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Bond enthalpy

Watch this to explore the energy of bond enthalpy for Higher Chemistry (and why it is crucial for marshmallows)

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I have many good memories of sitting around a fire with a group of friends. The sky is clear, and the stars are twinkling. The flames are dancing, and I can feel the heat warming my body. There have been stories told, songs performed, and marshmallows melted – a particular favourite with my children.

There’s something primal and meditative about watching and feeling this energy from the fire. It warms the skin and drives deep thought.

But how is it happening? Why are we able to experience this energy? Why does a chemical reaction make this feeling so intense? Let’s find out…

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Early on in chemistry you would have come across the terms exothermic, where heat energy is released, and endothermic, where heat energy is taken in. When wood burns in a fire, there is a combustion reaction, and heat is released to the surroundings, an exothermic reaction.

Let’s think about what is happening with the bonds in the molecules involved. Well use methane – the gas we burn in the lab as a simple example to work on.

Methane contains one carbon bonded to four hydrogens. Each of these bonds has an energy associated with it, also known as the bond enthalpy. As you’ll be aware, carbon to hydrogen bonds are very common and appear in all sorts of compounds. This means that the energy of this bond can vary depending on the environment it is in. Therefore, we use a value called the mean bond enthalpy – the average enthalpy for the C-H bond. This can be found in your data book with the value 412 kJ/mol.

The energy required to break a methane molecule into gaseous carbon and hydrogen atoms can be calculated by multiplying 412 by 4 giving us 1648 kJ/mol. As this value was determined using the mean bond enthalpy, it will be an estimate of the energy required.

For some molecules we can obtain the exact energy required to break one mole of the bond into separate atoms. For example, diatomic molecules contain bonds that can only be in one environment. For O2, the bond enthalpy is 498 kJ/mol – again this can be found in your data book.

Using bond enthalpies, we can determine the enthalpy change for a gas phase reaction. When combusting methane, two moles of oxygen are required, and will form one mole of carbon dioxide and two moles of water. This is important as we need to include this in the enthalpy change calculation. In total, 2644 kJ/mol is required to break up our reactants.

Using the same process, the total enthalpy of the products can be calculated, just ensure that you include all the bonds involved. From the values in the data book, the total enthalpy is 3338 kJ/mol. It’s important to note here that when the products are formed, energy is released. To show this with the enthalpy value, we make it negative.

To determine the enthalpy change, add the values together, 2644 + (-3338) gives -694 kJ/mol. The result is negative and so this tells us that it is an exothermic reaction.

In the data book you will also find the value for the enthalpy of combustion of methane, which is when one mole of methane burns completely in oxygen. This is the same reaction that we have been referring to, but the value calculated is different. This is due to it being a more accurate determination of the enthalpy change rather than using mean bond enthalpies.

If you find yourself sitting around a fire with friends and you get stuck for a fireside story, then tell them about enthalpy change and appreciate how this exothermic reaction allows you to warm up when it is cold and, perhaps more importantly, to melt marshmallows.

You never know, they might just be interested.

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