[BLANK_AUDIO]. Okay. So welcome everyone to the first of our three virtual practicals. Essentially, chemistry is a practical subject and therefore we've tried to introduce this element into our course. The practical should be informative and should help to clarify material you've seen in the lecture videos, but they should also be kind of fun, and provide a break from the lectures themselves. Just like lab classes during a chemistry degree. The thermodynamics experiment is entitled The Enthalpy of Hydrogen Peroxide Decomposition in Aqueous Solution. If you haven't already done so, I strongly suggest you download the experimental sheet that accompanies the practical, and review it before watching the remainder of this video. The objectives of this experiment are two fold. First, to illustrate the use of calorimetry in the experimental measurement of thermodynamic data. And second, to illustrate the use of thermodynamics and specifically Hess' Law calculation of enthalpies of reaction. The experiment uses a calorimeter comprising a geo vessel, fitted with an electrically driven stirrer, and a highly accurate mercury thermometer. The calorimeter is well insulated from its surroundings, so when an exothermic reaction is carried out inside the calorimeter, the temperature rises. In a typical experiment, the temperature inside the calorimeter is monitored for a period of five minutes. After exactly five minutes, the reaction is initiated; which releases heat into the calorimeter. The stirrer distributes the heat rapidly throughout the vessel. And after an initial significant jump, the temperature inside the calorimeter settles down once more, and is monitored for a further five minutes. [BLANK_AUDIO]. Extrapolated linear trend lines are then fitted to the pre and post reaction data sets, and the temperature rise in the calorimeter is determined by graphical means. This is all explained in detail on the experimental sheet. In this experiment, the reaction between sulphamic acid and sodium nitrate to give nitrogen, water, and sodium hydrogen sulphate, will be carried out in the calorimeter. The reaction has a known molar enthalpy change of minus 420.5 kiljoules per mole. Given the sulphamic acid will be in excess, the amount of heat released by the reaction will depend on the quantity of sodium nitrate, but this is known. Consequently, we can calculate exactly how much heat energy is released. If we measure the temperature rise, we can then determine the heat capacity of our calorimeter. In other words, how much heat energy it takes to raise the temperature inside our calorimeter by one degree. This step constitutes calibration of our calorimeter. We then perform a second reaction in the calorimeter, this time one with unknown molar enthalpy change, the decomposition of hydrogen peroxide into water and oxygen, catalysed by manganese dioxide. By measuring the temperature change, and using this in conjunction with the heat capacity of the calorimeter, we can determine how much energy was released. And since we know how much hydrogen peroxide was reacted, we can determine the molar enthalpy change. From this value, you will then perform a series of Hess' law calculations to determine the molar enthalpy of decomposition of pure hydrogen peroxide, and the oxygen oxygen bond dissociation energy in hydrogen peroxide. Or in other words the strength of the oxygen-oxygen bond. Now you need to pause this video and prepare the following table on a piece of paper, or in a notebook, or in electronic device. Although I imagine the latter option might be tricky to fill in as the experiment progresses. Once you have the table complete, restart the video and keep the list and your pen handy, as you're going to track the temperature of the calorimeter for ten minutes. [BLANK_AUDIO]
