2M Hcl Chemistry Lab Report PDF

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2M Hcl Chemistry Lab Report...

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2M Hcl Chemistry Lab Report The first step of the experiment was to prepare the 2M HCl solution for Part A. 200 mL of deionized water was placed into a 600 mL and then 100 mL of 6M HCl was added to dilute the acid. 100 mL of 3M NaOH and 50 mL of deionized water were then combined to form 150 mL of a 2M NaOH solution. The Lab quest then had to be set up so that the temperature changes taking place in the coffee cup calorimeter could be recorded. The temperature probe was plugged in and the “Time Based” data collection settings were adjusted so that there was a 15 s/sample interval so that there was an 180 second cycle for part B. The calorimeter was then set up so that the first reaction of hydrochloric acid and sodium hydroxide could take place. Two Styrofoam cups were stacked together and placed in a 400 mL beaker to which 50 mL of HCl were added. The calorimeter was then sealed and put on the Bunsen burner. The initial temperature readings were then gathered to provide a baseline. After the initial temperatures had been established, the NaOH was added quickly and stirred. After the max temperature had been reached and the reaction had stopped the data collection ended. The waste products were then thrown away in the proper container and the materials were cleaned. The second

reaction for Part B was Sodium Hydroxide and Ammonium Chloride. The data from the first reaction was saved it could be analyzed at a later time. The calorimeter was set up again with two stacked Styrofoam cups placed in a 400mL beaker. 50 mL of NaOH was assed to the Styrofoam cup and the calorimeter was sealed so that the initial temperature could be collected. After the initial temperature had been collected 50 mL of 2M Ammonium Chloride was added quickly and the calorimeter was then stirred and data was collected. After the reaction had stopped and data collection had ended the waste products were disposed of and the material were cleaned. The third and final reaction of part B which was hydrochloric acid and Ammonium Hydroxide was then set up. The data from the second reaction was saved on the lab quest and for later analysis. The calorimeter was set up by stacking two Styrofoam cups inside of a 400 mL beaker. 50 mL of 2M HCl was added to the Styrofoam cup and the initial temperature was then measured and collected by the lab quest. After the initial temperature had been collected 50 mL of 2M Ammonium Hydroxide was added to the cup. The calorimeter was stirred and the temperature changes were recorded. After the reaction was over and data collection had stopped the maximum/minimum and temperature change for each reaction was recorded in Table one and the lab data was saved onto a

USB drive. The waste products were then disposed of and the materials were cleaned for part C. Part C of the experiment included to reactions: Hydrochloric Acid and Magnesium Oxide, and Hydrochloric Acid and Magnesium Metal. The Lab Quest had to be adjusted so that the interval was 15 s/sample and so that the duration was 480 seconds. The calorimeter was then constructed by staking two Styrofoam cups together and lacing them in a 400 mL beaker. 100 mL of 6M HCl was added to the Styrofoam cups and the initial temperature was recorded by the lab quest. One gram of MgO was weighed out using scale and watch glass and all weights were recorded in table 2. After the initial temperature was recorded the MgO was quickly added and the cup was swirled while the temperature changes were recorded. After the data collection as over the reaction stopped the reaming MgO residue on the watch glass was recorded and all waste products were disposed of and all material were cleaned. The data form reaction four was saved on the lab quest so that reaction five could be carried out. The Styrofoam cups were restacked and placed in a 400 mL beaker to which 120 mL of 2M HCl was then added the calorimeter was then closed and the initial temperature was recorded by using the lab quest. 0.5 g of Mg

was then gathered and weighed using the scale and watch glass, all weights were recorded in table 3. Once the initial temperature was recorded 0.5 g of Mg was quickly added to the Styrofoam cup and the calorimeter was then swirled while the reaction occurred and the temperature changes were recorded. After the reaction had finished and the temperature changes had been recorded the lab quest was used to view the table with the data from reactions four and five. The maximum and minimum temperatures were recorded in table 4 for each reaction as well as the temperature change. The data was then saved onto a USB drive for further analysis and all the waste products were disposed of properly. All materials used in the experiment were washed thoroughly, dried, and put away. RESULTS & DISCUSSION This lab evaluates reactions for temperature change so that reaction enthalpies and enthalpy of formation could be calculated. The temperature readings that were gather from each sub reaction allowed for the use of multiple methods to calcite the enthalpy which include the theoretical calculation, Hess’s Law of Summation calculation and an experimental calculation. Through using all three methods, the experimental data calculations were able to be evaluated for percent

error. The different methods also gave a better understanding of what reaction and formation enthalpy is and how that helps to decide whether or not something is endothermic or exothermic. The theoretical enthalpy calculations for all three sub reactions in part B were found by using the provided net ionic equations as well as the provided table of enthalpy values for each of the compounds in those equations. Hess’s law was used to subtract the enthalpy of the products minus the enthalpy of the reactants, which then provided the total enthalpy change of the “theoretical” molar enthalpy. The reactions for part B were all found to be exothermic, meaning that they released heat into the surroundings. This was determined by looking at the theoretical values of reaction enthalpies and observing that they were all negative which means that the reaction must be exothermic. It was also proven to be exothermic through the Q values that were calculated from the experimental data, which were found to have positive Q values and thus a negative delta H making the reactions exothermic. The Q values in part B depended upon the data collection that was done so accurate temperature readings of the reactions and measurements of materials were needed. The first step to calculate required that the mass be found of each of the solutions since only the

density and molarity was given. The mass was calculated by adding together both substances used in the reaction and then multiplying it times the density thus converting it to mass in kg. Once the mass was found the number could simply be stuck into the equation of (M of water *Specific heat of water* Temperature change)+(Calorimeter value*temperature change). The only changing values for each reaction were the masses and the temperature changes. The enthalpy of reaction for part B was calculated first by using the Q value as described above. The limiting reactant was also found for each sub reaction. The limiting reactant was found by taking the volume of each substance in liter and multiplying it times the moles of the substance used and then the two values for each compound in the reaction were compared to find the lesser one which is the limiting reactant (both were equal in all three sub reactions for this experiment so .1 M was used regardless). The enthalpy calculation itself was done by dividing the Q values of the system by the moles of the limiting reactant and then converting the final answer from joules to kilojoules. This came out to be a negative enthalpy value for each sub reaction, once again confirming an exothermic reaction. The molar enthalpy change of reaction three was then found by using Hess’s

Law of Summation. This essentially allows for reactions or equations with know enthalpies to be added together in order to calculate an unknown enthalpy. However, the equations must cancel out any extra compounds so that only the compounds in the desired third equation remain as well as all the compounds must be on their correct side and this is where the multiplying a dividing as well as flipping comes into play . The equation that was desired in this experiment was hydrogen plus ammonia to yield ammonium and heat. The adding for this calculation was relatively simple as all that needed to be done was to flip reaction two which then placed all the compounds on their correct ides and allowed the additional compounds to cancel out. Thus, the enthalpy values for reaction two was flipped to become positive and was added to reactions one’s enthalpy which remained the same which made the enthalpy of reaction three -55.687 kJ. The three methods for determining the enthalpy of reaction three for part B showed values that all differed slightly form one another. The theoretical value for enthalpy calculated using the given values table 5 was -55.22 kJ. The enthalpy of reaction three calculated using the experimental data or the temperature changes recorded was done by using “q” values and limiting reactant molarity and thus that value was

found to be -60.70 kJ. The third calculation using Hess’s Law of summation was found using the experimental enthalpies for reactions one and two which were then combined to calculated a value for reaction three of -55.68 kJ. The actual experimental data was shown to be much higher than the other two calculations. The summation calculation was actually closer to the predicted value despite using experimental data. This could be due to inaccurate temperature measurements or slightly larger amounts of reactants added, however, the values were reasonably close together and can thus be considered accurate. Part C of the experiment focused on reactions of HCl and MgO and HCl and Mg, which allowed for the evaluation of enthalpy of formation as well as the similar calculations to those performed in part B. Two reactions were carried out and the temperature changes were recorded. The reactions in part C were all exothermic like part B. This is proven by the reaction enthalpies that were calculated both theoretically and experimentally all of which were calculated to have negative values meaning heat tis released form the system. It was necessary to use two different concentrations of HCl in order to have reactions of the same magnitude because if 6M HCl was used on the

second reaction of solid magnesium the reaction would have been to large. Similarly if 2M HCl were used on MgO then the reaction would have been too small to properly measure....