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:
The Science Behind Hydrogen Peroxide Decomposition
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.
Factors Influencing Natural Decomposition
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.
The Role of Environmental Conditions in Decomposition Rate
There are several environmental conditions that can influence the decomposition rate of hydrogen peroxide. These factors can either be naturally present and incidental, or they can be by design and intentional, and include the surface area of the container or the object which the hydrogen peroxide is in contact with.
For example, hydrogen peroxide can decompose faster on a concrete surface than one made of glass primarily because of the difference in the surface area.
Another environmental condition that can influence the decomposition rate is the temperature. Higher temperatures can make the decomposition rate faster because of the speeding up of molecular motions. Conversely, if you lower the temperature, you can slow down the decomposition rate as the molecular motion slows.
Hydrogen Peroxide Reactions
Hydrogen peroxide reacts chemically in a few different ways:
- Decomposition: How does hydrogen peroxide decompose? It’s thermodynamically unstable. Under higher temperatures and concentrations, it decomposes to form water and oxygen. Decomposition of hydrogen peroxide can be catalysed by other compounds, such as transition metals like silver and platinum.
- Redox reactions: Depending on the pH level, hydrogen peroxide has powerful reducing and oxidising (redox) properties. As an oxidising agent, it can remove electrons from other substances. As a reducing agent, it can reduce various inorganic ions.
- Forerunner to other peroxide compounds: As a weak acid, hydrogen peroxide forms hydroperoxide when combined with various metals and their compounds.
What Catalyst Is Used in the Decomposition of Hydrogen Peroxide?
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:
- Manganese (IV) Oxide
- Potassium Iodide
- Iron (III) Chloride
- Lead Dioxide
- Catalase
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.
Experimental Applications of Hydrogen Peroxide Decomposition
If you’re a chemistry teacher, you can have more engaging discussions about hydrogen peroxide by conducting experiments with your students. Here are some of the more interesting experiments you may want to try:
Genie in a bottle
For this experiment, you’ll need 50 to 100 mL of hydrogen peroxide with 30% concentration. Pour this into a two-litre bottle or container. Drop a small amount (about 2 to 5 grams) of manganese dioxide into the liquid. Oxygen will be released like a thick smoke.
Pasta rocket
For this experiment, use macaroni and a hydrogen peroxide solution with 3% concentration. The yeast in the pasta releases oxygen and provides thrust for your pasta rocket.
How Does Hydrogen Peroxide Decompose In Chemical Reactions?
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.
Demonstration of Hydrogen Peroxide Decomposition: Elephant’s Toothpaste
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.
Measuring Reaction Rates in Laboratory Settings
You can use water displacement techniques to accurately and precisely measure the reaction rate of hydrogen peroxide under different conditions.
Connect a tube or hose to an inverted graduated cylinder filled with water, and measure the rate of oxygen released per second or per minute based on the water displacement. Graphs can be useful in making extrapolations and spotting patterns.
Conclusion
Hydrogen peroxide is a relatively unstable compound. Extra oxygen tends to be released once the temperature and pressure increase beyond a certain range. Contact with surfaces that have microscopic crevices and cracks can increase the decomposition rate. Hydrogen peroxide has a wide range of applications, which include medical, industrial, and space rocket applications.
