Energetics - Bond Enthalpy and measuring energy changes

This topic will be split up into seperate posts because of it's length. This post covers syllabus statements 5.1 and 5.3

What is energetics?

Energetics is basically the study of energy in a reaction I guess, the amount of energy required or produced by reactions and more. It's quite a vital part of chemistry and links to almost every other topic in IB Chemistry.

The BAsics

Just a couple of basic things that you should know before you even begin this topic.

• Temperature is the measure of the average kinetic energy or all the particles/molecules in whatever you're measuring the temperature of

There are of course several different temperature measures. Never use Fahrenheit, it's just about the dumbest measure of temperature ever invented (sorry America). In Energetics we mainly use Kelvin because it means that we never get a negative temperature, and it's directly proportional to the average kinetic energy of the particles unlike Fahrenheit.

• 0 K is -273 C (never forget)
• Energy is measured in Joules
  

Energies are usually measured in Joules per Mole in chemistry, just because it's great to know how much energy is produced or needed per mole. Why per mole? because then we can directly relate it to other compounds or reactions without needing to do any conversions.

Awesome, lets begin!

What is enthalpy?

Enthalpy is formally defined as being the energy content of a system (a reaction) but basically just means 'energy'. It's mostly used to talk about how much energy is used or absorbed in reactions, which can tell us a lot about the reaction itself.

Enthalpy change

Enthalpy change should already be known to you if you've done GCSE Chemistry. Some reactions produce a lot of heat (think bombs) and some require energy input to react like photosynthesis (the energy is light) or the chemical reactions that occur in instant ice packs. 

• Reactions that release energy, usually in the form of heat, are called EXOTHERMIC reactions
  • ​​EXOTHERMIC reactions have a negative ΔH value because the reaction loses energy
• Reactions that absorb, use or  require energy are called ENDOTHERMIC reactions
  • ​ENDOTHERMIC reactions have a positive ΔH value because the reaction gains energy

Enthalpy change diagrams

Hopefully this part should be relatively straightforward. Endothermic reactions absorb energy from the surroundings (or system). Why? Because these reactions require more energy to break the bonds of the reactants then they release by forming bonds as the products. The reactants in an endothermic reaction therefore have (or had) a higher bond energy than in the products, and are more energetically stable. Exothermic reactions are basically the complete opposite. They release energy into the surroundings (or system) because they release more energy forming the bonds of the products then they absorb to break the bonds of the reactants. Their products are more energetically stable than their reactants.

The theoretical 'hill' a reaction has to climb as represented on this diagram is the activation energy.

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Why do exothermic reactions release heat and endothermic reactions absorb heat?

It's all to do with bonds.

• When bonds are made, energy is released
• When bonds are broken, energy is absorbed because energy is required to break those bonds​​

​If more energy is absorbed to break bonds than energy is released from making bonds, then the reaction is exothermic and vice versa.

How do we calculate the enthalpy change of a reaction in an exam?

That's when you get your data booklet out and flip to table 11. There you can find the bond enthalpies of many common bonds. Then you take a look at the reaction that you're dealing with and figure out exactly what bonds are broken and formed. Bonds that are broken are added together, and bonds that are formed are subtracted.

All the bond enthalpy values in this table are in kilojoules per mole. These values are the average amount of energy needed to break 1 mole of the bonds in question in a gaseous state.

Why average? because every different molecule will have a different bond enthalpy for their bonds. For example the enthalpy of the C-C bond in ethane will be slightly different to the enthalpy of the C-C bond in propane.

An example of a calculation:

Let's use the simple example of CH4 + 2O2 --> CO2 + 2H2O
First, all the bonds of the reactants need to be broken. We need to break 4 C-H bonds and 2 O=O bonds. If you check the data booklet, a C-H single bond is 414 kJ/mol. We're going to break 4 of these so we need to do
414 x 4 = 1656 kJ/mol to break all the bonds in 1 mole of CH4 molecules.
Then we need to break 2 O=O double bonds. They have an enthalpy of 498 kJ/mol.
498 x 2 = 996 kJ/mol to break all the bonds in 2 moles of O2 molecules.
We then add these 2 values together to get 
996 + 1656 = 2652 kJ/mol to break all the bonds of the reactants (1 mol CH4 and 2 mol O2)

Then we need to find the total enthalpy of all the bonds formed. To form CO2 we need to form 2 C=O double bonds. C=O double bonds have an enthalpy of 804 kJ/mol.
804 x 2 = 1608 kJ/mol to form all the bonds in 1 mole of CO2 molecules.
Then we need to do the same for H2O. We need to form 2 H-O single bonds to make water, but in our equation we have 2H2O, so we actually need to form 4. The enthalpy of the H-O single bond is 463 kJ/mol.
4 x 463 = 1852 kJ/mol to form all the bonds in 2 mole of H2O molecules.
We add these 2 values together to get the total energy released in forming all the bonds in our reaction
1608 + 1852 = 3460 kJ/mol to form all the bonds of the products.

To get the enthalpy change the entire reaction, we need to subtract the energy released in forming the bonds of the reactions from the energy absorbed in breaking the bonds of the reactants:
Enthalpy of bonds in the reactants - Enthalpy of bonds in the products
2652 - 3460 = -808 kJ/mol of CH4.
From the sign on the  enthalpy change, you can see that this reaction is EXOTHERMIC​.

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What if i want to find the enthalpy change in a real life reaction?

​The enthalpy change can be calculated using a formula you've probably heard of before and are probably bored of, Q = mcΔt. ​

I've written that the change in temperature here is in degrees celcius but it can actually be kelvin too. Seeing as the scale is exactly the same, it shouldn't matter.

This formula allows you to calculate the amount of energy absorbed or released in a reaction just from the temperature change, the specific heat capacity, and the mass of the solution you're dealing with.

What's the specific heat capacity?

It's the amount of energy you need to put into a unit mass to increase the temperature of it by 1 Kelvin. In chemistry (unlike physics) we work in grams, so the specific heat capacity will always be in Joules per gram. If your reaction is happening in solution, which it usually will be, then you use the specific heat capacity of water which is 4.2 J/gK.

Even though this formula is quite simple, it's easy to slip up on questions about it, here's why

Let me give you an example of a calculation. Let's say we have a simple reaction where the mass of the solution is 100g and the temperature rises by 6 C or 6K. They then ask you for the enthalpy change in kJ. Simply enough, you just put everything into the formula.
Q = 100 x 4.2 x 6
    =2,520 Joules.
Great, so then what people do is they put the enthalpy change of the reaction down as 2,520 Joules. This is WRONG!!! Look at your temperature change. The temperature has increased, meaning it must be an EXOTHERMIC reaction, but you've put the value down as a positive number, which would be the enthalpy change for an endothermic reaction!

• Always remember to reverse the sign when doing Q=mcΔt calculations, or just double check that the sign is correct for the temperature change that has been measured
  • ​An increase in temperature (positive Δt) means a negative ΔH value!
    
  • A decrease in temperature (negative Δt) means a positive ΔH value!
    

Ok, so the answer is -2520 Joules then?
No!!! It asked you for kJ! You need to divide the answer by 1000 to get the kJ!
2520/1000 = 2.52 kJ

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Enthalpies of combustion

How can we find out the energy content of a fuel? It's actually quite simple and it uses a process called calorimetry, not to be confused with colourimetry.

The heck is calorimetry? Sounds like something to do with dieting

Don't worry it definitely has nothing to do with diets (thank god). It uses the same sort of concepts we used earlier to find the enthalpy change of a reaction.
​Here's the setup:

Simple. Using this apparatus, you can pretty easily measure the enthalpy of combustion of the fuel if you know the temperature change and the mass of the fuel before and after burning. Then, you simply put it all into
​Q=mcΔt.

The mass in this case, is the mass of the water that was inside the calorimeter, and the specific heat capacity is 4.2 J/gK. Of course, the Δt is the change in temperature burning.

Example time:

Let's say we burnt propane. The original mass was 150g and after burning we had 148g. The temperature of the 100g of water in the calorimeter increased by 15 degrees.
Q=mc​Δt
   =100 x 4.2 x 15
   = 6300 J
This is how many joules were released when we burnt 2 grams (150  148) of propane. If we want to find the enthalpy of combustion in joules per gram, we simply divide by 2
6300/2 = 3150 J/g
If we want the enthalpy of combustion in joules per mole however, we need to find the number of moles of propane in 2 grams.
moles = mass/RFM
           = 2 / 44
           = 0.0454545..... moles
We then divide the joules released when 2g was burnt by the number of moles or propane in the 2g
6300 / 0.0454545.... = 138,600 Joules per Mole

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Hope this helped!

Let me know if you spot any mistakes or anything needs clarifying!

Over and out :)