Transcript for:
States of Matter Overview

[Music] thank you hello my name is Chris Harris and I'm from allery chemistry and welcome to this video on um topic 4 states of matter this is for the cie topic so this is the Cambridge Internationals paper so obviously in this topic as it suggests we're going to look at states of matter so solids liquids and gases and just kind of ear level it a little bit and kind of put some extra bits of information in there so um obviously that's the focus of this it is obviously dedicated to the Internationals qualification so if you're studying that then this is the right place you need to be um these slides here are available to purchase if you wish to purchase them they're great for your revision on the goal or you know for you to look back in your own time if you just click on the link in the description box below you'll be able to get hold of them there they're great value for money and I've bundled them together as well into um into the different topics so for example This falls into the physical chemistry topic for year one so I kind of bundled it there but you can't get the full range if you want for a year one and year two okay so without further Ado let's crack on um so I'm going to look at um states of matter first and let's have a look at an introduction now you might be looking at this thinking this is really easy this you know solids it was a gas does it take what you did you know when you were um 11 probably even earlier than that I would imagine but this is going to be a little bit trickier and you'll see there's some extra little bits that you need to be aware of and obviously at a level you've got to make sure that you're using your key terms and all you have to do in GCSE as well but at a level in particular you've really got to be precise and talk like a scientist I suppose so um let's have a look so obviously they exist a solid slippers and gases so we know that and solids have a regular particle arrangement they're tightly packed together and they can't be compressed as you're probably aware of that and liquids they are not tightly packed together they do have a bit of a random Arrangement um so they are tightly packed together they have a brand of arrangement um it's not a good start is it so they are tightly back together definitely if you try and put that in a syringe you won't be able to squash that um and they obviously fit the container they're in and gases have a random Arrangement so they do have a round of arrangement they fill the space and that they um they occupy they move around freely and they can be compressed um so if you put that into a syringe and put your thumb over the end of it then you'll be able to squash the gas and so obviously when we heat a solid um for example then heat is absorbed by the particles and they start to move more and more and more and they vibrate much more vigorously and we'll see a heat profile graph later on in this video um and they move towards a gas Arrangement because they have more kinetic energy okay so that's basically how we get from solid liquid to gas okay so let's start look at gases here now this is where it gets a bit technical okay so I did warn you it might be a bit might be a bit of it's not too difficult it's all right but it's just uh it's it's a bit more kind of more advanced than what you normally used to so we need to understand the basic properties of a gas first okay and you'll never look at a gas in the same way again after you've gone through this so gases have a lower density okay and that's due to the fact that the particles are obviously um spread out um however they can be compressed for the same reason because they have a lot of space between the particles then we can compress them they don't have strong attractions to each other and so they fill the container they're in quite well so this kind of feeds in from the previous topic topic three when we look at intermolecular forces so um you'll find a lot of the CIA topics that kind of merge one to the other so you'll see there'll be bits in here where I've kind of pulled it from topic three um so yeah so don't be kind of alarmed you're thinking I thought if you've watched topic three and think I thought you just said that is I've got Deja Vu and and actually yes it is basically so there's a lot of overlap between them so um so yeah so you'll probably see some similarities here um so but as a consequence up to the random um and Rapid movement of the particles then we can have a phenomenon called diffusion now if your apologist you you've probably seen this quite a lot with cells Etc and this is obviously no different there's no difference within chemistry but diffusion is affected with the movement of gas particles from a high concentration to a lower concentration and obviously if we look at this diagram on the left here and we can see in terms of the concentration our substance of interest is the red particles here so you can see obviously High concentration on the left as you can see there's loads on there and there's a partitioning line and on the right hand side you've got a low concentration of particles there and hopefully you can see that there so obviously um diffusion is the phenomenon where when we mix them so if we open that partition up then the particles will move from a high concentration to a lower concentration and obviously that will kind of diffuse across and this is quite a unique property fairly unique I mean liquids could do it as well but gases have this ability to kind of mix quite well and diffusion can occur in in gases okay so now we're going to look at something called ideal gases now it sounds sounds quite um I don't know um sounds very um surreal I think is a bit weird it's like well how can you have a gas that's ideal but you can so here we are um so to understand the properties of gas particles um we're going to look at something called an ideal gas molecule uh ideal gas models should I say um and this gets a little bit weird a little bit complicated but for gas to be classed as ideal we must make these assumptions now don't worry too much about these it's just kind of talking about you know the kind of properties of gas it builds up into an equation which we'll see later which you really do need to know um so gas particles have negligible volume okay so um they obviously don't have um you know any kind of um that'll occupy a large large space basically and you know the volume is quite small and the gas particles are equally sized they don't have any intermolecular forces of attraction you would have seen that in topic three um of repulsion with other gas particles so that's what we assume if something is ideal the gas particles will move completely randomly and continuously I.E they don't stop and the gas particles have perfect elastic collisions with no energy loss now for the physicists out there they'll probably know exactly what that means but for the non-physicists basically if you have a collision that is perfectly elastic there is no energy loss whatsoever now in reality when you Collide things together there is energy loss so that's either heat it could be sound it could be light it could be various different forms of energy um in an ideal gas we assume that this doesn't happen Okay something that's ideal so from these assumptions we can then say that ideal gases have the following properties so um the absolute temperature um in Kelvin is directly proportional to the average kinetic energy of the particles okay that's what we can kind of say and then pressure is caused by particles Hitting off the side of the walls of the container it's in so gases under ideal conditions are governed by two laws and you've got boil and you've got Charles okay these two um I suppose scientists physicists probably more likely to be and these two scientists um came with two laws um about gases and they loved gases if you really hate them these people loved gases so Robert Boyle was the first Chap and he said that the product of pressure and volume equals a constant at a given mass and a constant temperature so he said PV which is the pressure and volume equals a constant okay so basically he said there was some correlation there between the two again don't worry you don't need to know the kind of ins and outs of this it's just kind of set in the scene I suppose uh the second uh chap was Jax Charles who was a a French uh I think believe it was French yes he was a French uh scientist and he stated that the ratio of absolute temperature to volume is equal to a constant for specific amount of gas at a constant temperature so he said that um V over T equals some form of constant so these two people had these kind of ideas and like in science a lot of scientists will collaborate with each other and kind of Nick each other's ideas and use their ideas and be inspired by them and this was like the big thing back in the day of these these two chaps kind of existed a few hundred years ago and but for them it was you know they were like kind of the pop stars of the day and they basically come up with these two theories and actually if we combine them together we can actually come up with something called an ideal gas equation so we talked about what an ideal gas would look like or how an ideal gas would behave in reality obviously doesn't really exist Does it you know these ideal gases but it's a it's a it's a theory it's saying right we can have gases which are more ideal than others and we can look at the properties of them and you'll see you will have to use the equation for this as well and but basically it tells us um the number of moles in a specific volume of gas which is quite useful because gases are obviously an important part of chemistry and they basically match the two equations together and this is the one you do need to know this is the ideal gas equation PV equals nrt okay so p is the pressure in pascals V is the volume in meters cubed n is the number of moles R is a gas constant so remember them two chaps came up with the constant we've actually put a number to it this time which is 8.31 joules per Kelvin per mole and T is the temperature of the gas that we're looking at so we need to be able to use this equation and we have to be able to rearrange it and we have to be able to actually use the equation here as well so let's look at example so we're going to calculate the volume in centimeters cubed of 0.36 moles of a gas at 100 kilopascals and 298 Kelvin now PV equals nrt can be rearranged to give V equals nrt over P okay now one of the key things here I'm going to point out some key things as we go through these numbers here so obviously the end bit is the number of models which is not 0.36 moles fine so we know that 8.31 is the gas constant that's always going to be the same that isn't going to change in this example unless the exam board gives you something different but you know that's always the same 298 is the temperature in Kelvin which is here and we're going to divide that by the pressure now the pressure is 100 kilopascals remember though the unit has got to be in Pascal's very very important we have to if you put 100 in there you're going to get something that's out by a factor of a thousand so you've got to convert that to pascals really important the volume we get is in meters cubed okay because that's what the volume is measured in it's meters cubed so this is 8.91 Times by 10 to the minus 3. however it did ask in the question you've got to give the volume in centimeters cubed as well now there is um we looked at in topic two um how you convert from centimeters cubed to meters cubed to decimeters cubed if you're not familiar you're not comfortable with converting units make sure you go back and have a look at that video for topic too it's really really important so it's when we looked at the mole calculations and things like that so anyway so the to convert from meters cubed to centimeters cubed we need to multiply by a million to get that so we take 8.91 Times by 10 to the minus 3 multiply by a million and we get 8910 centimeters cubed so the volume of this gas 0.36 moles of this gas at this pressure and temperature would occupy that much space that's basically what this equation is looking at so um the next slide I've I've removed it from this just to kind of um keep it simple like I say look at topic two for that um as well so important so when you're using the ideal gas equation you should use the units above it's really important as I mentioned before make sure you can rearrange that formula as well you've got to be familiar with that the best thing to do with these is to again practice a lot of this is maths related and there's some equations here it can get very boring if I'm sitting here just working through loads of different examples but work through it have loads of exam questions and practice them until you literally go blue in the face well not literally because that would be quite dangerous but you know what I mean um so the gas constant will be given to you at 8.31 just for Kelvin per mole unless they give you some different but I don't think there will be 8.31 um and standard conditions is obviously 298 Kelvin and 100 kilopascals of pressure okay so right so let's look at some non-ideal gases in so we've looked at some ideal ones now let's have a look at some non-ones so real gases which are just like you know what like everyday gases I suppose they're classed as non-ideal in reality ideal gases don't really exist but the we have a kind of theory to show an extreme and then we can measure something against that so that's why that that equation is important but real gases are not ideal the assumptions were made obviously don't apply so when we increase the pressure of a gas we move the particles closer together don't we so when when we heat when we increase the pressure we're squashing them together and when we cool the particles down they have less kinetic energy so they don't move around as freely okay so obviously that does exist in reality and also we look at this diagram here we can see the difference in how particles behave as we cool or as we increase um the pressure and non-ideal gases do actually behave like liquids um to an extent so this means there must be forces of attraction between the particles so for example you might see now I'm not sure you might have done actually if you look at something like liquid nitrogen or or maybe even liquid oxygen you can actually pour it out when it's really cold and you'll see it kind of flow flow outwards and it looks like a liquid's been poured so for that to happen there's got to be forces between these particles now the ideal gas um scenario assumed that there aren't any forces between these which just really exists does it so so yeah so this has kind of showing you the difference here obviously it does behave like a liquid when we cool it down it kind of clumps together here they have less energy they're a bit more lethargic and they kind of just collect at the bottom if you tip that box or or you obviously get liquid you know behaves like liquid it can be poured so obviously the particles take less space as well and obviously that didn't fit with the ideal gas Theory um so obviously they're not taking as much space as say a normal um gas profile like this okay so just kind of put this in graph form then because the exam board again you don't really need to be too fussed about this it's just the example might give you graphs like this and you might think well what's it talking about and it's just kind of giving you some ideas of how the they could expand on this to use upon um so this graph basically shows something called the compression Factor so if we squash in gases and we're squashing various different gases at a fixed temperature and it's basically going to show what the compression factor is as we increase the pressure so how that changes how does the the factor of compression for these gases changes as we as we change the pressure on it so you can see an ideal gas really would follow a horizontal straight line upwards so to start from here and just work its way upwards directly you know as a diagonal line now you can probably see some of these gases clearly don't follow this profile I mean look at carbon dioxide when we increase the pressure the compression Factor actually decreases and then it starts to increase you know when following a typical ideal profile hydrogen as you can probably see here there we are hydrogen obviously follows probably a closer profile to that but real gases actually have greater ideality at lower temperatures and high pressures so you can see here the back end of these graphs here they're kind of more of a diagonal profile rather than at the start some of these occurring at lower pressures so um and they must have for them to kind of follow this ideality they must have a low molecular mass as well so a lighter the gas the better the more well the more closely it's going to follow an ideal profile and they must have weak attractive forces as well so anything with a bit of polarity in like um you know for example oxygen might have Vander valves between it and night vision will as well carbon dioxide will as well those classed as non-polar molecule but it can have Van Der waals forces as well so you know these forces can exist so if we look here basically hydrogen is the closest to an ideal gas that we actually have you can see it doesn't really deviate too much in the lower pressures and it just follow that classic kind of diagonal um kind of line there so it's the lightest one as well so it's just kind of going through the difference in ideal and non-ideal gases okay so let's move on let's look at liquids then so obviously we looked at gases and that's that's what we need two of our gases really uh liquids they are randomly arranged as we've seen before they are tightly packed together um you know much more than gases of course um and they actually kind of resemble more of a solid than than a gas again if you use a syringe and put your thumb against the syringe and fill that with the liquid and try and push the syringe you're gonna you're not gonna compress that you know that syringe the plunger's not going to kind of dip any further because the liquid is quite tightly packed whereas if you did that with a gas it would compress so um looking at liquid then so the spaces between the particles are much akin to solids and but their random Arrangement do allow diffusion just like you do with gases so they're kind of a bit of properties of gases and a bit with um with solids as well and although they do kind of diffuse at a slower rate because they are a little bit more compact it's not quite as kind of free to move around like gases are so melting is a physical state change from a solid to a liquid obviously you would have seen that so in order for this to happen obviously you've got to put heat energy in it's normally heat you know realistically so you've got to put energy in to break the bonds between the particles in the solid um and obviously the amount of energy you need is dependent on the strength of the bonds that's in there so this obviously feeds back to topic um topic three when we looked at bonding and and the forces between them so have a quick look there if you're not too sure on that so obviously solids go into liquid it's called melting and liquid going to solid is called freezing so that's basically increasing the energy or putting energy into sense of liquid and removing energy which is obviously freezing so I'm kind of going the other way um obviously going liquid to gas this is called vaporization and it's a physical state change um and obviously like with melting energies required to um obviously separate the particles of the liquid to turn it into a gas vaporization is the term that's used to Move It from a liquid to a gas and going from gas to liquid is condensation a bit like what you get in your windows so vaporization occurs on the surface of a liquid um it's very different to Boiling though so just be really careful okay so um yeah water can evaporate into the atmosphere but it doesn't need to be boiling to do that obviously we see that you know everyday life don't we when it rains um and you get water on the pavement let's say then when the sun comes out it evaporates but it's certainly not boiling because we wouldn't survive so you don't need boiling temperatures to turn a liquid into such as water for example into a gas you don't need that okay so that's that's what I want to get that distinct difference out the way okay it doesn't have to be a boiling point for vaporization to occur okay and so but it doesn't require a continuous supply of heat energy to evaporate it though okay so that's that's you know really important so if you had a big vat of liquid for example and you want to get rid of that as quickly as you possibly can and you really do want to apply an extreme amount of heat to that to get it evaporating quite quickly so evaporation happens at the surface of the liquid okay so boiling involves the difference between vaporization and boiling is boiling involves the transfer of heat into the liquid which has a limited rate of transfer so if you imagine a pan of boiling water and you boil the pan you get what vapor at the top of the water don't you see they get that water vapor coming off the top of the pan but there's actually heat energy within the main body of the water and that's because you can see the bubbles can't you see the particles kind of moving around like that so um but um obviously the rate of transfer and sometimes this can happen actually if you've ever done it with soup in the in the microwave it spits quite a lot there's a lot of kind of what we call bumping um and basically some of the energy that's within them in body of the liquid builds up and it spits out of spits the actual liquid out of the pot so um and obviously the reason why it's doing that is because it's just taken too long for the energy to escape from that liquid through vaporization so instead it just jumps out of the pan instead or the ball so um and that's basically where you've got something that's boiling something will boil if the rate of vaporization is not quick enough to get rid of the energy that you're putting in that's basically what boiling is okay so there we are you've learned something new um so vaporization is a lot slower than boiling as you can probably understand and particles on the surface of the liquid leave and they turn into a gas it's always on the surface and this leaves the remaining liquid cooler because some of the energy that was in the liquid is now absolutely been used to break the particles apart or move them further apart and turn it into a gas and this is effectively how when you do exercise and you sweat obviously you've got um you know when you you know when you're running you're you're doing any form of physical activity uh or if it's just a warm day your body temperature starts to get warmer now your body has to try and get rid of that heat energy in some shape or form otherwise you will you know you you won't you'll probably pass out with heat stroke or something so your body produces sweat as a liquid and what that does that liquid coats the surface of your skin and it's designed to absorb the heat from the body and evaporate now the whole point of having sweat glands all over your body over the vast majority of the body is to increase the rate at which that heat can be taken away from the body and I've see the more exercise you do or the harder you work the hotter your body gets the more sweat you produce and so all of this is designed to leave the skin cooler that's behind because the heat energy that was in the body is in the liquid which is now evaporated out of the body so it's a good idea if you ask wetness to try and move some of that sweat out of the way so you can get more you know more sweat that's produced and then that helps to cool the body a bit quicker okay so that's the difference obviously your body's not boiling um you know it doesn't boil but the water or the liquid does evaporate because of um vaporization okay so let's look at some kind of let's look at the kind of physical side of it so the heating and cooling graph so when you're melting a solid through to a liquid and then onto a gas there are changes in um in terms of how the heat energy is used you need to understand what is actually happening to the particles at these different phase changes um so we use something called a heating Heating and Cooling graph or sometimes it's called a heating and cooling curve and it's used to show the temperature of a substance changes as we heat or cool a substance you've got to be aware of kind of the the kind of physical change that's happened and what causes this so you can see on here we've got a graph here and it's showing the different um states of um you know the states of matter you've got your solid which is down here at the bottom you've got your liquids and you've got your gases here and you've got your different um phase change points so melting point and boiling point so this is a heating graph obviously we're increasing the temperature over time the graph can go the other way and go down instead it doesn't really make that much of a difference the kind of the the principles the same so but what you'll probably notice here is actually as you heat a solid into a gas you've got these vertical uh sorry these horizontal parts of the graph here where there isn't actually an increase in temperature at some of these points but the melting point and the boiling point you only actually get a rise in the temperature of the substance you're heating at certain points so what's actually going on here so if we've got if we look at the horizontal points here so this assumes obviously heating the energy at at the boiling point and at the melting point is being absorbed by the substance to weaken the bonds between the particles so the energy that's been putting in at this point I've been put in at this point is not actually being used to raise the temperature of the um material that you're heating up the energy is actually being used to just purely break the forces between the particles to get it into the liquid or gasket that we can see on here so and for that reason if you put a thermometer in your say if you take ice for example and then um you basically put ice in there and then you put a thermometer in you'll notice that the temperature doesn't change for a period of time if you plot that kind of ice melting over time so from obviously solid to liquid and then heated to a gas to steam and if you could monitor of the temperature of them stages you'll see there'll be points where the temperature just doesn't change even though you're continuously heating it it doesn't actually change so for substance that cools obviously we're going the other way we're just releasing energy as bonds are actually being formed between particles so we're just kind of reverse it and flip it around the other way okay so let's look at the the kind of the the other states change which is obviously solid so the final State um so solids obviously these are particles that are densely packed together they do have a regular Arrangement as we've as we've seen before and solids are held together by four different types of bonding as you've seen in topic three so you've got covalent you've got metallic you've got ionic and you've got hydrogen bonding is a type of force it's not strictly really a bond technically but um it's it's a type of force it's a very strong type of force that's why I've kind of included it in here so solids um in a crystalline structure can be represented obviously using a diagram again you might have seen bits of this in topic three and you've got to be able to draw these types of diagrams though it's really important so you can see that um obviously solids you have a positive and then you have negatives that are around them and negatives surrounded by positives so it's really important that they are regular they have a regular Arrangement they're kind of structured you can see these kind of lines and rules here and really important tightly packed together and have that form that classic solid structure so let's kind of summarize some of these different types of bonding again I'm not going to go into this too much um there's bits in here that you would have seen from topic three so I encourage you to kind of you know look watch the topic 3 video really um but this is just really summarizing the different types of bonding in solids here that you you really should be aware of obviously you've got your giant covalents simple moleculars Giant ionics and metallics again look at topic three for a lot more detail on these and that's what I encourage you to look at um you need to be aware though of the differences what the difference is between these you need to know what the state the usual state is at room temperature and pressure do they conduct electricity do they conduct electricity as a liquid for each one of these do they dissolve in water what's them melting and boiling points so make sure you're aware of that as well um just be aware as well that in terms of solubility um non-polar solvents a non-police substances will dissolve in non-polar solvents so for example hydrocarbons like coconut oil for example is an example of a non-polar substance it doesn't dissolve in polar solvents like water so for example if you add cooking oil into water it won't mix it'll separate um but um water will dissolve other non other polar substances such as ammonia for example so so just make sure you're aware of that but again I'm not going to go into it in too much detail here there is a lot more in topic three okay so let's look at some allotropes um again you would have seen some of these already from topic three I suppose you haven't heard of the word allotropes um in in this context so I am going to kind of bring in some of the examples from topic three into this topic here and just kind of put a different context to it so allotropes are basically made from the same element but their structures are different and that obviously leads to different properties um of these substances as well so one of them is um is graphite and again you've seen probably seen from topic three graphite is obviously made from carbon and there's loads of strong covalent bonds joining these carbon atoms together and in between you have delocalized set of electrons in between layers of these carbon Rings um and the Lords because it's got loads of strong covalent bonds then um graphite actually has a really high melting point now the layers between the graphite sheets um obviously connected by the weak forces with the delocalized electrons and this means the layers can slide very easily which is just as well because obviously in a pencil pencil is made from um made from graphite when you write when you put a pencil against paper or you know you write against something then basically what you're getting is is um sheets so layers of graphite are basically coming off um and um obviously that's quite useful because obviously that's what you want from a pencil um the delocalized electrons in um graphite do allow um conduction of electricity so obviously graphite will conduct electricity and the layers are actually quite far apart in comparison to the covalent bond length and obviously that means it's low density so um it's quite light which again is useful for a pencil you don't want something that's weighs an absolute ton um so obviously it's insoluble as well it doesn't dissolve in water the bonds are far too strong to Break um the other allotrope I suppose it's all the troll one of the allotropes um with carbon is diamond diamond has got a very different structure to graphite again I'm a lot of this you've gone through in topic three already but it's bonded four times in this tetrahedral um shape tightly packed and it conducts heat well um and obviously this is obviously diamond is a very very good uh sorry it doesn't conduct um so it's tightly packed Arrangement it's a rigid Arrangement and it allows Obviously good heat conduction and but it's also really strong as well it's a really tough uh material and obviously that's used in um construction for example in in saws like saw blades and drill tips Etc so you know anything where there's going to be you know potential a lot of heat or something which is um you know could could risk um being worn aware you would coat it with diamond to just toughen that up a little bit um the unlike graphite obviously diamonds can be cut to make gemstones so it's it's it holds its shape better than graphite to that extent when it's cut obviously an incredibly high melting point it's got so many strong coveted Bonds on there so it's a very hard substance and it doesn't conduct electricity um you know at all really it doesn't have any delocalized electrons at all um unlike graphite which does and obviously as you probably expected is insoluble in water as well because the bonds are too strong for water to break apart okay some other ones some other allotropes um which you may not have heard of before me or may not have done and graphene is is quite a revolutionary material actually and I think it's obviously it has changed the way in which particularly mobile phones for example and modern Technologies are modern kind of um digital technology um graphene is used quite you know relatively quite a lot but I think we can expect graphene to be used extensively in future technology as well um you know when continuing in with that graphene is just basically graphite but it's just taken one layer of graphite um it's one atom thick so it's incredibly thin made of hexagonal Rings that's that's basically all it is um and graphene is really lightweight and it's transparent and it makes it great for um anything that has to be carried around anything that doesn't shouldn't really be too heavy anyway if it's light it's a good thing it's transparent so that could be used for example on monitors and screens and things like that to touch screens for example um delocalized free moving electrons makes it a really good conductor of electricity so that's good because it's using Electronics as I mentioned before very similar to graphite as you can see um and it's got this delocalized electron to kind of sea of electrons as well that strengthen um so the same obviously delocalize electrons strengthen the covalent bonds and this gives graphene a really high strength property so it can be really robust and it can be used for example to strengthen materials it might be in cars it might be in um you know materials around the house it's lightweight and strong is a really good property to have you can see you know it can you can come up with quite a few uses here so like I said some of the uses are smartphone screens mentioned before and super computers so they can they can transfer a lot of a lot of electricity a lot of data very quickly and aircraft shells so if you make planes lighter they use less Fuel and you could probably go further because they don't have to carry as much fuel or the body main body of the airplane is as light as it can be um bookminster fullerene is another one another also known as Buckyballs um so Buckyballs are um carbon uh carbon atoms and you've got 60 carbon atoms arranged in a um football type structure um again um they're giant covalent structures they're used in a medicine um this is probably gonna be one of the next biggest areas of medicine and the drug drug delivery and development and whereby you can put a medicine in a in a like a cage like this like a buckyball and that cage can then open up when it's delivered to a particular site so for example cancer treatment this could be quite revolutionary in that in targeting specific areas for example cells cancer cells rather than um you know at the minute although we do have specific drugs that Target specific cells they are starting to come through um generally the treatment is quite generic and it can affect other cells um as well and can make people quite ill actually the treatment itself can make them quite ill because it's not specific to that area so this is a huge area of science where um you know chemistry like this so Buckyballs and you know in C60 and Nano tubes and things like this can help really help revolutionize in particular areas of medicine but obviously you know when areas of Technology as well so big big areas some something like carbon and the way it's structured can have big big differences um so the other type as well is um hydrogen bonding um so obviously we need to be aware of that now you again you would have seen some of this in in topic three I suppose this one just goes into some more specific examples here so obviously if you can remember from that topic hydrogen bonding occurs between hydrogen and um three of the most um some of the most electronegative elements in the periodic table and that's nitrogen oxygen fluorine and it's responsible for the unique properties of ice and water okay it's quite an interesting topic this actually there's loads of different kind of things which you can kind of see in everyday life but um so basically you've got water molecules here obviously we knew from the previous topic that they form hydrogen bonds with each other and that gives them a higher than normal boiling and melting point compared with other molecules with a similar molecular mass such as nitrogen and oxygen so what does not exactly a big molecule at all so really it shouldn't have very strong forces you know compared to its molecules of a similar size but they do have this obviously this property that they can hydrogen bond which gives them a bit of a different um you know a different uh different properties to their you know to molecules of a similar size so hydrogen bonding is actually responsible for um surface tension um on liquids and that explains why we get water that beads now if you look at say if you've got a car that's just been lovely and waxed been lovely validated it's got a wax on it and then if you tip water on on that wax the water will be together it'll actually Clump together instead of being spread out it actually beads together and collects in one place and water has a tension as a surface tension and this hydrogen bonding is responsible for creating that that um that surface tension you get certain um insects that can skate across the top of the water they can sit on the water because there's a surface tension there also if you look at divers obviously they've been the um the Commonwealth Games that's been on recently so you know if you look at divers when they try and dive in they'll dive obviously with their hands outwards so they can kind of glide into the water if you try and dive and do what like a belly flop I suppose and go in a bit like a starfish that's going to hurt quite a lot because the tension on the surface of the water applies a force against the body which will hurt and not really advised because you could get winded actually if you do that so don't do that um so hydrogen bonding is also responsible obviously for viscosity of certain liquids so for example um molten glucose and fructose and that's obviously found in fruits and honey um so obviously the reason why they're quite um viscous which just means thick is because of the hydrogen bond ending between the molecules so it just holds its form quite well treacles a good example of that it's just sugar that's been refined and and put into a liquid liquid form so just looking at a little bit more I suppose a little bit more detail into hydrogen button if you just come back onto water and in particular if we look at ice now ice is one of these bizarre um well a bit of a bizarre entity I suppose because if you look at um uh other solids or other liquids when you cool them they normally contract they take up less space and when you heat um say a solid or a liquid they expand so when you heat a solid in particular they expand and take up more space now water is a bit weird because when you cool it down it actually gets bigger ice takes up a larger amount of space than its liquid form and which is which is weird and this is down to purely down to hydrogen bonding um and to an extent where water is actually less dense than as a solid than what it is in liquid form and that is why ice cubes float on top of water because their density decreases as you cool them down and it's thanks to hydrogen bonding that kind of separates the molecules out it arranges them into a regular Arrangement but they're kind of more separated out than what they were as a liquid um and obviously that obviously transpires you can see ice floating on the water and hydrogen halides um so can also show how hydrogen bonding affects boiling points and we can use a graph to eventually look at the difference of these as well so this graph here has got boiling points in kelvin um let me bring the mouse in there we are so boiling point in kelvin um and you've got um hydrogen fluoride hydrogen chloride um hydrogen bromide and hydrogen iodide so this is basically going down group seven which you'll see later on in year one so there's bits here we see there's a lot of crossover but basically it's looking at the boiling points of each of these different substances here and what they are you know over time now one of the biggest things to note is obviously here traffic's got a higher boiling point in any of the other ones despite the fact that it's quite small but HF can hydrogen bond and so more energy is required to pull HF molecules apart than it is for HCL hbr hi which only have permanent dipole-dipole interactions these don't have hydrogen bonding and if you remember from topic three permanent dipole-dipole forces are weaker than hydrogen and so that's why these boiling points are lower than what they are compared to compared to HF there we are okay and some of these may have as we obviously go along here as well so hi doesn't really have that much of a difference in electronegativity so all of these will have Van Der valve's forces as well but um this HCL will have the permanent dipole dipole less so with some of these for example h i there isn't much of an electronegativity electronegative difference anywhere between hydrogen and iodine so yeah so hydrogen bonding obviously um does have an impact in terms of the melting and boiling points of these substances okay so let's look at um some um other examples of hydrogen bonding um this is obviously also responsible for the ability to dissolve Some solid substances into Solutions um so ionic salts for example are a classic example where salts are pretty soluble and a good chunk of them anyway are and they'll dissolve in water and the reason why salts can dissolve in water is actually because water can hydrogen bond so you can see how important water is actually in chemistry you pretty much dissolve everything in water to be fair so let's have a look at this then so we've got your ionic lattice so this is your ionic compound here so you might have seen this again from topic three um you've got your positive and negative ions that have obviously performed in a giant ionic um substance diagonic compound and you've got your water molecule here with your polarities on there so if what happens is actually the substance the bonds are broken to create free moving ions so these ions are then broken apart so this is when you um when you basically add it to the water the ions start to break apart um like that because obviously the charges on the the uh hydrogen and the oxygen and obviously the hydrogen bonds it then Dives straight in um and the water starts surrounding the individual ions and we say these are hydrated and this is effectively we're dissolving the salt into uh into um into water effectively so and it's the hydrogen bonds that exist between obviously the the Delta negative oxygens here are attracted to the Delta to the positive ion sorry and the Delta positive hydrogens are attracted to the negative ions now they keep these ions apart and that's effectively what allows water to uh what allows salts should I say so iron compounds to dissolve into into water so like I said most ionic compounds do do this and obviously we've seen the structure obviously starts to break down and obviously the surround the ions and we call that hydration and obviously for this to happen not because not all salts are soluble um but so for this to happen the new bonds formed here so these interactions here must be the same strength or greater than the ones that we use to break it down in the first place if not it's not likely to dissolve and actually sometimes um you'll know in you know if you do any form of cookery or anything like that that sometimes you want something to dissolve you might have to heat it up you might have to give it some energy to to kind of help it break apart okay so um let's look at the um I suppose another example which is iodine crystals um just looking at another example of solids so obviously these can be um iodine's a bit of a strange one because um iodine is um is a solid at room temperature but actually when you heat it up at room room pressure um it actually sublimes it goes straight to the gas it doesn't actually go to a liquid and but what we need to do is basically understand the structure of of iodine and again from topic three you'll understand that van der valve's forces are the weakest type of forces in a molecule and iodine just has Van Der valve's forces there are no other forces there um but these forces actually help iodine to form a crystal structure at room temperature um so iodine uh there we are so iodine is a is a nice kind of um grayish purplish um solid at room temperature and it's crystalline it looks like crystals um and crystals means that you have um you have a regular arrangement of molecules obviously it's a solid but you have these weak Van Der Val's forces um which hold these iodine molecules together in between the atoms though just to make a distinguish the difference between them you have strong covalent bonds that hold the two iodine atoms together as well so it's the um weak forces though that hold iodine as a solid together so just make sure you're aware of this and you can kind of describe the difference between the forces between the molecules and which are these here and the bonds that exist between atoms and and the reason why obviously iodine is a is a solid okay and then just the final bit really is metallic structure and again you would have seen this in topic three and I urge you to kind of look at topic three or watch the topic three video for a little bit more detail on this um but just briefly just to run through the metallic structures obviously Metals um are solids they will most of them are apart from Mercury um but for the purpose of this they're solids they have that positive metal charge with the delocalized Sea of electrons um they um they are obviously have this electrostatic attraction between the positive charge and the negative electrons um the more electrons you have between the meta lines um between the mat lines and the electrons obviously the more electrons that are floating around the stronger that attraction is and that can be classically seen with um sodium can give up one electron in the metallic structure whereas magnesium can donate two electrons per atom of magnesium and magnesium is a higher um boiling point or higher melting point should I say than sodium sodium is quite soft you can actually cut it with a knife it's like butter it's a bit bit weird um metals are obviously good thermal conductors as well they have that delocalized electrons so they can transfer that kinetic energy um across and the good electrical conductors again because they have this delocalized electron system a bit like giant ionic compounds with the delocalized ions and they have high melting points because you've got this really strong electrostatic attraction between the electrons and the positive metal lines they're also insoluble as well they don't have um the um well the bonds can't be broken the interactions between the delocalized electrons and positive metal lines um can't be um broken using water so they're kind of insoluble and they're my level as well so we can hit um say a row of these atoms so let's see if we hit the middle atom here and nudge the atoms across like further along down the line the electrons will kind of shift to kind of hold the um the bond in place so luckily we can Hammer metals and they will be mold into the shape that we want them to be okay and that's it so that is states of matter quite an interesting topic some quirky little bits in there I would say um make sure obviously like I say a lot of these Interlink with each other so it is obviously quite important that you kind of look at this in context um these slides are available like I say from my test shop if you click on the link in the description box below you can have a look at there great for revision and you can have a copy of these on your phone or tablet to use as and when you wish um for now that's it bye bye