Chemistry Laboratory Report Practical # 14 Heat Effects and Calorimetry
Partners: Kabimoldayev Ilyas
Section # 5
Dr. Eugene Douglass
Heat is an energy that can pass from high temperature to low temperature. Heat is measured by using of calorimeter. When heat comes to the substance, its temperature will increase.
q = S.H(specific heat) x m x
The S.H. of a metal can be easily calculated if the mass and temperature change of a sample are known (∆t – after mixing sample with water), since heat that is gained by water is equal to the heat lost by the metal, so qwater = - qmetal. Therefore, S.H. can be determined. Generally, S.H. is related to its molar mass, which is represented in the formula MM ≡ 25 / S.H. (where MM is molar mass of a substance; 25 is relative amount of heat required to raise one mole of many metals by 1 °C and S.H. is specific heat).
Results and Discussion
In the Experiment A, when metal was heated in the beaker with water, the level of water gradually decreased due to evaporation, so water was added each time to maintain the required level for covering all the metal, which was done to ensure that the metal would gain exactly the same temperature as water. When the metal was added to the water in calorimeter and it was swirled, the temperature of the water increased immediately. Using the data, heat absorbed by water was calculated (taking S.H. of water = 4.18 J/g°C). The heat that was absorbed by water is equal to the heat released by the metal, so using the givens the S.H. and approximate MM of metal were found. Also it can be said that the reaction is exothermic, since there was an increase in water’s temperature. According to the approximate molar mass found, it can be suggested that the unknown metal is either iron (~ 56 g/mol) or vanadium (~ 51 g/mol). However, vanadium is quite expensive metal, so it can be summarized that unknown metal is probably iron.
In the Experiment B, when unknown solid (# 3) was added to water in the calorimeter and it was swirled, the temperature of the water raised quite fast, which could also be noticed by warm calorimeter. Using the data, heat absorbed by water was calculated (taking S.H. of water = 4.18 J/g°C); and the heat for the reaction was deduced from it. However, the calculated heat was for the total mass of a solid, so to find the heat of solution per gram of solid, heat was divided by the total mass of solid. Here, water lost its heat, so the heat for the reaction has a positive sign, which means that the reaction is endothermic.
In the third experiment, when NaOH was poured into the calorimeter with HCl and it was swirled, the temperature of HCl increased, which was also sensed by hands. Using 1.02 g/mL as the density of solution, the mass was found by multiplying volume (50 mL) by the density. Then, the heat absorbed by the water was calculated. From that figure the heat for neutralization and the heat per mole of H+ and OH- ions reacting were calculated. Since the volume and molarity of solution are the same, the ∆H per mole of either is the same.
In the Experiment D, the same observations as in the Experiment C were observed (but instead of HCl CH3COOH was used).
Hess’s Law application (data from the first trial of C and D experiments)
Depending on the givens, ∆ Htotal is determined. If required, the second equation could have been inverted to get the inversed reaction.
During the practical, some errors could have been done. Firstly, it was assumed that no heat was lost to the surroundings (including the heat absorbed by the walls); however, some heat could have been lost through the hole in the cover or some of that could have been gone to the walls, so the original result (theoretical) can be different from the experimental. Also, after the first experiment it was found that there was a hole in the bottom of the inside-cup, so the liquid past to the second cup. Possibly, the hole was already there before the first experiment, so the results may be inaccurate. Due to the time limitation, two trials of the first experiment were not repeated. Also, regarding the first experiment, some water could have run into calorimeter from the test tube, which previously was in the water, so the results might have been changed. In addition to that, the initial temperature of the solutions in the experiments C and D is quite different, which could probably happened because the thermometer was not cleaned after each measurement but just dried by the tissue. So, for example, when measuring the temperature of NaOH, some of HCl could stay on the thermometer, which raised the temperature a little bit. Moreover, technical error could have been done as the thermometer itself could work incorrectly. Finally, some of the solution could have stayed in the beaker when the solution was poured into calorimeter, which could have a slight effect on the outcome.