# ?Thermochemistry Laboratory Report Essay

Published: 2020-07-15 05:50:05
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GET MY ESSAY AbstractionThe intents of these three experiments are to find the heat capacity of a calorimeter and with that informations, confirm Hess’s Law and observe enthalpy alterations within reactions. By mensurating the alteration in temperature that occurs with the interaction of two different reactants, we were able to find both the calorimeter invariable and the alteration in heat content of a given reaction. The consequences were instead assorted, as some Numberss more closely resembled the theoretical values than others did.
IntroductionThe first experiment is devoted to happening the calorimeter invariable for a polystyrene cup. Whenever a reaction takes topographic point inside a calorimeter, some heat is lost to the calorimeter and its milieus. In order to accomplish maximal truth, we must cognize precisely how much heat will be lost, so that the consequences of the following two experiments will be every bit right as possible. The equation used to find it is a simple use of the overall heat of the reaction equation, which is: Overall Heat = – [ ( Sp.Ht. hotwater * Mass of H2O * Change in temperature ) + ( Sp.Ht. coolwater * Mass of H2O * Change in temperature ) + ( Cp calorimeter * Change in temperature ) ] Since an mistake is bound to go on during the experimental procedure, three computations were done to happen an norm. This experiment is critical to the success of the undermentioned two thermochemistry experiments.
The 2nd experiment, entitled Hess’s Law, is a simple verification of said jurisprudence. To make so, we take three reactions, where one of them is the same as the other two, and step the heats of reaction for each of them. Hess’s Law states that the heat of reaction of the one reaction should be to the amount of the heats of reaction for the other two. The three reactions used in this experiment are: ( 1 ) NaOH ( s ) ? Na+ ( aq ) + OH- ( aq )
( 2 ) NaOH ( s ) + H+ ( aq ) + Cl- ( aq ) ? H2O ( cubic decimeter ) + Na+ ( aq ) + Cl- ( aq ) ( 3 ) Na+ ( aq ) + OH- ( aq ) + H+ ( aq ) + Cl- ( aq ) ? H2O ( cubic decimeter ) + Na+ ( aq ) + Cl- ( aq ) In order to happen the heat released by each reaction, we used a discrepancy of the overall heat of a reaction equation, which was q = – [ Sp.Ht. * thousand * Change in temp. ] . In add-on to happening the alteration in heat content, alteration in information was besides calculated utilizing theoretical values in given mention tabular arraies. Finally, the overall free energy released was calculated utilizing the equation: Change in free energy = Change in heat content – ( Temperature * Change in information ) . All of this is so used to verify Hess’s Law by ciphering the per centum mistake involved in the experiment.
The 3rd experiment, called Thermochemistry: Acid + Base, combines the constructs of the old two experiments. The chief construct is to detect the alteration in heat content that consequences from the assorted reactions between strong and weak acids and bases. There were four reactions used in this experiment, and they are:
( 1 ) HCl ( aq ) + NaOH ( aq ) ? NaCl ( aq ) + H2O ( cubic decimeter )
( 2 ) HCl ( aq ) +NH3 ( aq ) ? NH4Cl ( aq )
( 3 ) HC2H3O2 ( aq ) + NaOH ( aq ) ? NaC2H3O2 ( aq ) + H2O ( cubic decimeter )
( 4 ) HC2H3O2 ( aq ) + NH3 ( aq ) ? NH4C2H3O2 ( aq )By supervising the alteration in temperature that consequences from the reaction of an acid and a base, it is possible to cipher the overall energy for each reaction, besides known as ?H rxn/mole of restricting reactant. This experimental value can be compared with the theoretical to find how accurate the experiment was. The lower the per centum mistake, the more accurate we were at ciphering the energy involved in each reaction.
ExperimentalIn order to make any computation for energy, we foremost had to happen the calorimeter invariable. In order to make that, we foremost took and weighed a polystyrene cup ( our calorimeter ) and added about 100 g of warm H2O to it. The existent measurings are recorded in Table 1-1. The mass of the cup with the H2O in it were recorded to happen the exact mass of the H2O added. Next, a cylinder was weighed, like the cup, and about 48 milliliters of cool H2O was added. The sum was weighed and recorded in the same tabular array. Afterwards, temperature detectors connected through a LabPro device were suspended in the two containers and the calculator’s DataMate plan was used to enter temperature over a 90 2nd clip interval. After a few seconds of informations aggregation from the separate liquids, they were assorted together and stirred with the detectors until there was no clip left. By utilizing Graphical Analysis, a graph of the information was printed, exposing temperature vs. clip. Tangent lines were drawn on the graph in order to find the initial and concluding temperatures of the two liquids. The above process was repeated two more times for the interest of preciseness. Finally, we calculated the calorimeter changeless utilizing the expression listed in the Introduction subdivision.
Even though we conducted an experiment to happen the heat capacity of a calorimeter, we were given a new value for the invariable for experiment 2, due to inaccuracy in our consequences. For the lab called Hess’s Law, we foremost started by puting up the reckoner to roll up temperature informations once more. The process is the same as the one used in the last experiment, except that the clip interval is set to 4 proceedingss. Following, we obtain a polystyrene cup to utilize as our calorimeter and make full it with 100 g of H2O. The cup is placed within a 250-mL beaker to maintain it in a sustained environment. A temperature detector is placed in the H2O and is stabilized. Then, we obtained solid NaOH and weighed about 2 gms to the nearest thousandth denary point. This value is recorded, along with all other informations in Table 2-1. Afterwards, informations aggregation Begins and after approximately 15 seconds, the NaOH is added to the H2O. The resulting solution is stirred for the continuance of the clip interval and by utilizing Graphical Analysis a graph is produced. This process is repeated twice more for 0.5 M HCl in topographic point of H2O for one test, and so 1.0 M HCl and 1.0 M NaOH solution for the 3rd test. All of the measurings are recorded in the tabular array mentioned above.
For the concluding experiment, the process is really similar to its predecessors. We began by initialising the LabPro and DataMate to roll up temperature informations over clip ( this clip it is a one hundred eighty 2nd interval ) . First, we measure every bit near as we can to 50 g of a base of our pick in a 100-mL calibrated cylinder. A temperature detector is placed in the cylinder. Following, we weighed 100 g of a chosen acid in the calorimeter. The calorimeter is placed in a 1000-mL beaker for stableness and a temperature detector is submerged in the acid. After the detectors have a opportunity to equilibrate, we started to roll up informations. When about 15 seconds have passed, we poured the base into the calorimeter with the acid and stirred for the continuance of the clip with both detectors. Then, when clip was up, we used Graphical Analysis to publish the resulting temperature vs. clip graph. This processed is repeated three more times until every combination of strong and weak acids and bases is used.
AnalysisThe information we recorded for the first experiment appears to be accurate, though pulling tangent lines to happen concluding and initial points has its built-in inaccuracy. Using the expression discussed in the debut, our equation turned out like the followers:
0 = – [ ( 47.166 g * 4.184 J/g°C * 16.561 °C ) + ( 98.874 g * 4.184 J/g°C * -9.4139 °C ) + ( Cp calorimeter * -9.4139 °C ) ]Cp calorimeter = -66.522 J/°CThe norm of the three obtained values is every bit simple as adding them all together and spliting by three, the figure of values, which looked like this: ( -66.522 + 348.619 + 225.669 ) /3 = 169.255 J/°C. This figure is much higher than the default value we were given for the following lab, which was merely 15.0 J/°C.
For the Hess’s Law experiment, the Numberss looked much better. The first thing we did with the information was solve for the alteration in temperature, which was merely concluding temperature minus initial temperature. The consequence gave us something like this: 23.9 °C – 19 °C = 4.9 °C. Second, we calculated the heat released by each equation, which is shown as this: Q = – [ Sp.Ht. * m * ?t ]
Q = – [ 4.18 J/g°C * 99.524 g * 4.9 °C ]Q = -2.038 kJThen, the heat lost to the calorimeter was calculated utilizing the expression Q = – [ Cp * ?t ] . From that, we found that q = – [ 15.0 J/°C * 4.9 °C ] = -0.0735 kJ. Following, the entire ?H was found by adding both values of Q above, which merely peers -2.1115 kJ. In order to happen ?H/mol NaOH, we had to happen how many moles were used in each reaction based on the mass of NaOH weighed and recorded in Table 2-1. The format for happening the figure of moles looked like the followers: 2.0810 g NaOH * ( 1 mol NaOH / 40 g NaOH ) = 0.052 mol NaOH. This value is used to split the ?H to happen the ?H/mol NaOH value, which equaled -40.606 kJ/mol. Using the ?H of Reaction 2 as the theoretical value, and the combined ?H values of Reactions 1 and 3, we can happen out our per centum mistake, which is shown below as:
% mistake = acrylonitrile-butadiene-styrene ( ( theoretical – experimental ) / theoretical ) * 100% mistake = acrylonitrile-butadiene-styrene ( ( 79.56 – 94.87 ) / 79.56 ) * 100% mistake = 19.24 %The above values can all be found on Table 2-1. The above procedure was repeated with informations collected from the whole category, which yielded a 14.47 % mistake. Finally, utilizing theoretical Numberss, we calculated ?H, ?S, and ?G for reaction 2. For the first two, a similar equation of amount of merchandises minus amount of reactants peers ?H and ?S severally. ?G is calculated utilizing the expression in the debut, which looked like ?G = -98.8 – 298 ( 0.0580 ) = -116.062 kJ/mol.
With the informations collected in the 3rd experiment a battalion of computations were carried out. All of the undermentioned informations can be found in Table 3-1. First, we solved for ?H rxn, which is the same as the overall heat equation described in the debut. The computation looked liked the followers:
?H rxn = – [ ( 4.184 J/g°C * 98.781 g * 4.35 °C ) + ( 4.184 J/g°C * 48.5133 g * 4.0 °C ) + ( 15.0 J/°C * 4.35 °C )?H rxn = -2.68 kJFollowing, we needed to cipher the modification reactant for each reaction, which was merely the reactant that yielded the least merchandise. The method for finding it is like so:98.781 g HCl * ( 1 mol HCl / 36 g HCl ) * ( 1 mol NaCl / 1 mol HCl ) * ( 1 g NaCl / 1 mol NaCl ) = 2.744 g NaCl48.5153 g NaOH * ( 1 mol NaOH / 40 g NaOH ) * ( 1mol NaCl / 1 mol NaOH ) * ( 1 g NaCl / 1 mol NaCl ) = 1.213 g NaCl Then, we take the ?H rxn above and split it by the moles of restricting reactant, which we discovered above ( since each solution is 1.0 M, the moles used is the figure of gms divided by 1000 ) . This new ?H rxn / moles of restricting reactant is the experimental value to becompared to the theoretical value obtained with given Numberss. Comparing these two values utilizing the % error equation above, the % mistake of one of the reactions comes out to be merely 1.25 % . The remainder of the Numberss can be observed in Table 3-2. This concludes all of the computations that were involved in all of the experiments.
DecisionThe consequences of this experiment were a mix of both really accurate and nowhere near. For the first experiment, the values for the calorimeter invariable were really imprecise, runing from negative values to ten times greater than the theoretical 15.0 J/°C. This is most likely due to a series of misreckonings and human mistake. In experiment two, the Numberss were far more favourable, with a 19.24 % mistake for our informations and a 14.47 % mistake for the full category. This figure still seems excessively high to warrant the confirmation of Hess’s Law and should likely be redone with more attention in systematically mensurating reactants, but other than that, the experiment was completed good plenty. The consequences for the concluding experiment are besides rather assorted. While some experimental values had merely a 1.25 % mistake, others were grossly erroneous with approximately 65.1 % mistake. The most inaccurate informations was the 1s collected for the reaction of a weak acid and a strong base, which yielded an evidently flawed 300 % mistake. For the consequences that were inaccurate, the beginning of mistake was most likely to due a misreckoning on my portion, perchance in the computation of the theoretical values, or the experimental for that affair. Much more attention must be taken when reiterating this lab, for the possible mistakes are legion. The intent of these three labs were to detect the nature of heat and reactions, which the experiments do instead nicely. The processs described do an first-class occupation depicting the intent of each measure, though they are easy to make falsely. In the terminal, the experiments yielded second-rate consequences, a assorted bag of improbably accurate to merely really incorrect. Thermochemistry is so a instead elusive subject, but these experiments make it much more touchable.

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