Thermodynamics Of Borax LAB REPORT

Topics: Thermodynamics, Entropy, Energy Pages: 8 (1293 words) Published: April 13, 2015


Heats Effect on Borax
By: Alexis H. Prince
Department of Chemistry; Coastal Carolina University, Conway, SC 29526 April 13, 2015

Introduction
Borax has many uses, whether it’s being used as an antiseptic, helping to cure people, getting rid of pests, or even assisting fruit trees in their growth. It is actually the most important source of the element boron, and has been used for years as a “water softening agent.” Borax was found forming in saline lakes, or one may find it in Boron, California. Historically, “the first Borax specimens came from several dry lake deposits in Tibet” (The Mineral Borax). This experiment was conducted to determine the standard entropy and enthalpy of the dissolving reaction of borax in water. The thermodynamic properties of the reaction helped to determine the change in heat and spontaneity within the system. Entropy is said to be the tendency for the universe to move towards disorder. If the value of entropy is positive, then the amount of disorder would increase within the system, causing the reaction to occur spontaneously. However, if the value of entropy is negative, the amount of disorder would decrease, this could cause a spontaneous or non-spontaneous reaction, depending on the value of enthalpy. Enthalpy is the total energy within a system in relation to work and heat. If the value of enthalpy was negative, then the reaction is exothermic. But, if the value of enthalpy was positive, then the reaction will be endothermic. A doctor by the name of J. Gibbs came up with an equation, which combined contributions from enthalpy and entropy. This equation provided a way to measure the energy content within a system which allows one to evaluate the spontaneity of a reaction. If there is a lot of stored heat energy, then the substance has a lot of free energy. But, the more disorder and disruption the substance has, the less free energy it has.

The borax was tested to see what would happen when different temperatures of heat were applied. Thermodynamics then came into play when trying to explain what was happening with the borax solution as the temperature steadily increased. The higher the temperature was, the more disorder there was within the system. The more the temperature increased, the further away it was from reaching its equilibrium. The standard entropy and enthalpy of the dissolving reaction borax in water was later determined. The properties of thermodynamics in this reaction helped to determine the change in heat and spontaneity within the system. Experimental Details

To begin this experiment, 15-20 grams of borax was collected and added to a 100 milliliter beaker which contained 75-85 milliliters of water. A stirring bar was used to allow the borax to dissolve and reach saturation. Once the stirrer was turned off, the solution was given time to allow any solid to settle to the bottom so that the aliquot retrieved was only of the saturated solution, (if less energy were used, temperatures would decrease, causing larger amounts of solid to form, enabling one to retrieve aliquots). 5-7 milliliters were removed from the beaker and placed into a clean and dry 10 milliliter graduated cylinder, the volume and temperature of the aliquot was recorded. Immediately, the aliquot was transferred to a 125 milliliter Erlenmeyer flask, warm distilled water was used to ensure all of the borax had been transferred out of the graduated cylinder and into the flask. Then a few drops of bromocresol green indicator was added to the solution to indicate the endpoint during titration. A standardized hydrochloric acid solution with a 0.5082 molarity was titrated with the borax solution. Once the borax solution reached its endpoint, this was indicated by a change in color to a light green. If a yellow color appeared it meant the endpoint was surpassed and needed to be noted. After this, the final volume and dispensed volume was calculated and recorded.

While the titration was...

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