Purpose:
The purpose of this lab is to investigate the equilibrium of a saturated
NaCl solution, an acid-base indicator using HCl and NaOH, the complex ion
FeSCN+2, and two different complex cobalt ions by adding "stress" to these
systems by changing the concentration of one of the reactants or products or by
changing the temperature of the system.
Background:
Le Chatelier's Principle states that if an equilibrium system is subjected
to a stress, the system will react to remove the stress. To remove a stress, a
system can only do one of two things: form more products using up reactants,
or reverse the reaction and form more reactants, using up products. In this
experiment you will form several equilibrium systems. Then, by putting
different stresses on the systems, you will observe how equilibrium systems
react to a stress.
Acid-base indicators are large organic molecules that can gain and lose
hydrogen ions to form substances that have different colors. The reaction of the
indicator bromthymol blue can be illustrated as follows:
HIn(aq)  H+(aq) + In-(aq)
Yellow
Blue
In this reaction, HIn is the neutral indicator molecule, and IN- is the indicator
ion after the molecule has lost a hydrogen ion. Equilibrium reactions can easily
be forced to go in either direction.
An equilibrium solution can be formed in solution with the following
ions:
Fe+3(aq) + SCN-(aq)  FeSCN+2(aq)
Colorless Colorless
Red-brown
The iron ion (Fe+3) and the thiocyanate ion (SCN-) are both colorless; however,
the ion that forms from their combination, the FeSCN+2 ion, is a dark redbrown. It is the color of this ion that will indicate how the equilibrium system
is being affected. Phosphate ions have the ability to form complex ions with
Fe+3, which has the same effect as removing Fe+ from solution.
In Part 4, the equilibrium between two different complex ions of cobalt
will be investigated. The reaction is endothermic:
Co(H2O)6+2(aq) + 4Cl-(aq)  CoCl4-2(aq) + 6H2O(l) ∆H = +50kJ/mol
Pink
Blue
Solver and chloride ions combine to form a precipitate of AgCl.
Safety Precautions:
HCl is hazardous and should be handled with care. It also has strong
vapors that should not be inhaled. Wash spills off with lots of water, and if
spilled neutralize with baking soda. NaOH is also hazardous. If spilled,
neutralize with vinegar. Ethanol is flammable. Turn off all flames. Silver
nitrate causes stains on skin and clothing. Wash spills off with soap and water
immediately. Wear lab safety goggles & a chemical-resistant apron.
Materials:
 HCI, 12 M AND 0.1 M
 NaCl(s)
 KSCN, 0.002 M
 KSCN(s)
 Bromothymol blue indicator
solution
 NaOH, 0.1 M
 Fe(NO3)3, 0.2 M
 AgNO3, 0.1 M
 Ethanol
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Cobalt(II) chloride/
CoCl2*6H2O(s)
Disodium hydrogen
phosphate/ Na2HPO4 (S)
Test tubes, 13 x 150-mm
Test tube rack
Beaker, 100 mL
Graduated cylinders
Funnel, filter paper, and
holder for funnel
Procedure:
Part 1 – Equilibrium in a Saturated Solution
1) Pour some solid NaCl into a 13- X 100-mm test tube.
2) Fill the tube ¾ with distilled water.
3) Cork and shake to saturate solution. If all the NaCl dissolves, pour
additional NaCl into the test tube and shake again until solution is
saturated and excess solid can be seen at the bottom of the test tube.
4) Filter the solution into a second test tube.
5) To this saturated NaCl solution, add Cl- ions in the form of
concentrated HCl. Record observations.
Part 2 – Acid-Base Indicator Equilibrium
1) Fill a small test tube about half-full of distilled water. Add several
drops of bromthymol blue indicator solution.
2) Add 5 drops of HCl and stir. Notes the color of the indicator.
3) Add 0.1M NaOH drop by drop while stirring until no further color
change occurs. Again, note the color.
4) See if you can add the right amount of acid to this test tube to cause the
solution to be green in color after it is stirred (half of the indicator is
blue and half is yellow).
Part 3 – A Complex Ion Equilibrium
1) Pour about 25mL of 0.0020M KSCN solution into a beaker.
2) Add 25mL of distilled water and 5 drops of Fe(NO3)3. Swirl the
solution, and note the following: the color of Fe(NO3)3 solution, the
color of the KSCN solution, and the color of the resulting complex ion.
3) Pour equal amounts of the solution from the beaker into four test tubes.
4) Add 2-3 crystals of solid KSCN to the 2nd test tube. Observe.
5) Add 6 drops of Fe(NO3)3 solution to the 3rd test tube. Stir and observe.
6) Add small crystals of Na2HPO4, a few at a time, to the 4th test tube. Stir
and observe.
Part 4 – Equilibrium with Cobalt Complex Ions
1) Measure 10mL of ethanol into a beaker.
2) Examine solid cobalt(II) chloride and note its color.
3) Dissolve a small amount of cobalt(II) chloride (about half the size of a
pea) in the beaker of ethanol. The solution should be purple. If it is
pink, add a little concentrated HCl until the color is purple.
4) Put 2mL of alcoholic cobalt solution into each of the 3 test tubes.
5) To one of the test tubes, add 3 drops of distilled water, one drop at a
time while stirring, noting what happens with each drop.
6) Add 3 drops of distilled water to each of the other 2 test tubes. Observe
the effect of this stress on the system.
7) To the 2nd test tube, add 5 drops of concentrated HCl, 12M, one drop at
a time while stirring. Observe.
8) To the 3rd test tube add a few crystals of solid NaCl. Stir and observe.
9) Put the remainder of the alcoholic cobalt solution from the beaker into
a 4th test tube. Add 10 drops of 0.1M AgNO3, one drop at a time. Note
the color of the solution as chloride ions precipitate.
10) Obtain a sealed Beral pipet containing some of the alcoholic cobalt
chloride-water system. Note its color, and immerse the large end of the
pipet in hot water (about 60◦C). Observe.
11) Chill the Beral pipet in an ice bath. Observe.
Analysis/Observations:
Part 1 - The saturated solution of NaCl had a clear color, and some tiny salt
particles were suspended throughout the test tube after shaking. Adding Cl- in
the form of HCl turned the solution yellow by forcing the reaction to the left
and the dissociated Cl- ions to recombine into HCL. Fumes were visibly seen.
Part 2 – Adding HCl to the bromothymol blue solution shifted the reaction to
the left, thus turning the solution a yellow-orange. Adding NaOH shifted the
reaction to the right, thus turning the solution blue. After adding a few drops of
HCl, the solution turned solution green, coinciding with the reach in
equilibrium.
Part 3 – KCN solution was observed to have a clear color, as was Fe(NO3)3.
The color of the complex ion was observed as a burnt orange. The color of the
solution in test tube 1, the reference, was observed to have a burnt orange
color. Adding SCN- to test tube 2 in the form of KSCN shifted the reaction to
the right, which turned the solution a dark red. Adding Fe(NO3)3 to test tube 3
also shifted the reaction to the right and turned the solution red, although not a
dark a shade of red. The reason behind the different intensity of hue is that
more SCN- dissociated than Fe+3. Test tube 4, which had NaPO4 added to it,
turned clear.
Part 4 – Solid Cobalt(II) Chloride solution was observed to have a magenta
color. The top ¾ of test tube 1 was purple in color, while the bottom ¼ of it
was pink. This was the result of adding water, which shifted the reaction to the
left and turned the solution pink. Adding water and HCl to test tube 2 turned
the solution blue, as adding Cl- in the form of HCl forced the reaction to the
right. Test tube 3, which had water and NaCl added to it, turned a lighter blue.
The color change was to a lighter shade of blue because NaCl does not
dissociate as many Cl- ions as does the strong acid HCl. Test tube 4, which had
silver nitrate added to it, turned dark blue on top and had a pink colored
precipitate on the bottom. Alcoholic cobalt(II) chloride in the sealed Beral
pipet that was plunged into hot water changed color from a light blue to dark
blue. This is because increasing the temperature shifted the reaction to right
and the solution to turn blue. Alcoholic cobalt(II) chloride in the sealed Beral
pipet that was dunked into an ice bath changed color from a light blue to pink.
This is because a decrease in heat shifted the reaction to the left.
Conclusion:
Overall, the application of Le Châtelier's Principle explained the
chemistry behind the color changes taking place in the experiment. When
adding reactant into a solution, the volume of the solution (and as a result the
concentration of the reactants in the solution) was immediately changed. Le
Châtelier's Principle was then predicted which side the reaction would be
shifted. Because the reactants and products were both of distinctly clear and
yellow color (Part 1), yellow-orange, blue, and green color (Part 2), clear,
burnt orange color, and light and dark shades or red (Part 3), purple and pink
and shades of light and dark blue (Part 4), respectively, change in equilibrium
could be easily observed through visible, qualitative color changes. These
color changes were indications of Le Chatelier's Principle, which states that if
a system is subjected to a stress, the system will react to remove the stress. To
remove the stress, reactions either: shifted to the right and formed more
products, using up reactants; or shifted to the left and form more reactants,
using up products.