COMBUSTION AND ENTHAPHY

What are my objectives while studying this topic?


By studying this topic you’ll learn how to measure the bond energetics of fuel molecules.


Why is this important?


We know that fuels store and release energy so by studying and measuring the bond energetics of fuel molecules it’s possible to work out how much heat energy is released when a fuel burns. We can use this information to generate heat, light and cook food. What’s even more important is our ability to create electricity by using this knowledge. This is the mainstay of the Developed World, and our lives as we know them.


Can I get an introduction to this topic?


Yes, why not have a look at the following clip:



So, let’s get to it: what is combustion and enthalpy?


Fire has always been special for human beings. On one hand it is very dangerous and potentially lethal, but on the other it is the giver of warmth and light.

In order for things or materials to burn the fire triangle is essential.


Fire Triangle 2

 

The triangle illustrates the rule that in order to ignite and burn, a fire requires three elements: heat, fuel, and an oxidizing agent, usually oxygen. A fire naturally occurs when the elements are combined in the right mixture. The fire is prevented or extinguished by removing any one of them.


Heat


Without sufficient heat, a fire cannot begin, and it cannot continue. Heat can be removed by dousing with water; the water turns to steam and the steam is further heated, taking the heat with it. Introducing particles of powder or any gas in the flame removes heat in the same manner.


Fuel


Without fuel, a fire will stop. Fuel can be removed naturally, as where the fire has consumed all the burnable fuel, or manually, by mechanically or chemically removing the fuel from the fire.


Oxygen


Without sufficient oxygen, a fire cannot begin, and it cannot continue. With a decreased oxygen concentration, the combustion velocity gets lower. The heat content of things is called ‘enthalpy’.


Where can I see the fire triangle in action?



By using the fire traingle, you can put out fires. Take a look at the following link to see how:



What is the definition of enthalpy?


The standard molar enthalpy change of bond dissociation ( ΔHd° is the energy change when 1 mole of bonds is broken, the molecules and resulting fragments being in the gaseous state at 298K and a pressure of 100kPa).


This energy refers to a specific bond in a molecule, but if a molecule has 4 of the same bond (eg, the C-H bonds in methane), then different dissociation energies can occur.


CH4(g) => CH3(g) + H(g) _______ Hd° = +427 kJ mol-1


CH3(g) => CH2(g) + H(g) _______ Hd° = +371 kJ mol-1


So it is much more useful to know the average amount of energy needed to break a particular bond. In this case the process of breaking all the bonds in methane ending up with gaseous atoms.


So this process could be written as:


CH4(s) --------> C(g) + 4H(g)


The enthalpy change for this reaction is:

+1646 kJ mol-1

So the average bond enthalpy is:

+1646 / 4 = +412 kJ mol-1

They can be looked up in data tables.


It is important to stress that these are mean or average bond enthalpies. If average bond enthalpies are used to calculate an enthalpy change, the answer will be slightly out compared to a result obtained by other methods.


Why is this useful?


Their main use is in working out enthalpy changes for reactions. If we know the amount of energy needed to break a bond (endothermic), and the amount of energy we get back when a new bond forms (exothermic), then we can quite easily work out an approximate (because of the average nature of the bond enthalpies) value for the enthalpy change!


Remember: Breaking bonds requires energy (Endothermic) while making bonds releases energy (Exothermic).


Don't be scared by this equation below; it is just the correct terminology as used by chemists for working out enthalpy changes from bond energies.


Equation 1

This basically means that you add up all the energies of the bonds that are reformed and subtract one from the other. It’s another version of Hess’s law. Similar to the one that can be used when you know all the enthalpies of formation for the substances in a reaction.

An example:


The complete combustion of propane can be represented by the following equation:


Equation 2

Or we could redraw it to represent the bonds present:


Equation 3

We now need to work out how many of each bond type we have broken.


  • 8xC-H

  • 2xC-C

  • 5xO=O


And then how many bonds have been formed!


  • 6xC=O

  • 8xH-O

So using data tables we can look up then average bond enthalpies from, and calculate the enthalpy change of the reaction.


Bond Type

Average bond enthalpy /kJ mol-1

C-H

+413

C-C

+347

O=O

+498

C=O

+805

H-O

+464


Notice they are all endothermic.


So we can now do the sum, remember, sum of bonds broken - sum of bonds formed.

 

Hr°= [(8x413)+(2x347)+(5x498)] - [(6x805)+(8x464)]



= - 2054 kJ mol-1


The value for the enthalpy of combustion of propane from the data table is -2219kJ mol-1.


This apparent error is due to fact that we use Average Bond Enthalpies in our calculations and the fact that the values above relate to the gaseous state, while the standard combustion state of water is liquid. If we allow for this we get a value of -2226 kJ mol-1 which is pretty close to the value obtained above. The remaining difference must be down to the average bond enthalpy factor.


Useful questions to ask!


Look at the following links to base a class discussion on. Think about how combustion played a part in the burning of the fuels in the following fires.




What are the points to grasp?


Energy changes during combustion involve chemical reactions and are related to the making and breaking of chemical bonds.


Enthalpy, in terms of the total energy content of a system at constant pressure, is represented by the letter H.


The standard molar enthalpy of bond dissociation is the enthalpy change when one mole of bonds of a particular type are broken under standard conditions.


What are the key words to learn and use?


  • Combustion

  • Bonds

  • Enthalpy

  • Standard

  • Dissociation

  • Exothermic

  • Endothermic


What will I be expected to be able do after studying this topic?


Use enthalpy diagrams to represent the energy changes during combustion reactions.


Enthalphy Diagram

Use tables of bond enthalpy to calculate energy changes during combustion.


Energy Changes

Write balanced chemical equations for the combustion of carbon-based (fossil) fuels.


How long will this topic take to study?


On average you’ll be looking at about six hours of study time.


What are the main activities on this topic?


Experiment!


You’ll watch a demonstration showing the combustion of one of the following: crude oil, crushed coal or natural gas. (Make sure you follow the CLEAPSS Safety advice).


Questions:


What are the products of combustion – you should make observations and deductions (and know and perform the chemical tests for water and carbon dioxide).


Activity


Word & balanced symbol equations for fuel combustion.


Activity


Make Molymod models of combustion of fuels in oxygen – illustrating the breaking and making of bonds.


Activity


Define enthalpy and standard molar enthalpy of bond dissociation and formation. Define standard conditions. The following link will help you here:


Activity


Use enthalpy level diagrams to show enthalpy changes in reactions (exothermic and endothermic examples).


Activity


Use bond enthalpy tables to draw enthalpy level diagrams.


Experiment!


Measure the temperature change during reactions (both exothermic and endothermic) eg, displacement of copper by zinc from copper sulphate (exothermic) and the reaction of ammonium nitrate with water (endothermic).


Points:


  • You could use a datalogger to measure temperature with time during reactions

  • You should use and complete standard risk assessments using suitable CLEAPSS guidance (including Hazcards)

  • You will need to record your data and plot it graphically and calculate the temperature changes

  • You should evaluate your experiments. (eg, ILPAC Advanced Practical Chemistry, 2nd edition, Experiments 2.1 and 2.2)


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