### Calorimeter

.. ium hydroxide were chosen. Due to time restraints, our lab instructor told the class to perform only a single run of the reaction at each concentration. The experiment for the weak acid and base were conducted in Calorimeter One, and the experiment for the strong acid and base was conducted in Calorimeter Two. For the initial temperature, our team used the temperature of the solution already in the calorimeter before we added the second.

Our team always placed the acid in the calorimeter first before adding the base. As seen in Tables 4 and 5, all of the results gleaned from the weak and strong acid base experiments were negative. Therefore, the results indicate that these particular acid base reactions released heat into their surroundings. The experiment itself further reinforced our interpretation of the results as both calorimeters became warmer as the reactions progressed. Thus, according to our experiments, reactions involving these two acid base reactions are exothermic processes. The reactions caused an increase in temperature in the surroundings by releasing heat from the system.

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As seen in Table 4, as the concentrations of the weak acid and base increased, so did the heat released by the reaction. One molar solutions gave off the least heat, and the six molar solutions gave off the most heat. Therefore, according to our results, the concentration did affect the amount of heat given off by the reaction in this experiment. The more concentrated the solutions of weak acid and base, the more heat is given off by the reaction. The same general trend is seen in Table 5 with the exception of the one molar solution of strong acid and base. The concentration does have an affect on the amount of heat released by the reaction.

The one molar solution should have been the one that gave off the least heat, and the six molar solution should have given off the most heat. A discrepancy in this experiment is the one molar solution of strong acid and base in Table 5. The ?H was – 85.0 kJ/mol, more than the three and six molar solutions. This number is wrong. The ?H for the reaction of the one molar solution of strong acid and base should be below that of the three and six molar solutions, as concentration does affect the amount of heat released by the reaction.

This errant temperature could be due to a misreading of the thermometer during the actual experiment. This discrepancy might have been solved by conducting more than a single trial of the reaction under each concentration. However, due to time restrictions, our lab instructor gave us permission to conduct only one trial at each concentration and warned us that our results might be a little strange. A general weakness of the acid base reaction experiments is that, again, our team assumed that the density and specific heats of the solutions were the same as pure water. Also, only one run was conducted on each of the molarities.

If more trials had been done, our team would have achieved better results. Additionally, the initial temperature of the base was not factored into our equations. There should not have been much difference, as both solutions should have been at room temperature. However, due to these assumptions, our calculations of ?H for the acid base reactions are slightly inaccurate. As shown in Tables 4 and 5, the weak acids and bases used in the experiment gave off less heat than the stronger acids and bases used.

A weakness of the entire study was the assumptions we made in our calculations of ?H. Our team assumed that the density and specific heats of our solutions were the same as pure water. The density of pure water is 1.00 g/mL; the specific heat of pure water is 4.182 J/g. According to the lab manual, “This is not strictly true, especially as your solutions get more concentrated” (Cooper 106). None of our experiments used pure, distilled water. For convenience, our team utilized tap water, which is water mixed with various minerals and impurities, to conduct all of the experiments. Therefore, our results are not as accurate as they could have been, but under the time restrictions we were placed under, they are good approximations.

Conclusion After all of the experiments and calculations were concluded, the team was able to come to tentative assumptions. The data suggests that the salts (used in the salt and water reactions) were endothermic because they absorbed heat from the surroundings. The acid and base reactions, on the other hand, were exothermic because they released heat during the reaction. By comparing the strong acid and base reactions to the weak acid and base reactions, it can be determined that the strong acids and bases gave off more heat. It was also resolved that the concentrations of the acid and bases did effect the amount of heat that was released. It seemed that by increasing the concentration used in the reaction it caused the reaction to release more heat.

Experimental Procedures Calorimeter Construction: 1. Two Styrofoam cups were obtained 2. A sheet of aluminum was wrapped around one of cups 3. The cup with the aluminum wrapping was placed inside the other cup 4. A square piece of cardboard was cut to function as the lid 5.

A hole was poked through the cardboard to allow the insertion of the thermometer 6. The structure was labeled calorimeter one or two with a black marker; so the two different vessels would not be confused during the experiments 7. Then a second calorimeter was constructed using the same steps as seen above Finding the Heat Capacity of Each Calorimeter: 1. 50 mL of water were measured using a graduated cylinder and put into a calorimeter 2. The temperature was recorded every 30 seconds using the thermometer. This was done for two to five minutes to make sure the temperature had stablized.

3. The temperature of the cold water was recorded as the initial temperature. The thermometer should not be removed. 4. 50 mL of water was again measured into a beaker using a graduated cylinder 5.

The beaker was placed on a hot plate and heated for ten minutes 6. It was then removed and the intial temperature was recorded using a different thermometer 7. This procedure should be preformed twice on both calorimeters Solution of Salts in Water: 1. 50 mL of water was measured in a graduated cylinder and put into a calorimeter. 2.

The temperature was measured and recorded as the initial temperature of the water. 3. 5 grams of salt was measured 4. The room temperature was measured and recorded as the initial temperature of the salt 5. The salt was then poured into the calorimeter and the temperature was recorded immediately 6. The temperature was closely monitored and recorded every ten seconds until a rise in temperature was observed 7.

This procedure was to be preformed three times with each of the chosen salts Precipitation Reactions Ten milliliters of sodium chloride solution were measured in a graduated cylinder and put into Calorimeter Two. The initial temperature of the sodium chloride solution was recorded. Ten milliliters of silver nitrate were measured into a graduated cylinder and put into the calorimeter. The temperature was immediately recorded. The mixture was monitored, and the temperature was recorded every ten seconds until it began to decrease.

The mixture was properly disposed of. The above procedure was repeated two more times. All results were observed and recorded. Ten milliliters of barium chloride solution were measured in a graduated cylinder and put into the same calorimeter. The initial temperature of the barium chloride solution was recorded.

Ten milliliters of sodium sulfate were measured into a graduated cylinder and added into the calorimeter. The temperature was immediately recorded. The mixture was monitored, and the temperature was recorded every ten seconds until it began to decrease. The mixture was properly disposed of. The above procedure was repeated two more times.

All results were observed and recorded. Equations in the lab manual were used to calculate the ?H for precipitation reaction. Acid Base Reactions Weak Acid and Base Twenty milliliters of 1 M acetic acid were measured in a graduated cylinder and put into Calorimeter One. The initial temperature of the acetic acid sodium chloride solution was recorded. Twenty milliliters of ammonium hydroxide were measured into a graduated cylinder and added into the calorimeter. The temperature was immediately recorded. The solution was monitored, and the temperature was recorded every ten seconds until it began to decrease. The mixture was properly disposed of. The above procedure was repeated two more times using 3 M and 6 M concentrations of acetic acid and ammonium hydroxide.

All results were observed and recorded. The solutions were properly disposed of, and all results were observed and recorded. Equations in the lab manual were used to calculate the ?H for each experiment. Strong Acid and Base Twenty milliliters of 1 M hydrochloric acid were measured in a graduated cylinder and put into Calorimeter Two. The initial temperature of the hydrochloric acid was recorded. Twenty milliliters of sodium hydroxide were measured into a graduated cylinder and added into the calorimeter.

The temperature was immediately recorded. The solution was monitored, and the temperature was recorded every ten seconds until it began to decrease. The solution was properly disposed of. The above procedure was repeated two more times using 3 M and 6 M concentrations of hydrochloric acid and sodium hydroxide. All results were observed and recorded, and the solutions were properly disposed of.

Equations in the lab manual were used to calculate the ?H for each experiment. Science Essays.

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