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  1. Creating a Saturated Solution

    In this experiment, you will create a saturated solution of sodium tetraborate decahydrate, also called borax. In water, borax dissociates into two sodium cations and one tetraborate anion. When an aqueous borax solution is saturated, it means that it contains the maximum amount of dissolved solute, which is borax, for that volume of solvent, which is water. Any additional solid won't appear to dissolve.

    Since solubility is temperature-dependent, each group will create a saturated solution at a different temperature. More borax will dissolve at higher temperatures, resulting in a solution with a higher borax concentration.

    • Before starting the experiment, put on the necessary personal protective equipment, including a lab coat, chemical splash goggles, and gloves.
    • The instructor will assign you a temperature to use during the experiment. Based on this temperature, determine the amount of borax you will need to create a saturated solution using the table.
      Temperature (°C) Amount of borax (g)
      10 4
      20 6
      30 8
      40 13
      50 21
    • Measure the necessary amount of borax for your assigned temperature. For example, 8 g of borax is needed for a temperature of 30 °C. Record the mass in your lab notebook. Create this table to help you keep track of your data.

      Table 1: Temperature and Volume Data

      Assigned temperature (°C)
      Borax mass (g)
      Trial Solution temperature (°C) Volume (mL) of 0.5 M HCl added
      Click Here to download Table 1
    • Label a 150-mL beaker ‘borax solution’, and transfer the borax into it. Measure 50 mL of deionized water and pour it into the beaker. Add a magnetic stir bar to the solution, set the beaker on the hotplate, and turn on the stir setting.
    • Use a digital thermometer clamp to hold a digital thermometer inside the beaker and above the stir bar. Heat the beaker to about 10 to 20 °C higher than the assigned temperature. Note: Students assigned the 10 °C temperature will let the borax dissolve at room temperature.
    • For all assigned temperatures, add 200 mL of deionized water to a 400-mL beaker. Then, place a glass thermometer in the beaker of water. Set the beaker on the hot plate and heat the water to 20 °C higher than your assigned temperature. Note: Students assigned the 10 °C temperature should use room temperature water.
    • Label the five 250-mL Erlenmeyer flasks as ‘trials 1 through 5’. Once the borax has fully dissolved in the solution and it appears homogeneous, remove the beaker from the stir plate and set it on the benchtop to cool to the assigned temperature.
    • Once the temperature has been reached, place the beaker on a stack of paper towels to act as an insulator and maintain a constant temperature. Note: Students assigned to the 10 °C temperature should place their beaker containing borax solution into a larger beaker filled with ice and water.
    • Label another 400-mL beaker as ‘waste’, and pipette 5 mL of the hot water into your waste beaker. Repeat this several times to warm up the pipette and prevent it from getting clogged as the borax cools.
    • Record the exact temperature of the borax solution in your notebook. Then, pipette 5 mL of the saturated borax solution from the top of the beaker and dispense it into the first 250-mL Erlenmeyer flask. Do not allow the pipette to come into contact with the crystallizing borax at the bottom of the beaker.
    • Pipette 5 mL of the hot water into the Erlenmeyer flask containing the borax solution. Repeat this for each of the four remaining flasks, adding 5 mL of borax solution and 5 mL of water. Remember to keep an eye on the temperature of your borax solution, as it will cool on the benchtop. Note: The temperature of the saturated solution transferred into each of the five flasks should be within 2 – 3 °C of each other. Students working with the higher temperatures may need to reheat their solution before proceeding.
  2. Titration of Borax Solution

    Now you will determine how much borax has dissolved in your saturated solution. Recall the chemical reaction showing how borax dissociates in water, forming the tetraborate ion. Since the tetraborate ion is a base, it will react with acid following a neutralization reaction.

    When the amount of acid is twice the amount of tetraborate, the solution is neutralized. To do this, we will slowly dispense HCl into the borax solution until it is neutralized, meaning that the acid and base react to form water and salt and a neutral pH. We'll use the pH indicator bromocresol green to let us know when the solution is neutralized, as it turns from blue to pale greenish-yellow when the pH is neutral.

    • To begin the titration, first, measure 40 mL of deionized water and add it to one of the Erlenmeyer flasks. Swirl the flask to make sure all of the borax is in solution. Repeat this for the other four flasks.
    • Add 2 - 3 drops of 0.1% bromocresol green to each flask.
    • Clip a burette clamp on a ring stand and secure a burette in the clamp while making sure that the burette is vertical and as straight as possible. Make sure that the stopcock on the burette is closed.
    • Place a funnel in the open end of the burette and fill the burette with deionized water. Set your waste beaker under the burette and open the stopcock to let all of the water rinse out of the burette. Then, close the stopcock.
    • Obtain 0.5 M HCl from your instructor and pour 10 mL into the burette. Use the markings on the side of the burette as a guide. Open the stopcock on the burette to allow all of the HCl to drain into the waste beaker
    • Close the stopcock again, and then fill the burette with 50 mL of the HCl. Open the stopcock slightly to allow the liquid to fill the tip of the burette and remove any bubbles.
    • Record the initial volume of HCl in your lab notebook.
    • For the first trial, titrate the HCl into the borax solution in increments of 1 mL. Gently swirl the flask after each addition to make sure that the solution is well mixed. Note: At the endpoint of the titration, the indicator will turn the solution from light blue to a pale greenish-yellow color. If your solution turns a dark yellow color, this indicates that you have passed the endpoint.
    • When you have reached the endpoint, record the volume of HCl remaining in the burette.
    • Refill the burette to the 50-mL mark with more HCl and repeat the titration for all of the other flasks of borax solution. Make sure to record the final volumes of HCl used for each titration.
    • To clean up from the experiment, place the 400-mL waste beaker under the burette spout and open the burette to drain the remaining HCl into it.
    • Fill the burette with deionized water and allow it to rinse through the burette.
    • Use your remaining borax solution to neutralize the acid in your waste beaker. Swirl the beaker until it stops bubbling.
    • Add baking soda to the waste beaker and swirl the solution. Continue adding baking soda and swirling the solution until it stops bubbling. Note: You may not observe bubbling because the borax does most of the neutralization.
    • Wash the contents of the waste beaker down the sink with copious amounts of water. The flasks containing borax and HCl from your titrations are neutralized, so they can also be poured down the sink.
  3. Results
    • First, determine the reaction constant Ksp for borax dissociation in water. Determine the average temperature in Celsius and the average volume of 0.5 M HCl used to reach the endpoint of the titration.

      Table 2: Determining ΔG, ΔH, and ΔS

      Assigned temperature (°C) Tavg (°C) Vavg (mL) Tavg (K) 1/Tavg (K) Moles of 0.5 M HCl Molarity of Na2[B4O5(OH)4] Ksp lnKsp ΔG (kJ/mol)
      10 °C
      20 °C
      30 °C
      40 °C
      50 °C
      Click Here to download Table 2
    • Obtain data from other groups so that you have an average HCl volume for each temperature.
    • The solubility constant is determined from the stoichiometry of the dissociation reaction of borax and water. Since the reactant is solid, it is not included in the expression. Because two sodium cations are formed for every tetraborate, it is assumed that the concentration of sodium is equal to twice the concentration of tetraborate. Thus, the expression can be simplified as shown: Ksp = 4[[B4O5(OH)4]2-]3
    • The concentration of tetraborate can be calculated from the volume of HCl used to reach the endpoint of the titration and neutralize the base. The moles of HCl equal the volume of HCl times the molarity of the solution. Since two moles of HCl are needed to neutralize one mole of tetraborate, the moles of tetraborate can easily be calculated.
    • Calculate the concentration of tetraborate (based on 5 mL volume of saturated solution) and determine the reaction constant. Repeat the calculation for each temperature and compare your results.
    • Use the titration data to calculate the Gibbs free energy, ΔG, at each temperature value to determine whether the dissociation of borax is a spontaneous reaction. Recall the ΔG expression, where R is the gas constant, T is the temperature of the solution in Kelvin, and Ksp is our reaction constant. Using our reaction constants and the corresponding temperature, calculate ΔG for each temperature.
    • Remember that if ΔG is positive, the reaction is not spontaneous, meaning that energy needs to be put into the reaction for it to proceed. However, if ΔG is negative, the reaction proceeds spontaneously. In general, this reaction is not spontaneous at low temperatures, but it is spontaneous at higher temperatures. This supports the theory that borax prefers the salt crystal structure form at room temperature and lower, but it prefers to go into solution after a certain temperature is reached. Use Ksp to calculate ΔH and ΔS.
    • Plot the natural log of Ksp as a function of 1/T reported in Kelvin. The slope of this line equals -ΔH/R, so calculate ΔH. The intercept of the line equals ΔS/R. Determine ΔS.
    • Since the change in enthalpy of this reaction is about 90 kJ/mol and is positive, the reaction is endothermic, meaning that it absorbs energy. The change in entropy is positive and is about 290 J/mol·K, which indicates the favorable production of disorder. This is expected as the crystal structure of the salt breaks down.

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