Peroxide bubbles…a lot…
In the experiment where a few drops of an Iron(III) chloride solution is added to a
test tube of peroxide, sometimes a key observation may get overlooked.
When this part of the lab is done near the end of the period (some hustle and bustle)
I should take some responsibility for any missed observations. So here I’d like to try
and explain what some may not have noticed…and subsequently had some
challenges answering a question on the lab assignment
So what is this ‘key’ observation?
Good question, I’m so glad you asked, because I made this cool diagram with test
tubes with a 3D effect and…umm…OK, that’s probably enough about me…
You’re all probably familiar with peroxide, since it may have been in your bathroom
cupboard, ready to be poured on any wounds that you might have got as a child. It
makes a good antiseptic, and if you've had the pleasure of having peroxide used on
you, you might have noticed that it generated lots of bubbles when it contacted your
open wound…ehhh...trying not to picture a gaping wound here…damn it, that’s all I
can see now!!
…OK…recovering a bit…thinking of nice green meadows, with bunny rabbits and
butterflies…phew, I’m good. Back to the experiment…
Peroxide is somewhat stable as H2O2 in dilute aqueous solutions, and can stay
stable for a while (a month or two after opening), but breaks down slowly to oxygen
and water, as shown in the equation below:
Both the H2O2 and the H2O are liquids, so if bubbles form from this breakdown, the
O2 must be the source…being the only gas in the reaction.
So the questions that might follow from this would be:
1. Why did we see bubbles after we added the iron(III) chloride to our peroxide
solution?
2. Why does the peroxide bubble when it contacts our tissues?
Let’s deal with the first question
Below is a picture that shows the three stages of the reaction we observed, after the
addition of the iron(III) chloride
When we added the iron(III) chloride, the solution turned a reddish colour and
immediately started to bubble…a lot…which is what is happening in Tube A.
If we compare this tube to one with just some iron(III) chloride solution in water, we
might notice that the reddish colour starts to fade in the peroxide tube, to a yellowy
green, while it was still bubbling away - which is shown in Tube B.
Unfortunately, sometimes we are pressed for time and the sample gets quite hot, so
you may dump the tube contents into the waste container before the bubbling stops
completely. If we were to wait, we would likely see the tube return to the beginning
reddish colour, as if nothing had ever happened. (matching the tube with just iron(III)
chloride - which is shown in Tube C).
Obviously following along, something chemically had to have happened; heat and
gas were evolved, a colour changed - all proof that a chemical reaction(s) had taken
place.
Since we were following the decomposition of peroxide, the chemical equation
above (Eq. 1), showing the formation of oxygen gas must be the explanation for the
bubbles we observed. I would also guess that because of the speed of bubble
formation, this decomposition is the likeliest source for the heat that was generated.
But that darn colour change from reddish to yellowy and back to reddish? That has
to mean something.
…and it does.
We are dealing with a catalyst. This is a reactant that helps get a reaction going by
making the process easier; is only temporarily involved; and returns to its original
self so that it can keep helping the reaction to go forward. This is the definition of a
catalyst.
But how do we know that this happened in the reaction here?
Well, firstly the peroxide didn’t start bubbling until the iron(III) chloride was added,
so the iron must be involved. Since the peroxide, water and oxygen are all
colourless, and the iron(III) chloride solution was the source of the original reddish
colour, it’s probably not a stretch to assume that the iron is also the reason the
colour changes to yellow green…ish…
Below you can see the chemical reaction equations that show the changes in iron
leading to these colour changes, and how the original state of iron is regenerated.
If you look at Eq. 2 you can see that as the Fe3+ reacts with the peroxide to convert
it to oxygen, in that process the iron turns into its other cationic form - Fe2+ (which is
yellow). As this happens, the Fe2+ reacts with another peroxide (Eq. 3) and forms
the water, converting that iron cationic form back to Fe3+, which starts the cycle all
over again.
This can go on as long as there is peroxide, and when that runs out, the colour will
eventually change back to the Fe3+. This can happen because Fe3+ is the more
stable of the ions of iron, and there is plenty of oxygen around to help with that
(remember - red rust is iron(III) oxide).
What is interesting of course, is that if you add Eq. 2 and Eq. 3 together, and
subtract everything that is on both sides of the arrow…you get Eq. 1. Just as if the
iron wasn’t even there.
Why? Because the iron is just a catalyst, and is just there to make the reaction go
faster…it’s not actually a part of the net reaction….kind of…
Catalysts are important in chemistry and we’ll talk more about them in section 7.8,
but the real importance of their function is in biological enzymes.
In enzymic reactions (which most biological process are), the enzyme is the
facilitator/catalyst; designed to make reactions go faster, and make them more likely
to happen. Which of course brings us to the bubbling wound of a peroxide
antiseptic treatment.
Peroxide is a very concentrated form of reactive oxygen, and too much reactive
oxygen is not good…especially for bacteria. Because oxygen has such high
electronegativity, i.e electron greediness (remember oxygen is second only to
fluorine in that respect), too much of a good thing is not great for our cells either. To
combat that, we have an enzyme in our cells which is designed to deal with the
presence of peroxide. It’s called Catalase, and it quickly converts peroxide to
oxygen, which is much less reactive then peroxide. What’s also interesting is that
while catalase is a large protein - it has 4 active sites (kind of like haemoglobin) that
all contain iron Fe3+ ions. You can imagine that nature has invented a chemical
machine to do the same thing that our iron(III) chloride did, but with more of a
controlled biological process than our chemical one.
So when the peroxide was poured over that banged up knee wound we suffered
from falling off our bike, the catalase in our cells quickly dealt with it, producing the
bubbly oxygen molecules.
Unfortunately (for them), some bacteria aren’t so lucky to have such an enzyme, so
the peroxide can quickly react with their cell components and hopefully kill them in
the process.
…meaning, we could go on with our germ and infection free day…in that green
meadow, filled with bunny rabbits and butterflies…