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# Experimental Molar Enthalpy of Neutralization for Sodium Hydroxide Solution

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Introduction

ï»¿Yundi Wang October 22, 2012 Molar Enthalpy of Neutralization ________________ 1. For information regarding the problem, prediction, materials and procedure, please see attached Measurements Table for Molar Enthalpy of Neutralization for Sodium Hydroxide Solution Substance Instrument Used Measurement Sodium hydroxide 100mL graduated cylinder (±0.2mL) 50.0mL Sulfuric acid 100mL graduated cylinder (±0.2mL) 30.0mL Temperature of sodium hydroxide solution Thermometer (±0.2ËC) 26.0ËC Temperature of the sulfuric acid Thermometer (±0.2ËC) 24.0ËC Final temperature reached by solution Thermometer (±0.2ËC) 34.5ËC Initial and Final Temperatures of Solutions Temperature of sodium hydroxide solution (±0.2ËC) 26.0ËC Temperature of the sulfuric acid (±0.2ËC) 24.0ËC Final temperature reached by solution (±0.2ËC) 34.5ËC 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 Solution 1. The experimental molar enthalpy of neutralization for sodium hydroxide solution was found to be -64±3.3KJ/mol. Calculation of Uncertainties 34.5±0.2ËC – 25.0±0.2ËC =9.5±0.2ËC 50±0.2mL + 30±0.2mL =80±0.2mL 9.5 ± 0.4ËC = 4.210…% 80 ± 0.4mL = 0.5% 50 ± 0.2mL = 0.4% =5.11…% =5.1% Percent Difference Conclusion Through a pre-lab calculation the amount of sulfuric acid solution needed was found to be 30.0m±0.2mL. ...read more.

Middle

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. ...read more.

Conclusion

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. ...read more.

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