Sorry, but copying text is forbidden on this website!
1. For information regarding the problem, prediction, materials and procedure, please see attached
Measurements Table for Molar Enthalpy of Neutralization for Sodium Hydroxide Solution
100mL graduated cylinder (±0.2mL)
100mL graduated cylinder (±0.2mL)
Temperature of sodium hydroxide solution
Temperature of the sulfuric acid
Final temperature reached by solution
Initial and Final Temperatures of Solutions
Temperature of sodium hydroxide solution (±0.2ËC)
Temperature of the sulfuric acid (±0.2ËC)
Final temperature reached by solution (±0.2ËC)
Neutralization Reaction Taking Place
Pre-Lab Calculations – Volume of Sulfuric Acid Needed
Average Initial Temperature of Solutions Calculation
Experimental Molar Enthalpy of Neutralization for Sodium Hydroxide Solution Calculation
1. The experimental molar enthalpy of neutralization for sodium hydroxide solution was found to be
Calculation of Uncertainties
34.5±0.2ËC – 25.0±0.2ËC
50±0.2mL + 30±0.2mL
9.5 ± 0.4ËC = 4.210…%
80 ± 0.4mL = 0.5%
50 ± 0.2mL = 0.4%
Through a pre-lab calculation the amount of sulfuric acid solution needed was found to be 30.0m±0.2mL. Using this information, a calorimetric lab was conducted to find the molar enthalpy of neutralization for the sodium hydroxide solution. Through molar enthalpy calculations, the experimental molar enthalpy of neutralization for the sodium hydroxide solution was found to be -64.0±3.3KJ/mol; however, the theoretical (actual) molar enthalpy of neutralization for the sodium hydroxide solution is -57KJ/mol. In other words the experimental enthalpy change was -64.0±3.3KJ and the theoretical (actual) enthalpy change was -57KJ. This as a result produced a 12% difference. The various errors will be analyzed in the evaluation.
As discussed earlier in the conclusion, the experimental change in enthalpy is greater than the theoretical (actual) change in enthalpy. This result is quite rare. In general, a typical result for the experimental enthalpy change should yield an outcome lower than the theoretical (actual) value (the reason for this will be discussed later in the conclusion); however, this was not the case in this lab. There are a variety of reasons why the experimental enthalpy change for this lab was greater than the theoretical (actual) enthalpy change.
In general, the main reason for the result seen in this lab is due to the nature of the calorimeter. Due to the fact the calorimeter is an isolated environment there is no possible method to determine when the reaction is complete. As a result, the reaction may have been occurring in a concentrated area. With an increased concentration of reactants in one area, the rate of the reaction increases along with the temperature in the concentrated area. When this heat transfers to the thermometer, it causes an increased change in enthalpy. Normally, the concentration of reactants would be less, as they are not in a concentrated area. This would then cause a lower temperature increase because there is a smaller chance the particles will collide. As a result, the change in enthalpy in a normal situation would be much lower than if the reactants were all concentrated in one area.
In saying that, it is possible within this lab the reactants were concentrated in one area causing the experimental change in enthalpy to be quite large. Because it is impossible to see into the calorimeter to see if the reaction is concentrated or when the reaction is complete the reactants could easily have been concentrated in one area. Furthermore, by not knowing when the reaction is complete, the temperature might be measured too soon or too late causing inaccurate results. In general, because the calorimeter is an isolated environment it results in the experiment having many errors because how the reaction is occurring and when the reaction is finished is unknown. A way to eliminate this error is by inserting an electronic stirring rod to stir the reactants so they do not become concentrated in one area.
Furthermore, another reason contributing to the large enthalpy change is the impurity of the substances used. As a result, because the substances are impure, they could have had a higher concentration of reactants. With a higher concentration of reactants, the reaction rate will increase and there will be a greater reaction than wanted. With a larger reaction at an increased rate, the final temperature of the solutions will spike higher than wanted generating a larger enthalpy change. As a result, this is a reason contributing to the large enthalpy change in this lab; however, this reason is not very significant as the substances cannot be so impure the concentration on the label is extremely different then the concentration found in the bottle (it is illegal to put false information on chemical substances). As a result, the impurity of the substances cannot account for all the errors in this lab. Purifying the substances beforehand can easily eliminate this source of error.
Moving on, there is another reason contributing to the large enthalpy change. The theoretical (actual) value given is obtained at SATP conditions; however, when the following lab was conducted, the conditions were not at SATP. SATP conditions are at 100kPa and 25ËC. The conditions when the lab was conducted were at 101.9kPa and 25ËC.
By increasing the pressure, the reaction rate is increased and more reactions take place. As a result of a larger quantity of reactions occurring at 101.9kPa than at the standard SATP conditions, there will be a greater change in enthalpy at 101.9kPa. This as a result, contributes to the large difference in enthalpy change seen in this lab; however, like the previous reason, this is not a significant factor in increasing the enthalpy change. The pressure differences are not extremely different to cause the enthalpy change to increase to as much as they have in this lab. As a result, this is only a small contributing factor. Conducting this lab at SATP conditions will eliminate this source of error. In general, the main reason for the larger enthalpy change is due to not being able to tell when the reaction is complete and how the reaction is occurring in the calorimeter.
As mentioned earlier, the result in this lab is very rare. This is mainly due to the fact that the Styrofoam calorimeter used to conduct the calorimetric experiment most likely does not provide a perfectly sealed environment. A hole is needed to be made to insert the thermometer. And there were many holes between the lid of the calorimeter and the calorimeter itself. Due to this ineffectiveness of the Styrofoam calorimeter, some of the heat from the reaction would have escaped through the many holes causing a lower final temperature of the reaction and the experimental enthalpy change to be lower than the theoretical (actual) value. As a result, the experimental value is usually lower than the theoretical (actual) value.
Another reason includes the fact that some of the heat released during the reaction would have been transferred to the calorimeter itself instead of transferring to the thermometer. As a result, when the calorimeter and/or glass of the thermometer absorb the heat, it causes the thermometer to absorb less heat than it should. The final temperature will then be lower than it should be causing a lower enthalpy change. Even though this is not a main reason why the experimental molar enthalpy should be lower than the actual molar enthalpy it still contributes to it. As a result, with the combination of these factors the experimental enthalpy change should be lower than the theoretical value because a lot of heat is able to escape into the calorimeter and into the air due to there being holes in the calorimeter.