Ib Chemistry Experiment- Calculating Enthalpy Change

Chemistry Internal Assessment Determining the Enthalpy sort of a Dis aspirement Reaction AIM To determine the enthalpy convince for the reaction between copper(II) sulfate and coat. BACKGROUND THEORY Bond breakage is endothermic while bond forming is exothermic. The reaction between copper(ll) sulfate and zinc is exothermic as the nil required to form the bonds of the products is greater than the energy required to break the bonds of the reactants. In an exothermic reaction, inflame is given finish up to the surroundings thus, temperature of the surroundings will increase. By measuring the spay in the temperature and using the formula Q= mc?T, we throne calculate the enthalpy reposition of the reaction. Equation 1 CuSO4 + Zn ? ZnSO4 Ionic Equation Zn (s) + Cu2+ (aq) ? Cu (s) + Zn2+ (aq) MATERIALS/APPARATUS * 1 insulated Styrofoam cup * Copper(II) sulfate etymon * Zinc mill * 1 Thermometer * 1 S aggrandisementwatch * Weighing Boat * Electronic counterbalance VARIABLES Ind ependent Dependent Mass of zinc powder and concentration of copper(II) sulfate outcome used. Temperature of the solution PROCEDURE 1. Use a pipette to beat 25. 0cm3 of 1. 0 M copper(ll) sulfate to the insulated container. 2. Record the temperature every 30 seconds for 2. 5 minutes 3.Add the excess zing powder (6g) at exactly 3 minutes 4. Stir and record the temperature every 30 seconds for the following 10 minutes. DATA COLLECTION AND PROCESSING Time (s) Temperature (C) Time Temperature (C) 30 25 390 62 60 25 420 61 90 25 450 60 120 25 480 59 150 25 510 58 clxxx 25 540 56 210 45 570 55 240 52 600 54 270 56 630 52 300 60 660 51 330 61. 5 690 50 360 62 720 49 Therefore, base on the graph shown above (representing the raw data), the change in temperature if the reaction had taken place instantaneously with no heat loss ?T= 70. 5C ? 25C 45. 5C The volume of the copper(II) sulfate solution used was 25cm3, thus the mount of the solution is 25g. Given that the specific heat capacity of the solution is 4. 18 J/K and the temperature change is 45. 5C, as calculated above, thus, the heat, in joules, produced during the reaction derriere be calculated using the formula Q = mc? T =mass of solution ? specific heat capacity of solution ? temperature change = 25 ? 4. 18 ? 45. 5 = 4754. 75 J In the experiment, 25cm3 of 1. 0 mol dm-3 copper(II) sulfate solution was used. Thus, number of moles of the copper(II) sulfate solution used n(CuSO4) = (25? 000) ? 1. 0 = 0. 025 mol Therefore, the enthalpy change, in kJ/mol, for this reaction is ?H = Q ? n(CuSO4) = 4754. 75 ? 0. 025 = -190. 19 kJ/mol Theoretical value/ Accepted look upon= ? 217 kJ/mol Thus, percentage error = (? 217+190. 19) ? (? 217) ? 100 = 12. 35% CONCLUSION Thus, based on the experiment, the enthalpy change for the reaction is -190. 19 kJ/mol. However, as we passel see from the above calculations, the percentage error is 12. 35%. This means that the lead is in consummate from the theoretical value of -217 kJ /mol by 12. 35%.From the graph, we bear also see that once zinc is added to the solution (at exactly 3 minutes), the temperature of the solution increases until it reaches the terminal or maximum temperature of 61C. Then, the temperature of the solution bit by bit decreases until it reaches room temperature again (temperature of the surroundings). EVALUATION (WHAT CAN BE DONE TO correct THE EXPERIMENT? ) An assumption made for this experiment is that none of the heat produced by the exothermic reaction is lost to the surroundings and that the thermometer records the temperature change accurately. However, this is very unlikely to appen in reality, which would explain the percentage error. Thus, to improve the experiment, we can try to minimise the heat loss to the surroundings. This can be done by place a piece of cardboard (or any other insulated material) on top of the cup to cover the top of the cup. A hole can then be made in the cardboard for the thermometer. Another meas ure that we can take is to ensure that our eye is level with the thermometer when reading the temperature off the thermometer. We can also repeat the experiment a few times and hold fast the average of the results recorded. This would allow us to obtain a more accurate value.