Hydrogen peroxide (H2O2) is everywhere. It’s in your bleach, your hair dye, your cleaning cabinet and even your teeth whiteners. What you may not know is that this chemical is decomposing all the time. This happens very slowly but, with the magic of science, we can speed it up.
In this post:
Why Does Hydrogen Peroxide Decompose?
When it comes to determining exactly why hydrogen peroxide decomposes so easily, we have to look at the chemical structure of the H2O2 molecule.
Hydrogen peroxide contains a single oxygen-oxygen bond. Otherwise known as a peroxide bond, this is incredibly weak and unstable.
When its oxygen-oxygen bond breaks, hydrogen peroxide decomposes into water and oxygen. When this happens, it releases free radicals that are highly reactive with other substances.
While this decomposition reaction can be sped up by a catalyst, the instability of the peroxide bond means that decomposition also occurs naturally.
Natural Decomposition of Hydrogen Peroxide
From cosmetic to industrial applications, hydrogen peroxide is used for a variety of things. But there is always one thing these industries have in common: how hydrogen peroxide is stored.
Hydrogen peroxide has a finite shelf-life because, over time, it naturally decomposes into water and oxygen gas. Although this will take a while, UV rays from sunlight as well as warm conditions can actually catalyse the decomposition reaction.
This is why hydrogen peroxide is generally stored in dark plastic containers. The opaque colour protects the chemical from sunlight, while the plastic material accommodates for any build-up of oxygen gas that may occur.
A glass container, for example, has the potential to shatter if there is an increase in pressure. This is also why vented caps are fitted to hydrogen peroxide containers, as these provide an escape for any evolved oxygen.
Speeding up the Reaction
Sunlight isn’t the only thing that can speed up the decomposition reaction in hydrogen peroxide. In the lab, several catalysts can be used to accelerate the rate of reaction. These include:
In the body, the enzyme catalase is what catalyses the decomposition of hydrogen peroxide into water and oxygen gas. This process happens in nearly every living organism, including bees.
When doing the reaction in a lab, manganese (IV) oxide is generally the preferred catalyst to use. However, there is a wide range of catalysts to choose from and each one will have differing effectiveness.
Why Do Catalysts Speed Up Reactions?
Catalysts are able to lower the activation energy required for a reaction. This means that they can increase the rate of a reaction without being used up.
Therefore, at the end of a reaction, the leftover catalyst is able to be reused. This is very handy for commercial or industrial processes because less product is being consumed.
Decomposition of Hydrogen Peroxide Reaction
When you add a small amount of catalyst into a flask containing a solution of aqueous hydrogen peroxide, the first thing you will notice is an instant colour change.
In the presence of manganese (IV) oxide or iron (III) chloride, the clear solution will immediately turn black. As the catalyst works its magic, the hydrogen peroxide will begin decomposing very quickly.
When this happens, the solution will begin rapidly fizzing. This is caused by 2 things:
- Rapid decomposition means that a large amount of oxygen gas is being produced in the form of bubbles
- Decomposition reaction releases heat energy, making it an exothermic and fizzy reaction
As the decomposition of hydrogen peroxide continues, a lot of pressure will quickly build up in the flask due to the volume of oxygen gas being produced. The reaction will culminate in the mixture violently shooting upwards out of the flask.
This chemical reaction can be turned into a fun experiment for kids (and adults!). By adding some washing up liquid to the H2O2 solution, the final product is a thick foam that overflows out of the container – like squeezing elephant’s toothpaste.
The presence of soapy water is able to trap the oxygen that is released during decomposition. This creates a thick foam that is forced out of the container due to the build of pressure.
The most common catalyst used in this experiment is potassium iodide, but most catalysts will achieve the same effect if dish soap is added.
Measuring the Reaction
By adding a catalyst to an aqueous solution of H2O2 and recording the rate of reaction at specific time intervals, you can monitor the total volume of oxygen gas being produced. Here’s how to do it:
- Pour some aqueous hydrogen peroxide solution into a vertical flask
- Add a small amount of catalyst into the solution
- Quickly stopper the flask to prevent any evolved oxygen from escaping
- Record the volume of oxygen at specific time intervals throughout the reaction
- Using a graph, plot the volume of oxygen produced against the specific time
When you have completed your graph, you should be left with a curve. To calculate the rate of reaction, simply select a point on the curve and draw a line tangent to it.
Calculating the gradient of the tangent will give you the rate of decomposition at that specific point in time. The steeper the gradient, the faster the rate of reaction.
ReAgent not only sells top-grade hydrogen peroxide in a range of solutions, we also have various catalysts for sale in our online shop. From potassium iodide to iron (III) chloride, we’ve got the ingredients you’ll need to make your very own elephant toothpaste!
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