 Felly, rhaid i'n gwneud y cyd-lectu llwyfnol o'r cynedig cyd-lectwyr. Oeddaeth e'w amddangos a'ch cael ei gynaredd. Rhaid di'n gweithio'n gwneud ei wneud am lefel awl, o'r holl gweithio'r llefan, felly efallai'r llefi ari Institute i ni wneud am lefel cenedig, sydd lefanaeth fod yn mynd i chi zeithio. Ond oes, roedd hynny'n ceisio cael cael llefano agweithio, a'n mynd i gweithio gan y cerdau cyd-lectwyr o'n ymddangos, For a start we're going to look at what is a chemical reaction, you might know that but we're going to look at some very specific parts of what the definition of a chemical reaction is that we are going to need. Then we're going to review some basic concepts of kinetics and then look at what factors affect the rate of reaction. Now the reason that we're doing it in this order is that we're going to build on what a reaction is from an anatomic level, from what a molecule is with collision theory and so on, and then how the factors affect this. So we're actually really going to build upon previous things. Now if you watch that introduction video on Johnstone's Triangle, if you haven't go watch it, it's a really interesting subject. You realise that we're going to be mostly working in this whole submicroscopic realm of chemistry, the world of atoms and molecules and really small things as opposed to the macroscopic world. We'll do a little bit of maths and symbolic stuff soon. If this means nothing to you maybe go look up Johnstone's Triangle or watch that intro video. This is just to try and help orientate us what are we going to be looking at. We're going to be looking at it entirely from the molecules perspective. So what is a chemical reaction? We have a vague idea of what a chemical reaction is but from the microscopic perspective, from what atom is, it's a collision. It is literally molecules thumping into each other and forming a reaction. Now they've got to hit with the right orientation and speed as we'll see in a moment but it is a collision between chemical entities and the chemical reaction part of it comes from a rearrangement. So that rearrangement has got to be kind of important especially because if we rearrange the atoms around we can say get a new bond formed and that new carbon oxygen bond can be seen say in infrared spectroscopy. It means we can monitor the progress of this reaction. We need the right orientation. Now as chemists who do a lot of stuff on paper we often forget that molecules are actual entities that have to hit each other. We just write things down like this and then forget about it but in reality they've got to collide with the right orientation as well. They are physical objects so if they bounce around in the wrong place doesn't matter if they've got the right energy, doesn't matter how fast they're going or whether they do collide a reaction won't happen. It's got to be the right orientation. This each side of the hydroxide for instance can't do anything to the cl. It must entirely happen from the side of the molecule. So this is known as steric factors so if you have not looked this word up before, sterics do look it up. It is all about molecules getting in the way of each other purely because of their size. Don't worry if you don't know it, I went six months as an undergraduate without knowing this word and then I had to be cleared up in the tutorial. It was quite awkward anyway. So we need the right energy so if there is a collision and it is in the right place it's not always going to actually happen so if the molecule comes along with quite a lazy pace and they just tap each other there's not going to be a reaction. Now the reason for this is fairly straightforward. If you think of what an atom is it's got the nucleus in the middle and then it's got this negative field of electrons. Now another atom over here that's also electrons everywhere they're negatively charged with obviously a positive charge nucleus to balance it out but that's going to be a repelling force. Atoms by and large want to force themselves away from each other because the clouds of electrons repel and if they're repelling they're going to require some energy to overcome that repulsion and get close enough together so that the attraction between electrons and nuclei actually form a bond properly. So if you can't hop over that energy barrier there's not going to be a reaction. So from a microscopic perspective a chemical reaction is a collision between one sometimes more molecules. It's got to be in the right orientation and it's got to have a certain energy associated with it. So how does this all reply to kinetics? We're going to look at concentration. This is very specific for kinetics we always work in concentrations. We're going to look at rare constants then reaction orders and then finally the kind of the the complex one molecularity we'll do quite a bit on this later. So concentration. Concentration is an amount per volume. This is far more useful to us in kinetics especially than it is of an amount. We don't just want one mole of a substance or we don't want one gram of a substance for instance we want a specific concentration because that is what we're interested in. Now it's perfectly valid to measure concentration in mass or weight. Sometimes you see the word specific written down. So if you see a term that is specific that's what it's called. It usually means it's measured in mass and grams rather than moles or we do it in moles per volume and then moles refer to discrete chemical entities. So in this case in this box we have one, two, three, four, five, six, seven, eight, nine, ten, very discrete chemical entities, ten hydrogen molecules and this one we have one. So this is ten moles per the box. This has one mole per the box but this is neon here and that is 20 grams 20 moles 20 grams per mole sorry so both of these boxes would weigh the same even though there are a different number of moles. So they would react very differently even though they carried the same mass. So as chemists working in kinetics we are only really interested in moles per volume. We'd never really want to work in mass or weight. We're not interested in that we want definitely moles per volume. So reaction orders if you've done a bit of chemical connects before you'll know that rate that is the speed of the reaction is defined by a rate constant k multiplied by two concentrations. So a reaction order is its dependency on concentrations so a rate is proportional. So that's the proportional symbol if you've never come across it before to concentration but is it concentration raised to a certain power? So for instance if it's directly proportional its concentration is raised to power one. If it's it could actually be proportional to the concentration squared for instance and this is what we mean by first second or third order. So in this reaction we've got two things coming together to collide together and their rate law is k times two concentrations. Now we don't usually draw the little one in the corner just leave it implied but it's there if you have x and x squared or x cubed and so on that x there has a one in it that's just implied. So if we add these together one plus one that's two it is second order overall it works to the square of two concentrations so if this was a bimolecular reaction between two things it would be the equivalent of being proportional to the concentration squared so proportional. And individually it's first order so with respect to an individual reactant it is first order for both cases so first order with respect to the chloromethane and the hydroxide in there. In this case things look a bit different so we've actually got two plus one here that's three so it is third order overall this xenon flaring interaction and first order with respect to flaring and second order with respect to xenon. Now you might have noticed something a bit funny about that rate equation there. For a start this reaction is xenon plus three f2 it goes to xenon hexafluoride if you care. And you'll notice that there's a three here and a sort of an implied one there so why is the one next to flaring and the two next to xenon? Well they're simply because this rate doesn't necessarily reflect this reaction this stoichiometry of the reaction can't predict what the rate actually is only in the very simplest most elementary of cases does it have anything to do with it. So this can only be determined by experiment we have to run the reaction change the concentration of xenon change the concentration of flaring and see what happens when the rate doubles when we double a concentration it means it's first order so we'd stick a one left for instance. So we can only get this experimentally it has nothing to do with the stoichiometry it does give some hints to the reaction though because it means if the reaction order is like that it means you know its second order is with respect to xenon a really simple straightforward reaction like a dentition that stepwise is probably not what's going on something else is going to be going on in there instead. So there are a few hints it gets a bit complicated but it does hint at what could possibly happen in the reaction mechanism if we look at the reaction order and think a bit carefully about it. So rate constant this is the other part of that rate equation k. Now this k is effectively just a factor that converts the two concentrations to a speed we know that rate is proportional to the concentration we just don't know how much so k covers that and k is unique to a particular reaction for a particular temperature. So if we ran this reaction with chloramethane NOH repeatedly with different concentrations of the two reactants we would always get the same k we would calculate it out if we changed the temperature k might go up or down but k is generally completely independent of the concentration so this value of k and that value of k are not going to be the same number they are going to be different and unique and representative of that reaction entirely. In fact in these cases because one is second order and the other is third that value of k will have different units we'll do that as a curly k just to emphasise that it's small so they're going to have different units they're not going to be directly comparable but there will be unique and characteristic of that reaction. So how many molecules take part in the elementary reaction? This has a tendency of causing a little bit of confusion amongst people so we're going to take this quite slowly now we define a rate of something we write it like this so DOH that's a concentration change with the respect of time so we'll cover a bit of the maths about what that means later but for now this just means rate this is our definition of rate it is a differential equation and if one molecule of OH is disappearing here one molecule of chloromethane is also disappearing so the rate of change of that CH3Cl by dt are both equal they are equal and one of this disappeared one of this disappears and then at the same time one of these appears as well so that's equal to what I'll make the pluses explicit the change in the methanol dt and also equal to the change positive of CL dt so those are all equal it's very straightforward one disappears one reappears all the rates of change are the same er this second reaction oxygen plus hydrogen slightly different because when well we'll we'll do the simple one first two of these disappear two hydrogen molecules disappear two water ones appear in the reaction so that's who comes to say one goes away and another appears in action so DH2 by dt is equal to DH2O by dt so those rates are exactly the same what about oxygen what is that equal to it's obviously not going to be equal to sorry oh make sure I keep my negatives in here uh H2 by dt it is not going to be equal so we have to think about what are we going to measure this as now people get confused by this they wonder do we have to multiply by a store country do we have to divide by it the way to do this is to just look at the reaction think about what's going on take it slowly and kind of say it out loud what's happening so what's happening is when one of these disappears two of these disappear with it so if oxygen's concentration must go down by one mole then the concentration of hydrogen must go down by two moles so hydrogen is disappearing at twice the rate twice the rate of O2 how do we represent that mathematically oh we can therefore put a times two here so double the rate of oxygen we get the same rate as the change in hydrogen or we can bring that two to the other side and say it's times by a half the same these are two ways of expressing the exact same thing the rates of individual change remain the same what we're changing is how we're relating them by a little multiplication factor so say it out loud this disappears at half the rate of this or this is a twice the rate of that this is how you represent it mathematically so for a complete pattern here the plus H2O must be at the same rate as hydrogen so that is also times by half so whether they are threes or twos or anything around here we'll cover a couple of more examples a bit later but you have to remember that these rates are related by fractions on occasion if the stoicometry isn't one to one to one here it's one to one to one to one it's really easy so that is the summary of what we're going to be doing basic concepts of kinetics so now we're going to cover how do we influence the rate of reaction from the perspective of a molecule well firstly concentration this is effectively just a probability argument you can see here these these molecules for instance they could wander around this box for a while and they can keep piercing around and then suddenly they hit and have a collision these ones while they're a bit oh now it's collided here it can go now it's collided here we can wander around oh it's collided again you can see that as something is becomes more concentrated and this is concentration it's per amount per volume not an absolute amount as something becomes more concentrated the odd of a collision happening increase you still have to have the right energy and the right orientation of course but if you've got more collisions the odds of having a successful collision increase by default so we also have speed so this is why when we're looking at from the molecules perspective what determines whether it can have a successful collision it's got to be able to hop that energy barrier in order to have a chemical reaction so where does it get that energy from well literally at the microscopic level it is the physical speed of the molecule so these ones are going around quite loosely quite slowly these ones are zipping around with a lot of speed if they collide with each other the odds of a successful collision are much higher these ones are more likely just to slap off each other because they're going much more slowly so we represent this kind of graphically with something called an activation energy so if we have a graph here and energy goes up that's energy this is the activation energy ea sometimes it's um delta g or delta h depending on how you want to define it this little double dagger on top saying that it's a distance to a transition state that is different to the things you get from thermodynamics for instance which would be a equilibrium based idea so something that has a lot of speed and a lot of energy has enough energy tops straight over here something that has a little bit of speed will just lazily get up a certain way and then fall back down to reactants again so that's from a microscopic perspective in the lab that manifests as temperature so when we come to do the Maxwell Boltzmann distribution we will get a more concrete idea of what the relationship between the speed of a molecule and temperature is but until now you just need to know that temperature that we feel in the lab translates into the molecules moving faster therefore they collide with more energy they can hop over an activation barrier and react more efficiently now so as in complexity i said that chemists have a tendency to forget the molecules or physical things that interact and move so look at this reaction for instance this cl has to attach to a delta positive carbon whether that reaction goes ahead or not is something that you can ask organic chemists about we're just interested in this steric factors here so that's that word against sterics um let's have a look at this carbon hence why i've drawn them as spheres rather than as abstract diagrams well see how i can't get at this side of the carbon because the oxygen's in the way so that's pretty much out of the question i can't get at this side of the carbon because there is another carbon atom in the way it probably can't get at that side because there is another carbon atom in the way and side on probably won't do much so you can see there's a very tiny slither of this back sphere that the seal can actually hit so if you work that out that's probably less than 10% of the area of that sphere so think about it one in 10 times this seal just hits it randomly from any direction one in 10 times it can't do anything no matter what and whether how much energy it's got or how fast it's going it's not going to create a successful reaction it has to hit it in the right way not the wrong way so all else being equal that's going to be quite a slow reaction now compared to this faster one here cl they'll all draw it out cl radical plus o2 this is the kind of thing you might see in the upper atmosphere that's sort of a gas phase reaction and goes quite quick partially because cl is completely spherically symmetrical uh look at it you can attack it from any end it doesn't really care it can take it anywhere um oxygen meanwhile well if it needs to attack end on it's got that entire half that the cl can attack it's got this entire half that it can attack and to be honest it can probably come from the sides or in the middle as well anything so I wouldn't be surprised if this reaction could go ahead no matter what and the collision was so there is very little restriction in terms of orientation for this reaction to go ahead so all else being equal that kind of reaction will go very quick uh this kind of reaction will go very slowly so that's review if you were paying attention we learned that what is a chemical reaction well at the microscopic level the molecular world it is a collision between molecules it also requires energy and that energy is carried by speed of molecules then when we looked into your kinetics we found that the reaction order says that the rate is proportional to a concentration and then we found that the rate constant is basically turns that proportional sign into a k and then we looked briefly at molecularity knowing that one rate is not necessarily equal to another rate you have to take into account the fact that more than one thing can be reacting at once so factors that influence the rate of reaction from the atoms perspective and the molecules perspective then concentration that's a probability argument the more concentrated things are the higher the probability of a collision and then speed which at the lab scale the macroscopic world manifest as temperature so the faster they're moving the hotter they are the higher the chance of an actual successful collision and then size and complexity so simple molecules that don't really care about their orientation are going to react quite nippily and quite quick molecules that have a lot of bulk and steric hindrance around them to try and protect these areas they're going to move quite slowly so that's it for kinetics the very first screencast lecture this is just basic concepts revision that you need to know before we move on because we're going to study all this in a little bit more detail and in the lectures we're going to do a bit more interactive work and a bit more interesting data processing as well we're going to try and take these points turn them into problems and then solve those problems so I will see you there