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Bond enthalpy and enthalpy of reaction

Introduction to bond enthalpy, and how to use bond enthalpies to calculate enthalpy of reaction. 

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  • orange juice squid orange style avatar for user litsareth
    At where it mentions how bonds require energy in order to be broken and vice versa, why is it opposite for ATP bonds? Because when ATP bonds are broken, energy is released...
    (13 votes)
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  • mr pants teal style avatar for user Mah bab
    Why is she saying 3 C-C bonds are broken? Can't we just say 2 of the bonds are broken leaving just a single C-C bond?
    (1 vote)
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  • duskpin ultimate style avatar for user Mohit
    why is this method applicable only when gaseous forms are reacting?
    why is the value of a bond enthalpy measured only for a gas?
    thankyou
    (6 votes)
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  • blobby green style avatar for user Jade Ng
    why did she only focus on the carbon triple bonds and hydrogen single bonds? what about C-H bonds?
    (5 votes)
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    • piceratops seed style avatar for user Edson Broderick P. Lim
      She doesn't consider those C-H bonds which occur both in the products and the reactants since the energy released by the formation of those C-H bonds in the product would just be the negative of the sum of the bond enthalpies of the corresponding C-H bonds in the reactants. leading to no change in enthalpy, so the change in enthalpy of the reaction would only come from the new C-H bonds and the other bonds.
      (3 votes)
  • blobby green style avatar for user Rachel Ribeiro
    I learned in biochemistry that energy is stored in bonds and when you break a bond you release that energy and when you form a bond you are storing energy (ATP being formed and broken being an example). But in the video she says that breaking a bond requires energy and forming a bond releases energy. Could someone explain why making ATP from ADP requires energy, but here making a bond releases energy?
    (4 votes)
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    • piceratops seed style avatar for user RogerP
      I think you must have misunderstood in the biochemistry course. Breaking a bond always takes energy and forming a bond releases energy. I can't comment on your specific example of making ATP from ADP, but I assume that when you total everything up from breaking and making bonds, then the answer is a positive number - in other words, in the overall reaction, energy is needed to convert ADP to ATP. Conversely, energy is released in the reverse reaction.
      (1 vote)
  • leaf green style avatar for user Ghoshofx1
    bond energy and bond enthalpy are the same thing right?
    (3 votes)
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  • primosaur seed style avatar for user Arthur  Teymouartash
    Is there any specific way to know that there a triple bond between the carbon atoms in C3H4 or do I just have to memorize all of them?
    (2 votes)
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    • blobby green style avatar for user Amogh
      Propyne, which is C3H4, is an alkyne. You can figure out that Propyne is an alkyne because it ends with the suffix "yne". All alkynes have at least one triple bond between two of the carbon atoms in its carbon backbone. You also know that Propyne has three carbon atoms in its backbone because the prefix "Prop" means 3 in organic chemistry.
      (3 votes)
  • blobby green style avatar for user Andrew

    I'm confused because San in "Hess's law and reaction enthalpy change" video added the products side and then subtracted the reactants.why is she doing the either way around?
    (3 votes)
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  • leaf green style avatar for user Harasees Singh
    As we all must be knowing by the time that carbon is a more electronegative atom than the Hydrogen. While breaking the bonds why doesnt she consider electronegativities of corresponding atoms and gives both electrons one to each of atoms
    (2 votes)
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  • piceratops ultimate style avatar for user Bailan
    When the carbon and hydrogen broke their bond, why didn't the more electronegative Carbon take both the electrons??
    (1 vote)
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Video transcript

- [Voiceover] We're gonna be talking about bond enthalpy and how you can use it to calculate the enthalpy of reaction. Bond enthalpy is the energy that it takes to break one mole of a bond. So one mole of a bond. So different types of bonds will have different bond enthalpies. So as an example, we can talk about a carbon hydrogen bond, or a carbon hydrogen single bond. So this carbon is probably attached to some other stuff, because carbons usually have more than one single bond. But we're gonna ignore everything else attached to the carbon, we're just gonna represent it as a big blob, like popcorn, maybe it's a protein, it could be, it could be a sugar molecule, it could be a lot of things. But we're ignoring that blob. And one other thing I forgot to say earlier is that this is the energy it takes to break one mole of a bond in the gas phase. So it's a pretty specific definition. So in the case of our carbon hydrogen bond, the bond enthalpy of this bond, so if we break this bond... Let's do a sort of dotted line. If we break this bond... We have to add energy, and what we'll get as our products is we'll get our popcorn, and what happens is, when we break this bond, the two electrons that originally made up the bond, one of the electrons will go to the carbon, and the other electron'll go to the hydrogen. And we usually represent single electrons like that using a single dot, sort of like when you write Lewis structures, you can write lone pairs with two dots. So here's our carbon with one dot, or one electron, and our hydrogen with one electron, and these are both still in the gas phase. So the delta H of this reaction is the bond enthalpy, which I will abbreviate as BE. So some important things to remember about bond enthalpy are that bond enthalpy is always positive. So it's always going to take energy, you're always gonna have to add energy to break a bond. If we take the reverse of the bond, if we take the reverse of the bond enthalpy, so another way to think about this is to flip this reaction, so if we take the reverse of this reaction, that means we're making a bond. And since we know that breaking a bond always takes energy, that means making a bond always releases energy. So it will always be negative to make a bond. And that's another way of saying, it will always release energy. And then the third thing that we're gonna discuss about bond enthalpy is that you can use it, you can use bond enthalpy to estimate delta H of reaction. And delta H of reaction is, or the enthalpy of reaction, is something that chemists are often interested in. We wanna know if it's exothermic or endothermic. You might know that there's lots of other ways of calculating delta H of reaction, such as using Hess's law... Or another way is using delta H of formation. And then there are other ways too. So this is just another way that we can use to calculate delta H of reaction using bond enthalpies. So we're gonna go through an example of that next. So the example reaction is taking propyne, which is C3H4 gas... And reacting it with hydrogen, so hydrogen gas, to get propane, C3H8 gas. And I don't know about you, I'm pretty bad at looking at a chemical formula like this and knowing exactly what the molecule looks like, so I'm gonna draw out the Lewis structures. So the Lewis structure for propyne, propyne has three carbons, and one triple, one carbon carbon triple bond, and then it has four hydrogens. So that's propyne. And we also have hydrogen gas. And our product is propane, so propane has all single bonds. So three carbons with single bonds, and eight hydrogens bound to the carbons. So that's the reaction we are interested in, and what we wanna know here is what is delta H of reaction? And how can we calculate it using bond enthalpies? We said earlier that bond enthalpies, a bond enthalpy is the energy it takes to break a bond. So what we're gonna do next is look at our reaction in terms of what bonds are broken and what bonds are formed. And this is a lot easier to do using the Lewis structures. First let's talk about which bonds are broken. We, if we compare our reactants and our products, we're breaking this carbon carbon triple bond. If we're breaking this carbon carbon triple bond, and we're also gonna break this hydrogen hydrogen bond, and one thing we forgot to do earlier which is super important, is we actually need to make sure our reaction is balanced. And we have four plus two, six hydrogens on our reactant side, and we have eight hydrogens on our product side. That's not balanced. So we actually need two hydrogen molecules on the reactant side. So let's draw one more in. So yes, we said we are breaking a hydrogen hydrogen bond, we're actually breaking two hydrogen hydrogen bonds. It's important to keep track of how many of each type of bond we're breaking because the bond enthalpy is per mole, so if you have twice as many moles it'll take twice as much energy to break all of those bonds. And then we can look at the bonds that are formed. So we have, not, since we broke this carbon carbon triple bond, that means we needed to make a new bond, and the new bond we made in our product molecule is this carbon carbon single bond. Not only did we form a new single bond between these two carbons, but now these carbons are attached to a bunch of hydrogens, so we made four new carbon hydrogen bonds. So let's write that out so that we can keep track of them when we do our final calculation of delta H of reaction. So if we just look at the bonds broken... The bonds we broke... We have a carbon carbon triple bond, and we have a couple hydrogen hydrogen bonds. Let's also just write down how many of each we have, because we'll need that for our calculation. So we have one carbon carbon triple bond, and we have two hydrogen hydrogen bonds that are broken. And then we can also look up their bond enthalpies, which are in kilojoules per mole. Bond enthalpies you can typically look up in your textbook or online, and they usually come in a table of bond enthalpies. And so the units can be kilojoules per mole, sometimes you'll also see calories or kilocalories per mole. I already looked up these bond enthalpy values. So carbon carbon triple bonds have a bond enthalpy of 835 kilojoules per mole, and hydrogen hydrogen bonds have a bond enthalpy of 800, sorry, 436 kilojoules per mole. And then next, if we look at the bonds that are broken, we have a carbon carbon single bond. And we have one of those bonds forming. And the bond enthalpy for that, which is also in terms of kilojoules per mole, is 346. And last but not least, we have the carbon hydrogen bonds that we're forming, and we have four of those, and each of those, the bond enthalpy is 413 kilojoules per mole. So now we can take all of this information and put it together to calculate delta H of reaction. So delta H of reaction, if we're thinking about it in terms of bonds made and broken, it's a total energy change during a reaction. And so it's just the energy it takes to break all of our bonds in the reactants... So to break this carbon carbon triple bond and the two hydrogen hydrogen bonds, plus the energy it takes to make the bonds, to make new product bonds. We said earlier that you always have to add energy to break bonds, so bond enthalpy is always positive, so we know this part of our calculation should always be a positive number. What that means is that it always releases energy to make new bonds, and when energy is released, delta H becomes more negative. So this number here, when we're talking about adding up the energy it takes to make new bonds, these should be negative numbers. So now let's plug in the values we have for bond enthalpy for all of these bonds that are made and broken in our reaction. Let's start with the bonds that are broken. So we have our carbon carbon single bond, that will require 835 kilojoules per mole, and we have only one of them. And we also have to break two hydrogen hydrogen bonds, so two times 436 kilojoules per mole, which is the bond enthalpy of that bond. So that's all of the bonds we break. Now we have to add up the energy that's released when we make the new bonds. So we have this carbon carbon single bond. So that is 346 kilojoules per mole, and that's negative, because that energy is released. And then, the last bond is the carbon hydrogen bond, also negative because the energy is released, and we have four of them, and each of them will release 413 kilojoules per mole. So if we still all of this into our calculator to get our final answer, what I got was that the delta H of reaction for this, for this hydrogenation reaction between propine and hydrogen gas, is -291 kilojoules per mole. We can see that this overall reaction releases energy, because delta H is negative, so it's exothermic. And that's how you can use bond enthalpies to calculate delta H of reaction.