively straightforward if it had two electrons knocked off which would be quite rare and then the obviously the mass the charge ratio will be half as much um so that's quite that's quite important um because obviously the z bit um stands for charge so if you've got a double charge it's just the mass of the isotope divided by two okay um the bit on the side read this really carefully i mean this is the abundance um it's always shown on the left but this one it can be written as a percentage or a nominal value so it can be just a relative you Abundance this one's percentage abundance So this means that all your isotopes must give 100% if it's a percentage abundance because obviously you can't get bigger than 100% Okay, um so you can see here if we have the 75 and 25 and get 100% from there Okay, so this spectra shows two isotopes of one element So we've got one element going through here made up of two isotopes and we know this because we've only got two peaks So we've got one isotope that has a mass of 35 and one has a mass of 37 This is assuming obviously they have a 1 plus charge.

So you can see here that the most abundant, which means the most common isotope, is isotope 35, whatever this is. 37 isn't as common. Right, so, and from all this information we can work out the relative atomic mass, which we're going to look at now, which is pretty useful. Okay, so let's work out the relative atomic mass of these.

You need to know this formula. Relative atomic mass is the abundance of... isotope A times by the MZ of A, so that's the mass to charge ratio of A, plus the abundance of B times by the MZ of B. And if you had more than two, you would just literally keep adding up the abundance of A, abundance of B, and abundance of C, et cetera.

You just keep adding them up. So basically, you keep adding loads of these brackets, repeating it for each isotope, divided by the total abundance. Now, because this one's percentage abundance, our total abundance is going to be 100, but it might not be percentage, so kind of look out for that.

Okay, so let's look at this. Relative atomic mass is going to be 75 times 35, because the abundance of A is 75. The mass of A is 35, mass to charge, sorry, is 35. So we do 75 times 35 plus 25 times 37, so that's them two, divided by 100 equals 35.5. And if you're smart enough, you can look in the periodic table and you can identify that as chlorine.

So you've effectively identified your element from your mass spectrometer. So that's pretty nice and straightforward. You can also do it through tables as well.

You don't have to give it through spectra. Similar thing, this one's got more isotopes, as you can see. We don't know what the element is of the relative atomic mass, but we're going to try and work it out. So here's that equation again.

Look, all we're doing is taking the isotope, multiplying it by the abundance, 20.5. There it is, 70 times by 20.5 plus, and then here's your other isotopes I was talking about. So you just basically add them up. It is a percentage of abundance so we divide it by 100. If it's not then we just add these numbers up and divide by the total of their numbers.

Okay so the answer is in this case is 72.6. This means that this element is germanium and you can have a look in the periodic table and check that out. Okay right molecules. You don't need to know a lot about molecules at AS thankfully.

So not for the first year anyway. you need to know a little bit more for the second year but not for the first year just to show you look I've changed the axes on the left this is this is obviously relative abundance so I've changed it slightly it's not percentage abundance anymore so we just add up all these masses if we want to work out the total amounts okay molecules are different so we've got our mz there um Molecules are different because when we send them through the mass spectrometer, they actually break into little bits. We call them fragments. Don't worry too much about fragments at AS.

But the fragments have a mass. And basically, instead of isotopes that we have in an element, we have these bits. And these fragments obviously form the spectra that we can see here.

The most significant thing that we need to know is the M plus 1 peak. Basically, this is... this peak shows the fragments of the that hasn't been broken up so basically if we had say if we had ethane for example this is unfragmented ethane the whole thing is being ionized and the whole molecule has gone through because some molecules aren't fragmented and basically we get what we call a molecular ion peak which is always the last significant peak on the spectra so in this case the last one in here is 50 so we can know that this molecule has a MZ or mass of 50 and so that's pretty much all you need to know about that Okay, just looking at electron configuration, we need to know that electrons are split into four subshells. We have the S, the P, the D, and the F, okay? And we need to know how we write these electron configurations as well.

So S is only one orbital, it's spherical, it can only hold two electrons. Your P are like in a figure of eight, and they hold two electrons each orbital, but we have P. We have three P orbitals. So in total...

we can fit six electrons in the P subshell. The d orbitals are the, again, you only fit two in each orbital, but we have five of them. So in total, we could fit 10 electrons in the D subshell. And the F block is that funny block that's kind of detached away from the periodic table right at the bottom.

These basically have seven orbitals, and you could fit 14 electrons in total here. So let's look at the... The... The shell number first so if we look at an atom the first shell only has one s orbital That's all it has and the maximum we can hold is two so we've done that in green to match In shell number two in an atom and we have the 2s orbital you can see now We've still got an s orbital but now it's the 2s orbital and now we start to get into p orbitals electrons and p orbitals Again, we can fit a maximum of eight electrons in the second shell and because we can have six in the p in the third shell We can have we now have a d orbital and do orbitals member can hold ten electrons So if we have two in the s member, this is three s three p and three d because we have three in the s and Then and so because we have two electrons that can fit in the s we have a two there we have two lots of Three p orbitals or three lots of three p orbitals with two electrons each So and you can see here that we've got five 5d orbitals here which is 5 times 2 and that's 18. Okay so this just basically shows you how it's all structured. Okay so let's look at the electron configuration of an atom.

So basically we need to know that they're written like this 1s2. The first number here tells you the shell number. The letter bit here tells you the subshell that we've just looked at before and the number bit tells you the number of electrons in that subshell. So let's look at the electron configuration of iron.

Okay, so you can see here that we've got 26 protons. This is elemental iron, so we've also got 26 electrons. So basically we just need to look and see what the electric configuration is.

What I've done is I've drawn an energy level diagram to show you how these orbitals correspond to each other in terms of energy. So we have 1s2, basically we've got to get 26 electrons. 2s2, so you can see here that we've got the second shell now.

S orbital, we've got two electrons in the s orbital. And we represent them with a box. The boxes, the little arrows show electrons spinning. So spinning in opposite directions.

Okay, 2p6, 3s2. And you can see we're filling them up. Remember the p orbital?

We can have three orbitals in there. p subshell, so we can have three orbitals. 3p6 and then 4s2. Now it's a bit weird.

If you look here, your 4s2 is lower in energy than your 3d. It doesn't have to be in numerical order. But yeah, your 4S is lower than your 3D. So we fill that one first. And then we have 3D6.

And so obviously this tells us the electrical configuration of iron. And you can see that configuration there. Okay, so check. They must give or they should add to give 26. So just check.

Check them answers there. Okay. We always fill from the lowest energy upwards. So we start from the 1S first. We can't start from 2P or anything.

You have to start from the 1S. and we fill orbitals singly as well first then we pair them up so if you look in that top row there we're pointing to you can see that we've got these electrons that are single in other words they don't put that electron in that orbital they prefer to sit separately in their own orbital unless they have to pair up because there's no other orbitals left okay so and this is because of course you've got electrons off the same charge and we've got a bit of repulsion going on okay let's just look at some ions you So with ions you just have to remove the electrons from the highest energy level first. So transition metals behave a little bit differently but we'll have a look at them later. So let's look at the electric configuration for calcium 2+. All this does is this loses two electrons and you lose the two from the 4s.

So let's just have a look. There's the 4s2 and there it goes it disappears and what we're left with is your 3p6. because we've taken the electrons from the 4s2. Basically, we need to check the small numbers.

So they should give, these numbers here, if you add them up, it should give 20 minus 2, because we've taken 2 electrons away, is 18. If we add all of them up, it should give 18. Okay, so we lose from the 4s. Okay, there it goes and it disappears. Okay transition metals are a little bit different Um, you've got to be careful with these chromium and copper in particular.

Okay, so they behave differently So an electron from the 4s orbital actually moves into the 3d orbital to create a more stable half full or full three dubs 3d subshell respectively. Okay, so if we look at chromium okay so the electron configuration of chromium is 1s2 2s2 2p6 3s2 3p6 3d5 4s1 so what we've done is we've removed an electron from the 4s orbital or the 3d orbital should I say so we've removed an electron from there however an electron from the 4s orbital has moved into the 3d subshell so we've created a half full subshell so what we don't do is we don't take from the um basically we don't have this situation here where we take from the d orbital and we still have two in the s orbital so this is a bit unusual really so um so these things behave a little bit differently as well uh so your metal irons these behave a bit strange so if we look at the electron configuration for fe3 plus um So basically what it does is it loses three electrons and two from the 4s and one from the 3d Okay, which is a bit strange So normally you would think you would remove from this one first and because of the sign of energy But when you're removing from a transition metal because these things are so close in energy. We actually remove these ones first Um, it's more stable less energy to do that. Then we start picking away from the 3d.

So let's just have a look Um, there it is there. This is the electron configuration for uh iron and what we're going to do is we're going to remove there you go the electrons from there and we've got 3d5 let's just go back look there so we take three electrons the 4s2 goes 3d6 turns into 3d5 let's look at it uh in terms of the numbers if you add them up it should give 26 minus 3 which is 23 so just check the little numbers there look on this diagram there it is loosen the 4s there you go then from the 3d and there's the configuration Okay, so you're going to lose some 4s first. Right, let's look at ionisation.

So, this is the minimum amount of energy required to remove one mole of electrons from one mole of atoms in the gaseous state. You must, all these bits which are underlined, you've got to remember, it's always one mole. One mole of atoms, one mole of electrons, etc.

Always got to be in the gaseous state as well. I know it looks a bit weird when you're ionising sodium and sodium's in the gas state or gaseous state, but it's always got to be like that, okay? So let's look at sodium. Sodium is Na, forms Na plus plus one electron. This is the first ionization energy of sodium and it's given the value there.

You don't obviously have to remember these values. Always include your state symbols, like I say. And ionization always requires energy.

So these things are always endothermic. So they always have a positive value. That's always with all ionizations. We need to know about the effects of shielding.

This is quite important. Basically, the more shells or electron shells that we have between the positive nucleus and the outer electron, the less energy is required. We've got a weaker attraction here.

So there's your positive nucleus. Look at this atom. It's got loads of shells here.

There's the outer electron. So trying to take an electron from there is going to be a little bit easier than taking an electron from there because you've got a stronger attraction between the positive and the outer electron compared with this one. This one has more shielding. Okay. Atomic size that's obviously going to play a role as well So the bigger the atom the further away the electrons are from the nucleus the attractive force is weaker So therefore it takes less energy to remove that outer electron and the nuclear charge is pretty important as well So the more protons in the nucleus the bigger the attraction is between the nucleus and the outer electron So this means more energy is required to remove the electron This is particularly useful if you're looking at trends going across a period which we'll look at later on Okay, right.

So let's look at successive ionization energy. So this is basically removing an electron from an atom. So we're constantly taking the first electron, then we take the second, then we take the third, then we take the fourth, etc. So that's what we're going to do here. So the removal of more than one electron from the same atom is called a successive ionization.

So here's magnesium, we're going to remove an electron from something that's already positively charged. This is called the second ionization energy. it's a little bit bigger compared to the one which would be for the first ionization. So this is the second one here.

So it takes a little bit more energy. We're trying to remove something from something that's already positively charged. That's going to take some energy to do.

Okay so you can see we've got some distinctive jumps here. So we've got jump here and we've got a jump here. Now these jumps are because we're removing an electron from a shell that's increasingly closer to the nucleus.

Because it's close to the nucleus remember this is the nucleus that holds these electrons in So this is going to take a lot more energy to do you can see there's a general increase in energy Moving an electron from an increasingly more positive ion like I said um so yeah ok so let's have a look at some of these now you can see we're removing and these electrons here now these electrons are sitting in the 3s orbital remember these are the ones furthest away from the nucleus so we're starting here these are these two electrons here they sit in the 3s if we look along the next lot are the ones in the second shell now the second shell is much closer to the nucleus so we have obviously we've got six in the p orbital and two in the s orbital But this is what these electrons represent remember what this generally increasing shell Then if we want to remove the ones from the first shell the one closest to nucleus We're gonna need to put a significantly more energy in now for the exam You need to know these jumps you need to be able to explain them So it's all about trying to take an electron for something that's closer to the nucleus Okay, so we know if we look at this element here. So we know this is magnesium because this element has 12 electrons and you can see the 12 down there there is there okay right first ionization trends so we need to know about this in terms of the groups so this is going down group 2 in particular now just to kind of quickly show you the graph first ionization energy this is the energy required to remove one electron from each one of these elements so stop beryllium magnesium calcium strontium barium etc so Basically, the ionization decreases as we go down the group. And this is the reason why.

So we've got the atomic radius. As we go down the group, it gets bigger. And the electrons become further away from the nucleus.

So this means the attractive force between the outer electrons and the nucleus is weaker. And this means the energy needed to remove that electron is obviously going to be less. Because it's obviously got a weaker force. Also, shielding. How many times has this come up?

Shielding is so important. Shielding increases as we go down the group. More shells between the nucleus and the outer shell, so the attractive force is going to be weaker.

Really important, that one. And the energy required to remove an electron decreases. So we've got two things here, the ionic radius and shielding. Shielding plays a big role in here.

And if we just go back to our history of the atom, this data provides strong evidence for shells. Remember, Niels Bohr said there were shells. And so this is the evidence that would back that up. But, however, it didn't explain data showing going across a period.

So this is how we know Niels Bohr's model isn't quite, not quite there. It's not quite the finished article. So let's have a look at the one when we're going across a period. So that's going down a group.

So this is going across a period. Now, generally... The ionization energy increases as we go across a period.

So you can see from this graph here, there it is. Okay, you can see it's generally going up. I've picked the elements going from sodium all the way to argon. So it's going across a period. Remember, periods are going along the periodic table.

And all we're doing is just taking one electron from each one of these elements and measuring how much energy it takes. Now, this is a bit trickier. There's your general increase. Okay, the reason why we have a general increase is because as we go across the period, We have one more proton compared to the previous element.

So this increases that nuclear attraction that we were talking about. Shielding is similar. So actually, it's still an important point. We need to say that it's similar in your answers.

But it doesn't have any effect really. So yeah, it actually marginally decreases. So because the distance from the nucleus is effectively getting a little bit smaller. But it's not significant enough. to cause too many problems in terms of energy but the shielding is similar.

More energy is required to remove that electron so the ionisation energy increases. Now we do have some exceptions here and the examiners are going to pick up on these. So you can see here this one and this one.

So we're just going to look at what these exceptions are. So you can see that we've got a decrease in aluminium. That's the first one that we pointed to.

This is evidence. For having sub shells okay remember this is beyond Niels Bohr model So the outermost electron in aluminium sits in a higher energy sub shell Slightly further away from the nucleus than the outer electron in magnesium, so if you see here There's aluminium look 3p1 you can see magnesium doesn't have an electron in the 3p orbital But aluminium does now because it's a little bit further away from the nucleus and it's slightly shielded From the 3s orbital this is going to mean that we don't need much energy Basically to to remove it and it drops slightly so Magnesium is out electrons in that 3s orbital so the atomic model neil's ball came up with didn't explain this theory. That's quite important Okay, right, let's look at the next one along. This one's a little bit trickier.

This is sulfur. Now, there's a decrease at sulfur. This is evidence for electron repulsion this time in the orbital. So if we look at the element before, which is phosphorus, phosphorus and sulfur, they both actually have electrons in the 3p orbital. So the shielding is the same.

So we're not talking about shielding now in this one. The shielding is actually the same. However, This is the energy diagram for sulfur as you can see here. There's the that's the setup there So what the 3p even phosphorus phosphorus would just have three electrons in there And so they've got the same shielding but if we're removing electron from sulfur it involves taking it from an orbital with two Electrons already in it.

Okay, so there's the two electrons there now remember what we said last time electrons in the same Orbital repel each other that's not very good in terms of energy. So the electrons repel so less energy is needed to remove an electron from an orbital with two in compared to with one with phosphorus this is the configuration for phosphorus doesn't have that paired electron therefore it takes a little bit more energy this one is paired a bit of repulsion here won't put up much of a fight in terms of trying to take this electron so the energy drops and we don't need as much energy okay so um that is basically it um Like I say, this is just a very brief overview of atomic structure. These PowerPoints, like I say, they can be purchased. So if you just click on the link below in the description box, you'll be able to access them from there if you would like them. But I hope that helps.

Bye-bye.

Transcript for:AQA Atomic Structure Revision Notes

Transcript for:
AQA Atomic Structure Revision Notes