Determination of Thermodynamic Values for the Dissolution of Borax in Water. Goal: To experimentally determine the following Thermodynamic Values for the Dissolution of Borax: 1. 2. 3. Enthalpy, ΔΗ ο Entropy, ΔSο Gibb's Energy, ΔGο . Abstract: Borax or sodium borate is a naturally occurring mineral composed of Sodium, Boron, Oxygen and water. Vast deposits are found in the Southwestern US. The term borax is often used for a number of closely related minerals or chemical compounds that differ in their water content: 1. Anhydrous borax (Na2B4O7) 2. Borax pentahydrate (Na2B4O7 * 5H2O) 3. Borax decahydrate (Na2B4O7 * 10H2O) Borax decahydrate is the form found in most grocery stores and is generally . described as Na2B4O7 10H2O. However, it is better formulated as Na2[B4O5(OH)4] * 8H2O, since borax contains the Borate ion, [B4O5(OH)4]2−. In this structure, there are two four-coordinate boron atoms (two BO4 tetrahedra) and two three-coordinate boron atoms (two BO3 triangles). (As shown below.) Borax, Na2[B4O5(OH)4]*8H2O, dissolves slightly in water to give sodium ions, a borate ion, and water according as shown in the following reaction. The heating of the solution causes a shift of the reaction to the right with a corresponding increase in the Equilibrium constant, Ksp. Temperature affects the molar solubility of most salts: Solubility of borax @ 0oC = 2.01g/100mL Solubility of borax @ r.t. ≈ 6.3g/100mL Solubility of borax @ 100oC = 170g/100mL Notice that the concentration of the Sodium ions in Equation 1, above, is twice that of the Borate ions. The Equilibrium Constant expression for the above reaction is: Since there are two moles of sodium ions produced for every mole of the borate ions Equation 2 can be rewritten as follows: The calculation of the Equilibrium Constant, Ksp, is dependent only on evaluation of the concentration of the Borate ion, [B4O5(OH)4]2−. Since the Borate ion is a weak base, its concentration can be determined by titration with a standardized acid (HCl) solution according to the following equation: The HCl is standardized by titration with a Primary Standard Sodium carbonate, Na2CO3, using Bromocresol green as an indicator. The dissolving of Borax in water is an endothermic reaction, therefore the addition of heat will cause a shift of the reaction shown in Equation 1 to the right to increase the ionic concentration and the Equilibrium Constant value. (This means that at a warmer temperature more of the solid will dissolve and when the solution comes to equilibrium the value of Ksp will be larger.) The Ksp of the borax solution is related to the change in Gibbs Free Energy, ΔGo, by the following equation: ΔGo = - RT ln K R is the ideal-gas constant, 8.314 J/mol-K, T is the absolute temperature in Kelvin and K is the Equilibrium constant for the dissolution of the borax in water. The above equation allows the calculation of Gibbs Free Energy, ΔGo, for a solution if the equilibrium constant of the solution and the temperature of the solution are known. A series of titrations can be carried out in order to calculate Ksp of Borax Dissolution at different temperatures. This data will allow the determination of values for several Thermodynamic State Functions for the Dissolution of Borax. They are ΔH, the change in Enthalpy, ΔS, the change in Entropy and ΔG, the change in Gibbs Energy. In Chapter 7 (pages 256-806) of Shultz, (your text) the relationship between Gibbs Energy, Enthalpy, Entropy and Temperature is discussed and the following state function is explained. ΔGo = ΔHo - TΔSo Combining these two equations gives: Equation [7.19] , page 257 - RT ln K = ΔHo - TΔSo which can be rearranged to ln K = -ΔHo/ RT + ΔSo/ R or The above equation is in the form of the equation for a straight line, y=mx+b. From the data collected from the titration of different borax solutions, collected at different temperatures, a straight-line plot of lnK vs 1/T can be made. From this curve the slope (-ΔHo/ R) and the intercept ΔSo/ R can be determined and values for the state functions ΔHo and ΔSo can be calculated. From ΔHo and ΔSo, Gibbs Energy, ΔGo at several temperatures can be calculated. Literature values for enthalpy and entropy of the dissolution of borax in water are 110 kJ/mol and 380 J/mol.K respectively. Prelab Assignment: The following should be answered in your lab notebook before you come to lab. 1. A brief (2-3 sentence) introduction to the lab. 2. A table of safety information including the chemicals used in the lab and any safety handling precautions. This information can be obtained from the MSDS safety sheets. 3. Briefly give definitions for ΔH (change in Enthalpy), ΔS (change in Entropy), and ΔG (change in Gibbs Free Energy). 4. After reading the experiment, but before carrying out the procedures, predict whither ΔH (change in Enthalpy) and ΔS (change in Entropy) are positive and/or negative. (Be sure to explain your reasoning.) 5. Explain chemically why the dilute HCl solution, in Procedure: B, can be standardized using the salt Sodium carbonate, Na2CO3, as a primary standard. Give the information to your TA at the beginning of the lab. Chemicals • • • • Borax decahydrate (Na2B4O7_10H2O) Hydrochloric acid, 0.1M HCl Sodium carbonate ,Na2CO3, anhydrous Bromocresol green indicator Equipment and Supplies • • • • • • • Buret 10 mL Volumetric Pipet Pipet Filler Stir Bar Thermometer Ice Plastic Tray Procedure: Sheets of graph paper in pdf form A. Preparation of three saturated solutions of Borax at different temperatures. Prepare solutions of borax at three different temperatures and allow them to come to equilibrium. (All measurements that are to be used in the following calculations need to be made to 3 significant figures.) Specific temperatures are important, but it is very important to know exactly the temperature of the solution. Prepare three solutions: 1. A "room temperature" borax solution by adding about 10g of borax and a stir bar to 100 mL of distilled water in an Erlenmeyer flask. Place on a magnetic stirrer and allow the solution to stir for at least 30 minutes. 2. An "ice bath" borax solution by adding 5 g borax to 100 mL of distilled water and a stir bar in an Erlenmeyer flask. Place the flask in a small plastic dish, add a mixture of ice and water around the beaker and stir for at least 30 minutes. 3. A "cool water bath" borax solution be adding 10 g of borax to 100 ml of distilled water and a stir bar Erlenmeyer flask. Place the flask in a cool water bath on a magnetic stirrer. Periodically add a little ice and allow it to melt stir for 30 minutes and do not add any ice for the last few minutes. (Try to keep the water bath at a constant temperature that is about half way between the other two solution temperatures.) Be sure to record the temperature of this cooling bath. B. Preparation and standardization of dilute HCl solution. Obtain about 50 mL of 1 M HCl. Add it to a clean 500 mL plastic bottle. Add distilled water until the bottle is at almost full. Cap and mix the solution. At this point the concentration of the HCl solution is about 0.1 M but for the experimental calculations the concentration needs to be known to three significant figures Weigh out samples, between 0.1 and 0.2 grams, of the Primary Standard Na2CO3 to the closest milligram (+ 0.001g) into three clean Erlenmeyer flasks, add 50mL of deionized water and swirl to dissolve the solid. Add four or five drops of the Bromocresol green indicator to each sample. This should give an initial blue color. Titrate each sample with the dilute HCl to a yellow color. The three values should be close (+ 5%). If not perform other titrations until you have three that are close. C. Titration of samples from the three borax solutions. Prepare a data table for each borax solution. 1. Stop the stirring of the Borax solutions after 30 minutes and let them sit undisturbed for 5 minutes to allow the remaining solid borax to settle. 2. Measure and record the temperature of "ice bath" solution and carefully pipet three 10.00mL samples of the solution into three clean Erlenmeyer flasks. Be sure to not pipet any of the solid borax from the bottom of the flask. To each sample add 20 mL of distilled water, 4 drops of Bromocresol green and titrate with the standardized HCl to a yellow end point (no green tint). The volumes of the titration should agree (+ 5%). If not titrate another sample. 3. Measure and record the temperature of "water bath" solution and carefully pipet three 10.00mL samples of this solution into three clean Erlenmeyer flasks. Be sure to not pipet any of the solid borax from the bottom of the flask. To each sample add 20 mL of distilled water, 4 drops of Bromocresol green and titrate with the standardized HCl to a yellow end point (no green tint). The volumes of the titration should agree. If not titrate another sample. 4. Measure and record the temperature of "room temperature" solution and carefully pipet three 10.00mL samples of this solution into three clean Erlenmeyer flasks. Be sure to not pipet any of the solid borax from the bottom of the flask. To each sample add 20 mL of distilled water, 4 drops of Bromocresol green and titrate with the standardized HCl to a yellow end point (no green tint). The volumes of the titration should agree. If not titrate another sample. Post lab Assignment: Sheets of graph paper in pdf form Include the following information in a lab report to give to your TA. (All data to be clearly arranged in labeled tables. Be sure to include examples, with units, of each calculation.) 1. Concentration of the Standardized HCl solution + the deviation from the mean. (See Equation 4. in the Abstract) (For Example: [HCl] = 0.108 + 0.003 M.) (See Worked Example of Deviation from the Mean.) 2. Concentration of the Borate ion at the different temperatures. (See Equation 3. in the Abstract) 3. The calculated values of Ksp at the corresponding temperatures (T) in Kelvin. (See Equation 2. in the Abstract) 4. Values for ln Ksp and 1/T. 5. From a plot of ln Ksp vs. 1/T, calculate experimental values for the following Thermodynamic State Functions for the Dissolution of Borax and compare them to the literature values, (give percentage error): a.) the change in Enthalpy, ΔHo, b.) the change in Entropy, ΔSo, and c.) the change in Gibbs Free Energy, ΔGo, at each solution temperature. Copyright (c) 2012, the ICN Team.