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Course: AP®︎/College Chemistry > Unit 6
Lesson 8: Bond enthalpiesWorked example: Using bond enthalpies to calculate enthalpy of reaction
Bond enthalpies can be used to estimate the change in enthalpy for a chemical reaction. In this video, we'll use average bond enthalpies to calculate the enthalpy change for the gas-phase combustion of ethanol. Created by Jay.
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- How graphite is more stable than a diamond rather than diamond liberate more amount of energy.(3 votes)
- 1:36He breaks all the bonds on the Ethanol, but wouldn't the oxygen be able to keep one of its bonds?(2 votes)
- When using bond enthalpy like this to estimate the enthalpy of a reaction, we’re assuming that the reaction proceeds by breaking all the reactant bonds.(3 votes)
- what do we mean by bond enthalpies of bonds formed or broken? Does it mean the amount of energies required to break or form bonds? If so how is a negative enthalpy indicate an exothermic reaction? According to my understanding, an exothermic reaction is the one in which energy is given off to the surrounding environment because the total energy of the products is less than the total energy of the reactants. If we scrutinise this statement: "the total energies of the products being less than the reactants", then a negative enthalpy cannot be an exothermic. In these eqauations, it can clearly be seen that the products have a higher energy than the reactants which means it's an endothermic because this violates the definition of an exothermic reaction.
I'm sure you, as an expert, are correct, but could you please clarify this? I'm always thinking of the same thing and cannot fathom why a negative enthalpy indicate an exothermic reaction and for this reason i'm always LOST.(1 vote)- The energy is the energy that the molecules have. If the molecules are losing energy, or have negative enthalpy, that energy has to go somewhere, which is seen as the release of energy to the surroundings in the form of heat.(2 votes)
- Why did the moles between the hydrogen-oxygen bond change from 6 to 5(1 vote)
- 615 for a carbon double bond breaking, plus an oxygen hydrogen bond is broken.
then a carbon oxygen bond is formed @ 351 KJ + a carbon hydrogen bond @ 413. a carbon carbon bond isnt formed. the answer should be
1078 -764
the question acted as if a new carbon bond was formed which was not the case. qeuestoin is wrong(1 vote) - can you use bond energies to calculate delta H knot tho cus doesn't the knot mean standard state so can u use bond enthalpies?(1 vote)
- I think you mean nought (British English) or naught (American English), instead of knot. Nought is another name for zero in English, a knot is something you tie with pieces of string.
But yes, you are correct ΔH° means the standard enthalpy change. Bond enthalpies are average bond dissociation energies for substances in the gaseous state. The two issues with bond enthalpies is that it is an average of bond energies, and also it applies only to gases.
A particular type of bond can have different bond enthalpies in different molecules. A good example is the carbon-hydrogen bond, C-H, which is commonly found in organic molecules. The C-H bond enthalpy in methane, CH4, is 438 kJ/mol. But in trifluoromethane CHF3 the C-H bond enthalpy is 446 kJ/mol. And there are similar differences for other C-H bond enthalpies in other molecules. The problem becomes which of the multitude of C-H bond enthalpies do we use for the overall C-H bond enthalpy? The solution is to use an average of a large number of compounds containing the C-H bond. That’s why the average C-H enthalpy is listed as 422 kJ/mol, even though it doesn’t correspond to any one particular bond. So trying to use bond enthalpy to find the standard enthalpy of a reaction yields on an estimate since the listed bond enthalpies do not correspond to the actual bond being investigated.
The other issue is that bond enthalpies apply only to substances in the gas state. However, not all substances are in the gas state as their standard state. So we can’t use bond enthalpies for substances in the liquid or solid phase directly. To get around this problem we have to add in extra enthalpy terms like heat of vaporizations (which is the enthalpy change going from liquid to gas) to account for the nongaseous states. This allows us to treat nongaseous substances as gases and use the bond enthalpies.
Hope that helps.(1 vote)
- Does anyone know how you tell there is a double bond versus a single bond between two molecules in a reaction?(1 vote)
- The best way is to draw an Lewis dot diagram to determine how many electrons each element needs. Then you can determine which have single and double bonds. Other than that, you can use periodic trends. For instance, when oxygen typically forms two covalent bonds, so when it is only sharing with one other element (like in CO2) it will form a double bond.(1 vote)
- So uhm so like uhm yeah you know like uhm that that I uhm so uhm that so I just want to know how to decide which bonds must be broken easily bro. Thank you bro.(0 votes)
- I mean, if we’re using bond enthalpies to determine the enthalpy of a reaction, we’re only breaking the reactant bonds.(1 vote)
Video transcript
- [Educator] Bond enthalpies can be used to estimate the standard
change in enthalpy for a chemical reaction. Let's use bond enthalpies to estimate the enthalpy of combustion of ethanol. Looking at our balanced equation, we have one mole of ethanol reacting with three moles of oxygen gas to produce two moles of carbon dioxide and three moles of water
in the gaseous state. To find the standard change in enthalpy for this chemical reaction, we need to sum the bond enthalpies of the bonds that are broken. And from that, we subtract the sum of the bond enthalpies of the bonds that are formed in this chemical reaction. To figure out which bonds are broken and which bonds are formed, it's helpful to look at the dot structures for our molecules. So let's start with the ethanol molecule. We're gonna approach this problem first like we're breaking all of
the bonds in these molecules. And we're also not gonna worry
about units until the end, just to save some space on the screen. So looking at the ethanol molecule, we would need to break
a carbon-carbon bond. So let's go ahead and
write this down here. Right now, we're summing
the bond enthalpies of the bonds that are broken. So we have one carbon-carbon bond. So let's write in here, the bond enthalpy for
a carbon-carbon bond. Next, we have five carbon-hydrogen bonds that we need to break. So to this, we're going to write in here, a five, and then the bond enthalpy of a carbon-hydrogen bond. Next, we have to break a
carbon-oxygen single bond. So we write a one, and then the bond enthalpy for a carbon-oxygen single bond. And then for this ethanol molecule, we also have an
oxygen-hydrogen single bond. So we'll write in here, a one, and the bond enthalpy for an oxygen-hydrogen single bond. We saw in the balanced equation that one mole of ethanol reacts with three moles of oxygen gas. So to represent the three
moles of oxygen gas, I've drawn in here, three molecules of O2. And we can see in each molecule of O2, there's an oxygen-oxygen double bond. So we would need to break three
oxygen-oxygen double bonds. So to this, we're going to add a three
times the bond enthalpy of an oxygen-oxygen double bond. And this now gives us the
sum of the bond enthalpies for all the bonds that need to be broken. It takes energy to break a bond. So the summation of the bond enthalpies of the bonds that are broken is going to be a positive value. And since it takes energy to break bonds, energy is given off when bonds form. So next, we're gonna
sum the bond enthalpies of the bonds that are formed. And notice we have this
negative sign in here because this energy is given off. So we're gonna write a minus sign in here, and then we're gonna put some brackets because next we're going
to sum the bond enthalpies of the bonds that are formed. In our balanced equation, we formed two moles of carbon dioxide. So to represent those two moles, I've drawn in here, two molecules of CO2. And we can see that in
each molecule of CO2, we're going to form two
carbon-oxygen double bonds. So that's a total of four
carbon-oxygen double bonds. So down here, we're going to write a four
times the bond enthalpy of a carbon-oxygen double bond. We also formed three moles of H2O. And in each molecule of
water that's drawn here, we form two oxygen-hydrogen single bonds. And since we have three moles, we have a total of six
oxygen-hydrogen single bonds. So to this, we're going to add six
times the bond enthalpy of an oxygen-hydrogen single bond. The next step is to look
up the bond enthalpies of all of these different bonds. For example, the bond enthalpy for a carbon-carbon single
bond is about 348 kilojoules per mole. You might see a different value, if you look in a different textbook. However, we're gonna go
with 348 kilojoules per mole for our calculation. And we're gonna multiply this by one mole of carbon-carbon single bonds. Next, we look up the bond enthalpy for our carbon-hydrogen single bond. And that's about 413 kilojoules per mole of carbon-hydrogen bonds. And we're multiplying this by five. And we continue with everything else for the summation of
the the bond enthalpies of the bonds broken. When we do this, we get positive 4,719 kilojoules. Next, we do the same thing for the bond enthalpies of the bonds that are formed. So the bond enthalpy for our carbon-oxygen double
bond is 799 kilojoules per mole, and we multiply that by four. The bonds enthalpy for an
oxygen hydrogen single bond is 463 kilojoules per mole, and we multiply that by six. When we add these together, we get 5,974. So for the final standard
change in enthalpy for our chemical reaction, it's positive 4,719 minus 5,974, which gives us negative 1,255 kilojoules. Notice that we got a negative value for the change in enthalpy. And that means the combustion of ethanol is an exothermic reaction. And 1,255 kilojoules
of energy are given off for the combustion of one mole of ethanol. Also notice that the sum
of the bond enthalpies of the bonds formed, which is 5,974, is greater than the sum
of the bond enthalpies of the bonds broken, which is 4,719. And since we're
subtracting a larger number from a smaller number, we get that negative sign for the change in enthalpy. If the sum of the bond enthalpies of the bonds that are broken, if this number is larger than the sum of the bond enthalpies of the bonds that have formed, we would've gotten a positive value for the change in enthalpy. And that would be true for
an endothermic reaction. We did this problem, assuming that all of the bonds that we drew in our dots
structures were broken and all of the bonds that we drew in the dot
structures were formed. However, if we look
closely to dots structures or just look closely
to what we wrote here, we show breaking one oxygen-hydrogen
single bonds over here, and we show the formation of six oxygen-hydrogen
single bonds over here. So we could have just canceled out one of those oxygen-hydrogen single bonds. So we could have canceled this out. And instead of showing a six here, we could have written a
five times the bond enthalpy of an oxygen-hydrogen single bond. We still would have ended
up with the same answer of negative 1,255 kilojoules. So if you look at your dot structures, if you see a bond that's the
same on the reactant side and the same on the product side, you don't have to show the breaking and forming of that bond. You can make the problem
a little bit shorter, if you want to. Finally, let's show how we get our units. So to get kilojoules as your final answer, if we go back up to here, we wrote a one times 348. The one is referring to breaking one mole of carbon-carbon single bonds. And the 348, of course, is the bond enthalpy for a carbon-carbon single bond. So this was 348 kilojoules per one mole of carbon-carbon single bonds. When you multiply these two together, the moles of carbon-carbon
single bonds cancels and this gives you 348 kilojoules. And so, that's how to end up with kilojoules as your final answer. You also might see kilojoules
per mole of reaction as the units for this. And, kilojoules per mole reaction means how the reaction is written. So for the combustion of one mole of ethanol, 1,255 kilojoules of energy are released. To get kilojoules per mole
of reaction as our units, the balanced equation had
a one as the coefficient in front of ethanol. Therefore, you're breaking one mole of carbon-carbon single bonds per one mole of reaction. So we can use this conversion factor. Now, when we multiply through the moles of carbon-carbon single bonds, cancel and this gives us
348 kilojoules per mole of reaction.