Lab: Thermodynamics of Saturated Borax Solutions
Background: Borax is commonly called sodium tetraborate decahydrate and has the formula
Na2B4O7•10H2O. However, based upon its chemical behavior as the sodium salt with the
conjugate base of weak boric acid, a more descriptive formula for borax is Na2B4O5(OH)4•8H2O,
also called tincal. Borax dissolves and dissociates in water according to the equation below:
Na2B4O5(OH)4•8H2O(s)  2Na+(aq) + B4O5(OH)42-(aq) + 8H2O(l)
At equilibrium, the solubility-product constant for borax is determined by: Ksp = [Na+(aq)]2
[B4O5(OH)42-(aq)]
The borate ion, B4O5(OH)42-(aq), can accept two protons from a strong acid in an aqueous
solution. Therefore, the molar concentration of the borate anion in a saturated solution can be
determined during a neutralization titration with standardized hydrochloric acid. The
concentration of the sodium ions in the saturated solution will be twice the molar concentration
of the borate ions.
The free energy of the dissolution of borax is related to the equilibrium constant by the equation:
Gº = -RT ln Ksp , where R = 8.314 x 10-3 kJ/mol•K and T is the Kelvin temperature
The free energy is also related to the changes in enthalpy (Hº ) and entropy (Sº) by the
equation:
Gº = Hº - TSº
Setting these two free energy equations equal to each other gives the expression: -RT ln Ksp =
Hº - TSº
Re-arranging and solving for the ln Ksp results in an equation in the form of a linear
relationship, y = mx + b. By plotting a graph of ln Ksp versus 1/T, the slope of the line equals Hº/R and the y-intercept equals Sº/R.
Problem: Can you determine the Ksp values for borax at various temperatures, then use this
information in order to determine the free energy, enthalpy and entropy changes for the
dissolution of solid borax?
Standardization of HCl Procedure:
1.
Dry 2-3 g of the primary standard, anhydrous Na2CO3(s), for several hours in a drying oven
set at 110ºC. Cool the sample in a dessicator.
1.
Obtain 100 mL of ~0.5 M HCl(aq) that was prepared by diluting a sufficient amount of
concentrated 12 M HCl(aq). (Caution: HCl causes skin burns. Neutralize any spills with
baking soda.) Clean a buret and rinse twice with 3 mL portions of the ~0.5 M HCl(aq). Fill
the buret with the HCl. Record the initial volume, using all significant digits
3.
Calculate the approximate mass of anhydrous Na2CO3(s) needed to neutralize 15-20 mL of
~0.5 M HCl(aq). Measure this mass to  0.001 g and transfer to a 125 mL Erlenmeyer flask.
Dissolve the solid in 50 mL of distilled water, then add 1 mL of methyl orange indicator.
4.
Begin the titration by adding HCl to the Na2CO3(aq) until the apparent endpoint of the
indicator is reached. This occurs when the methyl orange changes color from yellow at a pH
> 4.4 to orange at a pH < 3.1. Heat the flask to near boiling to drive off the carbon dioxide
gas. Then add additional drops of HCl until the endpoint is reached and the orange color
persists for 30 seconds. Record the final volume of HCl.
5.
Enter your data for the HCl titration into the class spreadsheet. Calculate the experimental
molarity values for the HCl(aq). The trials should agree within  1%. Perform the Q test to
determine if any suspect molarity values can be excluded from the class average. Calculate
the average M HCl and the standard deviation, then record these values.
Preparation of Saturated Borax Solutions Procedure:
1. Saturated borax solutions must be prepared at six different temperatures, ranging from 0ºC
to 60ºC. The amount of borax that will dissolve before the solution becomes saturated will
increase with increasing temperature. The following chart indicates the approximate mass of
borax needed to completely saturate 100 mL of distilled water at the various temperatures.
Temp ºC
0
20
30
40
50
60
g borax
5
9
12
16
23
34
Prepare one of the saturated solutions, as assigned by your instructor, by adding the mass of
borax to 100 mL of distilled water in a 125 mL Erlenmeyer flask. Place the flask in a 400
mL beaker filled with water heated to ~5ºC higher than your assigned temperature. (For the
0ºC trial, allow the solution to set at room temperature.)
1. After 5 minutes of heating, remove the flask from the water bath, stopper it and shake for 2
minutes. Remove the stopper and return the flask to the water bath. Allow the water bath to
cool down to your assigned temperature and for the borax solution to reach thermal
equilibrium. After another 5 minutes, stopper the flask and shake for 2 minutes again, then
return it to the water bath. There should be some solid remaining that will settle to the
bottom of the flask. Allow the borax to settle until the solution is clear. The solution is now
saturated.
2. Record the exact temperature of the saturated solution. Prepare the borax solution for
analysis by using a volumetric pipet to remove 5.00 mL of the saturated solution and transfer
it to an Erlenmeyer flask. (Do not transfer any of the solid borax!) Rinse the pipet three or
four times with warm distilled water and combine the washings with the sample. Add 1 mL
of bromocresol green indicator to the flask and titrate with the standardized HCl until a
yellow-colored endpoint is reached. (The indicator is yellow at a pH < 4.0 and blue at a pH >
5.6. A green color is present within this pH range.)
3. Complete the HCl titration using samples from each of the six saturated solutions prepared.
Results & Calculations:
1.
2.
3.
4.
5.
6.
Calculate the molar solubility of borax at each of the measured temperatures.
Calculate the solubility-product of borax at each of the measured temperatures.
Plot the natural logarithm of the solubility-product versus the reciprocal Kelvin temperature
for each sample and draw the best-fit straight line through the data points.
Determine the slope of the linear plot and calculate the standard enthalpy of solution for
borax.
Determine the y-intercept (at x = 0) of the linear plot and calculate the standard entropy of
solution for borax.
Calculate the standard free energy change for the solution of borax.
Conclusions:
1.
2.
3.
If the sodium carbonate used for the standardization of HCl was not anhydrous, but rather
contained some water of hydration, explain how this would affect
a) the calculated molarity of the HCl
b) the calculated solubilty-product of borax based upon this “erroneous” molarity of
standardized HCl.
If some solid borax was inadvertently sucked up into the pipet from the saturated solution
and transferred to the Erlenmeyer flask for analysis, explain how this would affect the
calculated solubility-product?
Explain why the value for the standard free energy determined by using the ∆Hº and ∆Sº
values from the plot of the data may be a better method for the analysis than by calculating
its value from the solubility-product, Gº = -RT ln Ksp