 In this lesson we will learn about how a catalyst affects the rate of the decomposition of hydrogen peroxide. Fairly dilute solutions of hydrogen peroxide approximately 3-30% volume by volume can be used as a household cleaning agent to remove stains on furniture and clothing. It can also be used to bleach or lighten darker hair. Whatever the real life uses, hydrogen peroxide is always stored in a dark plastic container away from sunlight. This is because in the presence of warm conditions and UV light, hydrogen peroxide can decompose into water and oxygen because the oxygen-oxygen bond in the molecule is unstable. Since aqueous hydrogen peroxide and water are clear and colorless liquids, we will monitor the rate of this reaction by measuring the amount of oxygen produced. This decomposition reaction can be catalyzed by manganese foroxide. The oxygen produced can be measured using a gas syringe. Place a set amount of hydrogen peroxide and add a small spatula of manganese foroxide, a fine black powder. Quickly stopper the flask so to prevent any evolved oxygen from escaping into the surroundings. You will see that fizzing happens which shows that oxygen gas is being produced. This is an exothermic reaction, meaning that heat energy is released. The volume of oxygen evolved can be recorded at specific time intervals. Plotting the volume of oxygen produced right against time will give a graph that is similar to this one right here. We can measure the rate of this reaction at any time in this reaction by selecting a point on the curve and drawing a line tangent to the curve. Calculating the gradient of the tangent will give you the rate of the decomposition of hydrogen peroxide at that point in time. A steeper gradient is representative of a faster rate of reaction whereas a less steep gradient is representative of a slower rate of reaction. As you can see at the start of the reaction the rate of decomposition is very fast.