A 2M HCl solution was prepared for certain parts of the experiment by carefully adding 100 mL 6M HCl to a beaker with 200 mL deionized water using a graduated cylinder, and the solution was stirred. Then, a 2 M NaOH solution was prepared adding 100 mL 3 M NaOH to 50 mL of deionized water in a beaker, and stirring. These reactions generated heat, and the solution was allowed to cool down before further reactions were performed. The LaqQuest was turned on, and the temperature probe was connected to Channel 1 of the LabQuest. The settings for the data collection were changed as needed so that the “Mode” was set to “Time Based,” the “Interval” was set to “15s/sample,” which caused the Rate to adjust automatically, and the “Duration” was set to …show more content…
The LabQuest was made to initiate Run 2 with the same parameters as Run 1 by clicking on the File Cabinet icon in Graph view. Then, the steps for collecting data using the LabQuest in reaction 1 were repeated for reaction 2, but instead of NaOH being added to HCl, 2 M NH4Cl was added from the graduated cylinder to the 2 M NaOH in the Styrofoam cups. After the data collection, the calorimeter, temperature probe, and graduated cylinder were cleaned, and the reaction solution was disposed of, similar to reaction 1. The calorimeter was then set up again for the third …show more content…
To obtain the final reaction, the first reaction was added to the reverse of the second reaction. As a result, the sign of the ΔHrxn of reaction two was changed from negative to positive. The sum of the ΔHrxn of reaction 1 and the reverse of reaction two was: -55.84 kJ/mol + 3.62 kJ/mol = -52.22 kJ/mol This value (-52.22 kJ/mol) was also calculated, using the provided enthalpies of formation, to be the ΔHrxn of reaction 3, the sum of reactions 1 and 2. Hess’ law was hence applicable to this reaction. The temperature measurements taken in part B were used to calculate the amounts of heat energy absorbed by the various reactions. The heat (q) for the various sub-reactions was calculated by using the formula q = mCΔT + CΔT where q was the heat, m was the mass of the solution, ΔT was the change in temperature, C was the specific heat of water, and C was the calorimeter constant. The mass of the solution was calculated by multiplying the density of water by the volume of water. In part B, this resulted in the general equation: (1.03 x 100 mL x 4.184 x ΔT) + (25.5 x
First, 100 mL of regular deionized water was measured using a 100 mL graduated cylinder. This water was then poured into the styrofoam cup that will be used to gather the hot water later. The water level was then marked using a pen on the inside of the cup. The water was then dumped out, and the cup was dried. Next, 100 mL of regular deionized water was measured using a 100 mL graduated cylinder, and the fish tank thermometer was placed in the water. Once the temperature was stabilizing in the graduated cylinder, the marked styrofoam cup was filled to the mark with hot water. Quickly, the temperature of the regular water was recorded immediately before it was poured into the styrofoam cup. The regular/hot water was mixed for a couple seconds, and the fish tank thermometer was then submerged into the water. After approximately 30 seconds, the temperature of the mixture leveled out, and was recorded. This was repeated three
Theory of Water of Displacement: A volume of water was measured. A second volume of water was measured when the metal cylinder was added. The initial volume was subtracted from the second (total) volume to get the volume of the metal cylinder.
Next, we measured 1.07 g of magnesium oxide, using a balance in the fume hood, added it to the HCl in the calorimeter, and shut the lid quickly to conserve heat. This mixture was “swirled” and allowed a few moments to react. The final temperature was recorded and DT determined. GRAPH GRAPH
Experimental: The experimental procedure outlined in the OU Physical Chemistry Laboratory Manual was followed without any deviations.
•The sum of the energies released to form the bonds on the products side is ◦2 moles of H-H bonds = 2 x 436.4 kJ/mol = 872.8 kJ/mol
Experiment is to investigate the rate of reaction between hydrochloric acid and calcium carbonate Hydrochloric acid + Calcium Carbonate Þ Calcium Chloride + Water + Carbon Dioxide 2HCl (aq) CaCo3 (s) CaCl2(s) H2O (aq) CO2 (g) There are a number of variables in this experiment and these are listed below as input variables and outcome variables.
To start this study, nine labeled test tubes were setup with precise amounts of 2mL of deionized water, and 1ml of 50-50 corn syrup to water mixture. The addition of 1mL of yeast would also be added, but this will not be added until the fermentation apparatus is assembled in the water bath and ready to begin the reaction. The assemble of our apparatus included submerging and combining of the test tube and tubing with a stopper to ensure no air is in the apparatus. Then the assemble would be put this apparatus with water inside a Styrofoam cup, to ensure temperature is conserved best, and prepare to add the test tube with controlled substance to the test tube and stopper. The water baths at different temperatures are the only variables changed. One water bath was set up as the control group at room temperature, 28°C. The second water bath was setup to 0.4°C by use of ice water, and third bath used hot water at 49°C. Right before adding the test tube with control substance, the yeast would be added to create the reaction that produced the gas. To ensure best accuracy of fermentation, an initial test tube with all substances but yeast was performed to obtain an initial equilibrium time. Measuring of this time occurs till no more air is bubbling out of tube. This time is where we would mark are initial measuring line for each of the following reactions. As the gas pushes the water out of the test tube
" This means that therefore the enthalpy change of a reaction can be measured by the calculation of 2 other reactions which relate directly to the reactants used in the first reaction and provided the same reaction conditions are used, the results will not be affected. We have the problem set by the experiment to determine the enthalpy change of the thermal decomposition of calcium carbonate. This is difficult because we cannot accurately measure how much thermal energy is taken from the surroundings and provided via thermal energy from a Bunsen flame into the reactants, due to its endothermic nature. Therefore, using the enthalpy changes obtained in reaction 1 and reaction 2 we can set up a Hess cycle.
For this experiment, you will add the measured amount of the first sample to the measured amount of the second sample into its respectively labeled test tube then observe if a reaction occurs. In your Data Table, record the samples added to each test tube, describe the reaction observed, if any, and whether or not a chemical reaction took place.
Tf-Ti). Next, subtract the initial temperature, 25 degrees from the final temperature, 29 degrees putting the change in temperature at 4 °C. To calculate the heat absorbed by the water in calorimeter, use the formula (q = mCΔT). Plug in 50 mL for (m), 4.184 J for (C) and 4 °C for the initial temperature (ΔT), then multiply.
Investigating Factors that Affect the Amount of Heat Produced in Neutralisation I am going to investigate factors that affect the amount of heat
During this reaction the solution gained heat. This is what we were monitoring. The reason why the solution gained heat is because the reaction lost heat. Energy is lost when two elements or compounds mix. The energy lost/ gain was heat. Heat is a form of energy as stated above in the previous paragraph. The sign of enthalpy for three out of the four reactions matches what was observed in the lab. For the last reaction, part four, the reaction gained heat not the solution like parts one through three. The negative enthalpy value for part four indicates that the reaction gained
Input variables In this experiment there are two main factors that can affect the rate of the reaction. These key factors can change the rate of the reaction by either increasing it or decreasing it. These were considered and controlled so that they did not disrupt the success of the experiment. Temperature-
In a 100ml beaker 30mls of water was placed the temperature of the water was recorded. 1 teaspoon of Ammonium Nitrate was added to the water and stirred until dissolved. The temperature was then recorded again. This was to see the difference between the initial temperature and the final temperature.
... model for the thermodynamics and fluid mechanics calculations for this system need to be presented.