Determination of Thermodynamic Values for
the Dissolution of Borax in Water
Goal:
To experimentally determine the following thermodynamic values for the
dissolution of Borax: Enthalpy (ΔH0), Entropy (ΔS0), and Gibb’s Energy (ΔG0). 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 only in their water content: Anhydrous borax (Na2B4O7),
Borax pentahydrate (Na2B4O7 Ÿ5H2O), Borax decahydrate (Na2B4O7 Ÿ10H2O).
Borax decahydrate is the form found in most grocery stores, and the formula is
usually written as Na2B4O7 Ÿ10H2O. However, it would better indicate the
physical structure by writing the formula as Na2[B4O5(OH)4] Ÿ8H2O, since borax
contains the Borate ion, [B4O5(OH)4]2-. In this structure, there are two fourcoordinate boron atoms (two BO4 tetrahedra) and two three-coordinate boron
atoms (two BO3 triangles), as shown in the diagram below.
1 Borax, Na2[B4O5(OH)4] Ÿ8H2O, dissolves slightly in water to give sodium ions and
borate ions, as shown in the following reaction:
Na2[B4O5(OH)4] Ÿ8H2O(s) ⇔ 2Na+(aq) + [B4O5(OH)4]2-(aq) + 8H2O(l)
Equation 1: Borax Dissolved in Water The heating of the solution causes a shift in the reaction to the right with a
corresponding increase in the Equilibrium constant, Ksp. Temperature affects the
molar solubility of most salts. In the case of borax, the solubility is 2.01g/100mL
at 0°C, about 6.3g/100mL at room temperature, and about 170g/100mL at 100°C.
The Equilibrium Constant expression for the above reaction is:
Kc = [Na+]2 [B4O5 (OH)42-]
Equation 2: Equilibrium Constant for Borax in Water
2 Since there are two moles of sodium ions produced for every mole of the borate
ions, Equation 2 can be rewritten in terms of borate concentration as follows:
[Na+] = 2 x [B4O5 (OH)42-]
and so
Kc = {2 [B4O5 (OH)42-]}2 [B4O5 (OH)42-]
or
Kc = 4 {[B4O5 (OH)42-]}2 [B4O5 (OH)42-]
or
Kc = 4 [B4O5 (OH)42-]3
Equation 3: Equilibrium Constant for Borax in Water
The calculation of the Equilibrium Constant, Ksp, is dependent only on 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
hydrochloric acid (HCl) solution according to the following equation:
-
B4O5(OH)42-(aq) + 2 HCl(aq) + 3H2O(l) ⇒ 4 H3BO3(aq) + Cl (aq)
Equation 4: Neutralization of Borate Ion with HCl The HCl is standardized by titration with a Primary Standard, to determine its
concentration. The standard to be used is Sodium Carbonate, Na2CO3, with
Bromocresol green as an indicator. The neutralization reaction is:
Na2CO3(aq) + 2HCl(aq) ⇒ 2 NaCl(aq) + CO2(g) + 3H2O(l)
Equation 5: Neutralization of Sodium Carbonate with Hydrochloric Acid
3 The dissolving of Borax in water is an endothermic reaction; therefore, the
addition of heat will cause a shift of the reaction, 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, ΔG0, by the equation:
ΔG0 = − RT lnK
Equation 6: Relationship of Gibbs Free Energy to
Equilibrium constant where R is the ideal gas constant, 8.314 J/moleŸK, T is the absolute temperature
in Kelvin, and K is the Equilibrium constant for the dissolution of the borax in
water. (See your text, Moore,Stanitski,Jurs, Chapter 18, pages 864-873.)
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 the
change in Enthalpy (ΔH0), the change in Entropy (ΔS0), and the change in Gibb’s
Energy (ΔG0). In your text the relationships between Gibbs Energy, Enthalpy,
Entropy and Temperature are discussed, and the following state function is
explained.
ΔG0 = ΔH0 − T ΔS0
Equation 7: Relationship of thermodynamic values Combining Equations 6 and 7 gives:
− RT lnK = ΔH0 − T ΔS0
which can be rearranged
lnK = −ΔH0/RT + ΔS0/R
or
lnK = (−ΔH0/R) 1/T + ΔS0/R
which is in the form of a linear equation
y = mx + b
4 From the titration data collected for borax solutions at different temperatures, a
straight-line plot of lnK versus 1/T can be made. From a linear fit to the data
points of this graph, the slope equals −ΔH0/R and the y-intercept equals ΔS0/R.
From these values, ΔH0 and ΔS0 can be calculated, and using Equation 7, ΔG0 at
several temperatures can be calculated. Literature value for ∆H0 and ∆S0 for the
dissolution of borax in water are 110 kJ and 380 J.K.
Pre-Lab Assignment:
In your lab notebook, prepare the following information.
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
Material Safety Data Sheets (MSDS).
0
0
3. Give brief definitions of change in 0Enthalpy (ΔH ), change in Entropy (ΔS ),
and change in Gibb’s Energy (ΔG ).
4. After reading the experiment, but before carrying out the procedures,
predict whether ΔH0 and ΔS0 will be positive or negative. Explain
your reasoning.
Chemicals
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Borax decahydrate (Na2B4O7 Ÿ10H2O)
Hydrochloric acid, approx. 0.1M (HCl)
Sodium carbonate (Na2CO3)
Bromocresol green indicator
Equipment and Supplies
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Buret
10mL volumetric pipet
pipet filler
stirrer and stir bar
digital thermometer
ice
Procedure:
A. Preparation of three saturated solutions of Borax at three
different temperatures.
1. Prepare a “room temperature” borax solution by adding about 10g of borax
and a stir bar to 100 mL of deionized water in an Erlenmeyer flask. Place on
a magnetic stirrer and stir for at least 30 minutes, to allow the solution to
5 reach equilibrium. Since this is to be a saturated solution, there must be
un-dissolved borax in the bottom of the flask.
2. Prepare an "ice bath" borax solution by adding about 5g of borax and a stir
bar to 100 mL of deionized water in an Erlenmeyer flask. Place the flask
in a large beaker containing a mixture of ice and water, on a magnetic
stirrer. Stir the solution for at least 30 minutes. Make sure that some ice is
kept in the bath at all times.
3. Prepare a “cool water bath” borax solution by adding about 10 g of borax
and a stir bar to 100 mL of deionized water in an Erlenmeyer flask. Place
the flask in a large beaker containing cool water, on a magnetic stirrer. Stir
the solution for at least 30 minutes. You may want to periodically add a
little ice to keep the water bath at a constant temperature somewhere
between room temperature and the ice bath. Do not add any ice for the
last few minutes.
While the borax solutions are stirring, do Part B. Standardization of the
HCl solution.
B. Standardization of a dilute HCl solution
1. Obtain a bottle of 0.1 M HCl. Rinse and then fill the buret with the HCl
solution. For the experimental calculations, the concentration needs to be
known to three significant figures; therefore, you will be determining the exact
concentration by titration against a primary standard.
2. Weigh out three samples, between 0.1 and 0.2 grams, of the Primary
Standard Na2CO3 into three clean labeled Erlenmeyer flasks. Record the
exact mass of standard for each sample (+ 0.001g). Add approximately
50mL of deionized water and swirl to dissolve the solid. Add four or five
drops of Bromocresol green indicator to each flask (this should give an
initial blue color).
3. Titrate each Na2CO3 sample with the HCl to a yellow color (no green tint).
Calculate the number of moles of base (Na2CO3) present in each flask, and
then find the molarity of the acid for each titration (to 3 significant figures).
The three values should be close (within 5%). If the difference is >5%, you
should perform additional titrations until you have good agreement for the
determination of the concentration of the dilute HCl. Calculate the mean of
6 the concentration. It is to be reported along with the calculated deviation
from the mean (See Worked Example of Deviation from the Mean)
C. Titration of samples taken from the three borax solutions
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 the "ice bath" borax solution and
carefully pipet three 10.00mL samples of the solution into three separate
clean Erlenmeyer flasks. Be careful not to pipet any of the solid borax from
the bottom of the flask. To each sample add approximately 20 mL of
deionized water and 4 drops of Bromocresol green, then titrate with the
standardized HCl to a yellow end point (no green tint). The volumes of HCl
added should agree (± 5%). If not, titrate another 10mL sample.
3. Measure and record the temperature of the "cool water bath" borax solution
and carefully pipet three 10.00mL samples of this solution into three separate
clean Erlenmeyer flasks. To each sample add approximately 20 mL of
deionized water and 4 drops of Bromocresol green and titrate as in step 2.
4. Measure and record the temperature of the "room temperature" borax
solution and pipet three 10.00mL samples into Erlenmeyer flasks with water
and indicator, and titrate as in step 2.
5. Based on your titration volume of HCl, calculate the average concentration
of the Borate ion at each of the three temperatures. Then using Equation
3, calculate the Ksp for each temperature (a data table of Ksp, ln Ksp, T in
Kelvin and 1/T is recommended). Using Excel® or similar plotting
program, create a graph of ln Ksp versus 1/T and include the equation for
the linear fit (trendline) on the graph. From this equation for the line, you
can determine the experimental values for ΔH0 and ΔS0. Then, using
Equation 7, ΔG0 at several temperatures can be calculated.
7 Lab Report Rubric: Determination of Thermodynamic Values for the Dissolution of Borax in Water Introduction: (5 pts.) Data, Results, Evi d enc e:
' Scientific data that supports the claim should be presented in understandable tables, graphical plots, examples of calculations The data needs to be appropriate and sufficient to support the claim. ' (25 pts. total) Introduction: Provide background information to put the experiment in context, summarize the purpose of the experiment. Include definitions for change in Enthalpy, Entropy and Gibbs Free Energy. Procedure: You can refer to the procedure on the ICN page and provide the reference of the web address. Be sure to note any changes made to the experimental procedure. Results: Organize and present the data you collected that supports your claim including all observations, tables of data, and sample calculations. Report your standardized HCl concentration ± deviation from the Mean. Report your experimental values for borate ion concentration at each temperature, ΔH0, ΔS0, and ΔG0.
Analysis of Evidence (Reasoning) Scientific explanations that use evidence and appropriate chemistry concepts to construct claims. (30 pts. total) Discussion: Explain how the evidence you presented supports your claim. Claim (s) Statement(s) derived from evidence, using scientific reasoning. (15 pts. total) Summary of Claims: State clearly the major conclusions or claims that answer the question: What experimental values for ΔH0, ΔS0, and ΔG0 were determined? How does changing the temperature of the solution change the solubility and the equilibrium constant? Explain chemically why a dilute strong acid such as HCl can be standardized using the salt Sodium carbonate. Explain all the calculations you did, and compare your values for ΔH0 and ΔS0 to literature values, give percent error.
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