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.