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Current time:0:00Total duration:12:35

Video transcript

we saw in the last video that if we defined enthalpy H as being equal to the internal energy of a system plus the pressure of the system times the volume of the system and this is in almost arbitrary definition but we know that this is a valid state variable that it no matter what you do in terms of how you get there you're always going to have the same value because it's the sum and product of other valid state variables but this by itself isn't that useful or that intuitive but we saw in the last video that if you assume constant pressure and that's a big assumption but it's not a unreasonable assumption for most chemical reactions because most chemical reactions you know we're sitting at the beach with our beakers and they're exposed to just standard temperature and pressure or at least some pressure that it's not changing as the reaction occurs if we assume constant pressure we saw that the change in enthalpy that the change in enthalpy becomes the heat added to the system at that constant pressure that P there is just to show you that hey this is just when we're just what this is this is assuming that we're dealing with heat being added at a constant pressure fair enough so how can we how can we use these concepts in any useful way let's say that I had some carbon and in its elemental form is graphite and I add to that I have to have a mole of carbon and I add to that two moles of hydrogen in its elemental form it's going to be a gas it's going to be as a molecule right if I just have a bunch of heart of a bunch of hydrogen in its gaseous state let's say in a balloon I'm not going to have individual atoms of hydrogen they're going to bond and form these diatomic molecules and if I react them I'm going to produce a mole of methane a mole of methane ch4 but that's not all I'm going to produce I'm also going to produce some heat I'm going to produce I'm going to produce 74 kilojoules of heat plus 74 kilojoules of heat when I produce when I produce that one mole that can't do a lowercase K for the Killough when I when I produce that one mole of meth so what's happening here so first of all how much heat is being added to the system and let's assume that this heat just gets released from the system that this isn't an adiabatic process I haven't insulated the system from anything but this just just gets released it just goes away it gets released so my question is how much you know I started off with this container I guess we could call it that's a the standard you know it sort of fixed pressure and maybe I had a bunch of well I wanted to do the carbon I'll do it in grey I have a bunch of solid carbon laying around maybe some type of dust and then I have some hydrogen molecular hydrogen gas each of those dots there's a two hydrogen atoms and I don't know maybe I shake it up or something make them react and then I get a bunch of methane and then I get a bunch of methane I get a bunch of methane gas I'll do that in green so now I just have a bunch of methane gas and I released 74 kilojoules I released 74 kilojoules so how much heat was added to the system well we released heat from the system we released 74 kilojoules so the heat added to the system the heat added to the system was minus 74 kilojoules minus 74 right if I asked you that heat released then I would have said 74 but remember we care about the heat added to the system is 74 kilojoules and I just showed you that that's the exact same thing as the change in enthalpy as the change in enthalpy as the change in enthalpy so how can we think about this what is the enthalpy of this system relative to this system well it's going to be lower right because if you take enthalpy so the change in enthalpy is the enthalpy of your final system minus the enthalpy of your initial system and we got a negative number we got minus 74 kilojoules so this has to be lower than this by 74 kilojoules so H this enthalpy right here is less than this enthalpy right here so if we are actually draw it on a diagram if I actually draw the reaction let's say that this is just I know this is just time or something this is as the reaction proceeds that axis and on the y axis I'll draw enthalpy so the reaction starts off at your initial enthalpy H I and that's this state right here so you start there I'll do it in the yellow of that container so this yellow I'll do it right there you start there and then now toner you shake it up or I'm not going to go into the activation energy so it might have a little hump and all that but who knows but then we end up at our final enthalpy we have this final enthalpy right here after the reaction is occurred that's this date right here this is H final so you can see you've had this drop off and enthalpy and what's interesting here is is that not so much what the absolute value of this enthalpy is here or the absolute value of this enthalpy here is but now that we have enthalpy we can kind of have a framework for thinking about how much heat energy is in this system relative this system and given that there's less heat energy in this system than that energy system we must have released energy and it you know to some degree I told you that from the beginning right I told you that energy is released and the word for this we use this exothermic exothermic now if you want to go the other way let's say we wanted to go from methane and go back to its part you have to add heat into the reaction you would have to if you wanted to go backwards through this reaction go upwards you would act to add you would have to add that heat content to get that positive Delta H and then you would have an endothermic reaction so if a reaction releases energy exothermic if a reaction needs energy to occur its endothermic now you might be asking Sal where did that energy can't come from so I started at this enthalpy here and enthalpy has this weird definition right here and then end up at that other enthalpy here and as you see enthalpy you know the pressure we're assuming is constant let's say the volume isn't changing much in this situation or maybe doesn't change at all so most of the change is going to come from the change in internal energy right there's some higher internal energy here and some higher internal and some lower internal energy here that's causing the main drop in enthalpy and that change in internal energy is really a conversion from some potential energy up here into the heat that's released so there was some heat that was released 74 kilojoules and so our internal energy dropped and what all of this does is it gives us a framework so that if we know how much heat it takes to form or not form certain products then we can kind of predict how much heat will either be released or how much heat will be absorbed by different reactions and so here I'm going to touch on another notion the notion of heat of formation or sometimes it's change in enthalpy of formation so the way they talk about it is the change in enthalpy of formation and it's normally given at some standard temperature and pressure so you put a little usually it's a nought sometimes it's just a circle in there and what that is is how much what is the change in enthalpy to get to that to get to some molecule from its elemental form so for example if we want it for for methane if we have methane there and we want to figure out its if we want to figure out its heat of formation we say look if we form methane from its elemental forms what is the Delta H of that reaction well we just learned what the Delta H of that reaction was it was minus 74 kilojoules which means that if you form methane from its elemental I guess building blocks you're going to release 74 kilojoules of energy that this is an exothermic reaction exothermic reaction because you released heat you also have this you can kind of say that the methane is in a lower energy state or it has lower potential energy than these guys did and because it has lower potential energy it's more stable I mean one way to think of it is you know if you have a guy you know you have a mountain here and then it's down here and you have a ball you have a ball and this isn't you know a complete direct analogy but the analogy to potential energy is look when you're down at a lower potential energy state you tend to be more stable and so in in the everyday world if you have a bunch of methane sitting around the fact that it has a negative heat of reaction or heat of sorry a negative heat of formation or a negative I should say standard heat of formation because I have that not here or a negative standard change in enthalpy of formation those are all the same things tells me that methane is stable relative to its constituent compounds and actually you can look these things up you'll never have to memorize them but it's good to know what they are and I copied all of this stuff actually let me get the actual get the actual tables from Wikipedia down here I did all of these directly from Wikipedia these give you that the standard heat of formation of a bunch of things and if you look if you look down here for let's see if they have methane right there this is what we were dealing with they're telling us essentially the Delta H of that of the reaction that forms methane they're telling us they're telling us that you know this point table right there tells us that if we start off with some carbon in a solid state plus two moles of hydrogen and a gaseous state and we form one mole of methane that if you take the enthalpy here minus the enthalpy here so the change in enthalpy for this reaction at a stet standard temperature and pressure is going to be equal to minus 74 kilojoules per mole and this is all given per mole so if you have a mole of this two moles of this and to form one mole of methane you're going to release 74 kilojoules of heat so this is a stable reaction now there's a couple of interesting things here and we'll keep using this table over the next few videos you see here mono atomic oxygen monta has a positive has a positive standard heat of formation which means it takes energy to form it right that if you have a reaction let me just say the reaction I'll write it this way one half of molecular oxygen as a gas to go to to go to one to go to one mole of oxygen just as a kind of in its gaseous state this tells us this tells us that this state has more potential than this state and in order for this reaction to occur you have to add energy to it you have to put the energy on the other side so it has to be you would put a plus so right here you would have to say plus 249 joules so you might say hey Sal that doesn't make sense oxygen is just oxygen why why is there a heat of formation of oxygen and that's because you always use the elemental form is your reference point so oxygen if you just look if you just had a bunch of oxygen sitting around it's going to be it's in the o2 form if you have a bunch of hydrogen it's going to be h2 if you have a bunch of nitrogen is going to be n 2 carbon on the other hand is just C and it tends to be in its solid form as graphite so all heats of formation are relative to the form that you find that element when you have pure version of it not necessarily its atomic form although sometimes it is it's atomic form now in the next video we're going to use this table which is a very handy table I cut and pasted parts of it to actually solve problems in this last video I gave you the heat of formation and we just thought about it a little bit in the next few videos we're going to use this table that gives a standard heats of formation to actually figure out whether reactions are endothermic meaning they absorb energy or exothermic meaning their release energy and we'll figure out how much
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