Show your students that decomposition reactions are a key part of our everyday lives
The word oxygen comes from ‘oxygène’, the name that French scientist Antoine Lavoisier gave to what he called the ‘eminently respirable’, or breathable, part of air. Lavoisier realised there was a connection between the air we breathe and the chemistry of combustion. He showed that by heating a chemical called mercury calx – a red powder we now know as mercury(II) oxide – he could separate it into two components, one of which was a breathable gas, oxygen.
This type of reaction is called a decomposition reaction. In a decomposition reaction, substances are broken up into simpler parts (elements or simpler compounds), usually with heat, light or electricity. Lavoisier’s reaction – an example of thermal decomposition (using heat) – was of great interest to the scientific minds of the time because it helped prove the existence of a new element. While the decomposition of mercury oxide is of limited use today, there are other decomposition reactions that are a key part of our everyday lives, and others that could become increasingly important.
One example that illustrates the concept of a break-up reaction particularly well is the reaction that occurs inside a glow stick when you snap it. Glow sticks contains two chemical reactants kept separate by a glass cylinder – until you break the glass and the two mix. This mixing triggers an immediate reaction producing a very unstable compound (chemists have been trying for decades to confirm its identity). It’s so unstable that it rapidly decomposes, releasing carbon dioxide (CO2) and, crucially, energy. The coloured dye in the glow stick absorbs the energy, causing it to luminesce (or glow).
In your class
Demonstrate decomposition reactions such as the catalytic decomposition of hydrogen peroxide or the thermal decomposition of sodium nitrate to grab learners’ attention. Link these reactions to everyday life using the contexts from this article and follow up by testing their recall of word equations, symbol equations and conservation of mass in relation to simple decomposition reactions with our differentiated classroom resource.
Demonstrate decomposition reactions such as the catalytic decomposition of hydrogen peroxide (rsc.li/3HmJDIs) or the thermal decomposition of sodium nitrate to grab learners’ attention. Link these reactions to everyday life using the contexts from this article and follow up by testing their recall of word equations, symbol equations and conservation of mass in relation to simple decomposition reactions with our differentiated classroom resource.
Some decomposition reactions split compounds into their most basic components, as in the case of the decomposition of water. This particular reaction is attracting a lot of scientific attention because it can be used to produce hydrogen for fuel. Also known as water splitting, it’s a reaction you may have experimented with in a science class yourself. You can achieve it very simply, but only with the right equipment. ‘It doesn’t just happen spontaneously: it has to be done in a very specific way using a process called electrolysis,’ says Anna Hankin, head of the Electrochemical Systems Laboratory at Imperial College London.
As Anna points out, we can’t achieve this decomposition reaction by heating because the water just evaporates. So, in the school laboratory, we take a beaker of water and place a couple of electrodes in it. When the electrodes are connected to a battery, the current causes water to break up into its constituent parts, hydrogen and oxygen. It’s a neat experiment because you can watch the two gases bubbling off – oxygen at the positive electrode and hydrogen at the negative – and even capture them in test tubes. This only became possible around the time that Lavoisier’s ‘oxygène’ was gaining acceptance, after the invention of the battery by Italian physicist Alessandro Volta in 1800.
Most hydrogen is produced from fossil fuels today, so one exciting application of the decomposition of water is as a route to producing clean hydrogen – a fuel that produces no carbon emissions when it burns – for hydrogen-powered vehicles. Of course, making hydrogen in this way requires electricity for electrolysis. ‘The snag is that the hydrogen you get from water electrolysis is only environmentally benign if the electricity you’ve put into it is not made from burning fossil fuels,’ says Anna. So we can’t plug a beaker of water into a standard electricity socket and claim we’ve made ‘green’ hydrogen if that socket is powered by coal-fired power stations. It is possible, however, to do the same decomposition reaction and make hydrogen while plugged into a solar panel or a wind turbine. In fact, on the Orkney Islands they power several zero emission vans in this way through electrolysis driven by excess wind energy from its turbines.
Scientists want to integrate water decomposition with solar power in a single device that can directly use light to decompose water and this, it turns out, is much more difficult to do. ‘It makes sense to do this integration,’ says Anna, ‘but it means that you now have to solve all the tricky things that scientists and engineers have had to solve for each of the separate systems in one system.’ It’s a tough ask, but Anna thinks devices for the decomposition of water could at least be used in small communities and niche applications like, say, hydrogen-powered drones and lawnmowers.
Breaking up bleach
While we may be a few years away from having our own clean hydrogen-powered drones, there are decomposition reactions that we make use of every day. When we clean our bathrooms, for example, or get our hair coloured or even fill our printers with crisp, white sheets of printer paper. Hydrogen peroxide (H2O2) is the basis of some disinfectants. You may also know it as bleach, or the strong-smelling peroxide used in hair dye to lighten hair. As a bleach, it uses the chemical power of single, reactive atoms of oxygen to destroy molecules called chromophores that give things, like your hair, their colour. This oxygen, along with water, is produced in a decomposition reaction.
Hydrogen peroxide was discovered in 1818 by another French chemist, Louis Jacques Thénard. In one of Thénard’s experiments, he mixed hydrogen peroxide with blood, demonstrating that something in the blood caused the decomposition reaction to ‘go’, releasing the oxygen. Although, at the time, catalysts hadn’t been discovered, it wasn’t long afterwards that Swedish chemist Jöns Jacob Berzelius coined the term catalysis to describe this ‘go’ factor, using the decomposition of hydrogen peroxide by blood as an example. It turns out that blood contains an enzyme, catalase, that relies on iron to speed up the reaction.
This same reaction can be seen in the fizzing on a fresh wound if using hydrogen peroxide to clean cuts and grazes – although it’s not recommended by medical professionals these days. Hydrogen peroxide (also known as non-chlorine bleach) is intended for cleaning kitchens, bathrooms and surgical instruments, rather than for direct use on skin. The reactive oxygen released in the decomposition process kills disease-spreading bacteria and other nasties, but it doesn’t make a distinction between those single-celled invaders and the cells in our own bodies, even at the lower concentrations we use in our homes, so it can damage skin.
Decomposition reaction worksheet, for age range 11–14
A differentiated worksheet linking word equations, symbol equations and conservation of mass to simple decomposition reactions.
Decomposition reaction worksheets, for age range 11–14
Use this differentiated worksheet to link word equations, symbol equations and conservation of mass to simple decomposition reactions. Download the scaffolded and unscaffolded worksheets as MS Word or pdf, and the teacher notes including answers as MS Word or pdf at: rsc.li/3HH29N2.
Besides bleach, one of the most important applications of the hydrogen peroxide decomposition reaction is in paper mills, where the concentrations used are generally much higher. As Carl Houtman, a chemical engineer with the US Department of Agriculture’s Forest Service explains, the need for paper bleaching only developed when demand for paper began to outstrip the supply of white fibres for making it. ‘In the early days of paper production, naturally white fibres were often used, for example, cotton or linen rags,’ he says. ‘As the demand for paper outgrew the supply of white fibres, wood became the fibre of choice.’ Then, chemicals were needed to destroy the colour-carrying compounds in wood, turning it from brown to white.
While chlorine-based bleaches can also be used to turn paper white, their waste is more polluting to waterways and therefore hydrogen peroxide decomposition has become the first choice for the paper industry.
Under the conditions used in paper mills, decomposition of hydrogen peroxide is generally safe. However, the reaction needs to be properly controlled or it can go too fast, according to Carl. ‘As with most chemical reactions, hydrogen peroxide decomposition goes fast at higher temperatures, so a little reaction leads to heating of the solution, which results in significantly more reaction,’ he says. ‘This auto-acceleration in reaction rate can lead to explosions.’ It can happen when heat and pressure from the production of oxygen build up very rapidly. In one Texas paper mill, for instance, a 100-kilogram metal component was thrown 100 metres through a brick wall by the force of an explosion.
Thanks to Carl’s work, there’s now more guidance for paper mills on how to carry out the reaction safely. In a recent article, Carl and co-author Peter Hart showed that limiting the concentration of hydrogen peroxide used in paper bleaching to 10% should keep the process safe. The Texas plant has now put this into practice, ‘so they never have to suffer that explosion again,’ Carl says.
- Demonstrate the thermal decomposition of sodium nitrate by revealing an invisible message.
- Explore the concept of conservation of mass with a practical video and supporting resources for 14–16 year-old learners.
- Help students navigate the law of conservation of mass with this classroom activity and questions.
- Compare the thermal stability of different metal carbonates with this class practical.
- Reveal an invisible message to demonstrate the thermal decomposition of sodium nitrate for your students: rsc.li/40cxiiL
- Explore the concept of conservation of mass with a practical video and supporting resources for 14–16 year-old learners: rsc.li/3jczsOS
- Help students navigate the law of conservation of mass with this classroom activity and questions: rsc.li/3RpPJwn
- Compare the thermal stability of different metal carbonates with this class practical: rsc.li/3DNZU8D
There’s one decomposition reaction that anyone with a sweet tooth will be familiar with – even if they don’t know the chemistry. It involves sodium hydrogen carbonate, or bicarbonate of soda (NaHCO3), the active ingredient in baking powder.
Bakers worry endlessly about the ‘rise’ of their baked products. In yeasted breads, that rise is provided by bubbles of carbon dioxide gas produced in fermentation. In cakes and other sweet products, though, the carbon dioxide comes from the decomposition of sodium hydrogen carbonate. The reaction starts at around 80ºC, so it’s another example of thermal decomposition, but one we can experiment with in our ovens at home.
The process can also be seen very clearly when used in making honeycomb, a simple mixture of sugar, syrup and bicarbonate of soda. Heating the sugar and syrup makes toffee, then the bicarbonate decomposes when added to produce carbon dioxide, putting holes in the sweet treat that stick when the honeycomb rapidly sets.
All this goes to show that breaking stuff up can be pretty useful. And if you ever need an excuse for baking a cake then a scientific demonstration of a decomposition reaction is a very good one.
Hayley Bennett is a freelance science writer and editor based in Bristol, UK