Download Chem 30 – Thermochemistry

In Chemistry 30 - Thermochemistry we talk about 2 types of processes:
Endothermic Processes (those that take heat in)
Such as boiling water OR dissolving ammonium nitrate in water (this causes the water to cool immediately in a cold
pack), etc.
and
Exothermic Reactions (those that release heat)
Such as condensing steam (water gas turns into liquid) OR the explosion of dynamite, etc.
Endothermic Reactions can be represented as
A+B →C
A + B + heat → C
ΔH = +ve (enthalpy is positive)
Where the difference in energy (essentially heat and pressure/volume changes) is represented as ΔH
We can graphically show an endothermic process - the energy inside product C is greater than that of A and B:
C is at a higher energy than A and B - this means that the products have more potential energy than the reactants. This
also means C is a less stable compound.
Take a look at pages 4-5 of the Chemistry 30 Databook – this is a table of the Standard Molar Enthalpies of Formation.
This information tells you the amount of energy released when various compounds are formed from their elements.
Notice that the majority of the values are negative – meaning they are exothermic and forming the compounds releases
energy. The universe favors lower energy and greater stability.
Endothermic reactions are contrary to this – they create less stable compounds – one example of which is the synthesis
of explosives.
Endothermic processes include the synthesis of trinitrotoluene (TNT). TNT is made by repeatedly nitrating (adding a
nitro group) to toluene (benzene with one methyl on it). This process requires energy and has been performed on stove
tops – however you’re making an explosive using HEAT. As you can probably tell a lot of people have hurt themselves by
making some unstable endothermic compounds – they require heat to make, but too hot and they are likely to blow up
in the process of manufacture. It’s not a smart idea to try at home.
Exothermic Reactions can be represented as
A+B →C
A + B → C + heat
ΔH = −ve (enthalpy is negative)
Graphically, we can show the energy inside product C is lower than that of A and B:
Examples of endothermic processes are
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the boiling of water, water(liquid) → water(gas) ∆H = +ve (positive)
the dissolving of ammonium nitrate (which is used in cold packs)
Notice that whenever we have a reaction there is a bit of extra energy required to get it going – it’s the ‘bump’ in going
from A and B to C. That’s called the activation energy. The best way to think about it is this – when you take a match out
of your pocket, does it just spontaneously light on its own?
What do you need to do to light it?
That’s right – you need to strike it. It needs to form something unstable before it can decompose in an exothermic
reaction.
So – how is this information used? First, you must be able to identify endothermic and exothermic processes.
Let’s do an exercise – identify the following as either endothermic or exothermic when looking at the chemical/physical
process:
1.
2.
3.
4.
5.
A plant forms carbon dioxide and water into sugar using photosynthesis.
An athlete pours water on himself during a hot day and feels cooler.
Water freezes into ice during winter.
A small pond dries up during a hot summer.
Sodium hydroxide is added to hydrochloric acid in water and the temperature of the water increases.
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Solutions:
1. A plant forms carbon dioxide and water into sugar using photosynthesis. This is endothermic because
the plant takes in energy in the form of sunlight during for this process.
2. An athlete pours water on herself during a hot day and feels cooler. This is an endothermic process – the
water on the athletes skin absorbs heat and evaporates, causing the athlete to feel cooler.
3. Water freezes into ice during winter. This is an exothermic process – the water goes from liquid to ice
(high energy to low energy) and hence its surroundings are warmed. This is actually quite cool because
in the winter time orange trees are watered to both warm them up and protect the fruit in an insulated
igloo of ice. Actual damage to the fruit occurs when temperature fluctuates and the ice freezes/thaws
repeatedly.
4. A small pond dries up during a hot summer. This is an endothermic process – the water absorbed heat
to evaporate.
5. Sodium hydroxide is added to hydrochloric acid in water and the temperature of the water increases.
This is an exothermic reaction that released heat causing the water to warm up.
Once we know the energy change of a process we can manipulate it. For instance:
We can say that
A → B ΔH = −100 kJ
2A → 2B ΔH = −200 kJ
We saw that by doubling the amount of reagent we doubled the quantity of energy exchanged (in this case produced).
The analogy is you have a stick of dynamite, and it blows up making 100 kJ of heat. If you have 2 sticks of dynamite you’d
expect the explosion to be twice as large and generate 200 kJ of heat.
Now let’s say the reaction is reversible. First, before speaking of reversible reactions, let’s define an irreversible one:
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*Picture above is composed of images from royalty free archives of Microsoft Office.
If you make an egg into an omelet, can you reverse the reaction and make that omelet back into a chicken? No, you
CANNOT reverse the denaturing of the protein (amongst other processes) in the egg. the process is irreversible
regardless of whether your original egg is insane or not (as pictured).
Now as far as reversible processes go you already know a couple – you can vaporize water and then condense it back
into a liquid:
*Picture above is composed of images from royalty free archives of Microsoft Office.
Let’s see how this relates to our system:
If we flip the relation, we get:
A → B ΔH = −100 kJ
B → A ΔH = +100 kJ
This makes sense as we’ve just flipped the direction of our reaction coordinate:
If we have half as much material we get the following relation:
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1
1
B → A ΔH = +50 kJ
2
2
To summarize, whatever factor you multiply the coefficients by, you multiply the enthalpy by:
A → B ΔH = −100 kJ
2A → 2B ΔH = −200 kJ
3A → 3B ΔH = −300 kJ
1
1
A → B ΔH = −50 kJ
2
2
1
1
B → A ΔH = +50 kJ
2
2
B → A ΔH = +100 kJ
This concept will be continued in a future article: Chemistry 30 – Thermochemistry – Enthalpy and Hess’s Law, stay
tuned!
© Copyright Pavel Sedach 2013, Page 5 of 5