Conversely, if some iodine is removed from the container, there will be fewer collisions per second between hydrogen molecules and iodine atoms, resulting in a slower forward reaction. The reverse reaction will predominate until a new equilibrium is established with less hydrogen iodide, more hydrogen, and some what less iodine than in the original mixture. Any such concentration changes will affect the "balance" of opposing reactions in accordance with Le Châtelier's Principle. When some stress is applied to a system originally in equilibrium, the system (reaction) automatically will shift in such a direction as to relieve the stress (in this example a change in concentrations) and restore the original conditions as much as possible.
The common ion effect is a special case of the application of the law of chemical equilibrium to ionization reactions. For example, in a solution of the weak base, ammonium hydroxide, there is the equilibrium
NH4OH <-----> NH4+1 + OH-1
The addition of NH4Cl, (NH4)2SO4, or any other soluble ammonium salt, will increase the concentration of NH4+1, and therefore increase the number of collisions per second between NH4+1 and OH-1. The equilibrium will be shifted to the left, and the concentrations of the OH-1 will be decreased. The NH4+1, since it is common to both the ammonium hydroxide and the added ammonium salt, is called a "common ion".
In the same way, salts that are only slightly soluble can be made even less soluble by increasing the concentration of a common ion. For example, the equilibrium between the slightly soluble salt gypsum, CaSO4.2H2O, and its ions in solution is represented by the equation
CaSO4.2H2O(s) <-----> Ca2+ + SO42- + 2 H2O
The addition of either Ca2+ or SO42-,
from any soluble salt containing one of these ions, would shift the
equilibrium to the left and decrease the solubility.
1. The Shifting of an Equilibrium. The Common Ion Effect.
a) The Chromate Ion-Dichromate Ion
2 CrO42- + 2 H+1 <-----> 2 HCrO4-1 <-----> Cr2O72- + H2O
At present, we need consider only the over-all reaction,
2 CrO42- + 2 H+1 <-----> Cr2O72- + H2O
To 3 mL of 1 M K2CrO4 add several drops of 2 M H2SO4. Mix this, and observe any change. Now add several drops of 6 M NaOH with mixing, until a change occurs. Again add H2SO4. Interpret the observed changes in terms of "shifting the equilibrium."
b) Weak Acids and Weak Bases
To 3 mL of 0.1 M HC2H3O2 add a drop of methyl orange, then add 1 M NaC2H3O2, a few drops at a time, with mixing. Explain your observations in terms of the equilibrium equation and the "common ion" effect.
To each of two 3 mL samples of 0.1 M NH4OH add a drop of phenolphthalein. To one sample add 1 M NH4Cl, a few drops at a time, with mixing. To the other add 6 M HCl, a drop at a time, with mixing. In each case, note any changes in colour and in odour of the solutions. Write the equation for the equilibrium in dilute NH4OH and interpret the different observed results, including the additional overall net ionic equation for the reaction with HCl, in terms of shifting the equilibrium.
c) Saturated Solution Equilibria
(2) BaCrO4(s). To 3 mL of 0.1 M BaCl2,
a few drops of 1 M K2CrO4, and then a little 6 M
Explain your observations in terms of the equilibrium equations
d) The Thiocyano-iron(III) Complex Ion
Fe3+ + SCN- <----> Fe(SCN)2+
Add 1 mL of 0.2 M Fe(NO3)3 to 2 mL of
0.1 M KSCN, and dilute this until a moderately red colour is attained;
25 to 35 mL of distilled water will be required. To a 5 mL portion of
add a little 0.2 M Fe(NO3)3; to a second 5 mL
add a little 0.1 M KSCN; and to a third 5 mL portion add a few drops of
M NaOH. (Fe(OH)3 is quite insoluble.) Correlate your
with Le Châtelier's principle and the above equilibria. Do your
prove the correctness of the above equation?
2. A Quantitative Study of the Equilibrium
Between Iron(III) Ion and Thiocyanate Ion
You will first prepare a standard comparison solution of red thiocyano-iron(III) complex of known concentration by mixing a known low concentration of SCN- ion with a large excess of Fe3+ ion. You can assume essentially complete reaction, so that the concentration of the red complex can be calculated as equal to the total SCN- concentration in the mixture.
You will then prepare a number of test tube mixtures, with successively lower Fe3+ concentration, resulting in a corresponding shift of the equilibrium and a corresponding decreased concentration, and colour, of the red complex. We can measure this decreased concentration by removing solution from the first known standard sample until the colours, viewed lengthwise through the tubes, exactly match. The concentrations will then be inversely proportional to the respective depths of solution.
Colourmetric Concentration Measurements
Remove small amounts of standard solution from test tube No. 1 (putting this into a small clean beaker or Erlenmeyer flask, since you may need it again), until the colour intensities appear exactly the same when viewed lengthwise through the tubes. Reverse the viewing positions of the tubes for better comparison and make final adjustments of the solution in test tube No. 1 with a dropper pipette. Measure the depth of each solution to the nearest millimetre. Repeat this comparison procedure with test tubes No. 1 and 3, No. 1 and 4, and No.1 and 5.
Calculate your quantitative observations so as to
evaluate the possible equilibrium involved, in the following way. First
equilibrium constant expression, assuming that the red complex is
Fe(SCN)2+, from the equation given. Calculate the initial
concentrations (before reaction) of Fe3+ and of SCN-
in each of the test tubes. (Express concentrations to two significant
figures, in exponential notation, such
as 5.0 x 10-4) Assume that in tube No. 1 all the SCN-
reacts to form Fe(SCN)2+; then, from the relative depth
in each tube No.2-5, compared with tube No. 1, calculate the
concentration of Fe(SCN)2+ in each tube. You can then
the equilibrium concentrations of Fe3+ and of SCN-,
by subtracting this Fe(SCN)2+ concentration from the
Fe3+ and SCN- initial concentrations for each
No. 2-5. Finally calculate the value of Keq for each mixture
Lab Report Reversible Reactions and Chemical Equilibrium
1. The Shifting of an Equilibrium, The Common Ion Effect
a) The Chromate Ion-Dichromate Ion Equilibrium.
Rewrite the equation for this equilibrium
Describe the colour changes obtained on the addition the H2SO4
and NaOH, respectively, and interpret these effects.
b) Weak Acids and Weak Bases
The equation for the equilibrium in dilute acetic acid is:
The equation for the equilibrium in dilute ammonium
hydroxide solution is:
Note any observed odour or colour changes when 0.1 M NH4OH (+ phenolphthalein) is treated with:
NH4Cl ________________________ HCl _______________________
Which direction, right or left, does each reagent shift the
above equilibrium? Explain fully.
c) Saturated Solution Equilibria
(1) NaCl. Calculate the molarity of saturated NaCl, bases on the data in the experimental directions.
______ M NaCl
Write an equation for the equilibrium present in saturated
NaCl, and explain your observations when concentrated HCl was added to
(2) BaCrO4 Write the equation for the
equilibrium established when a few drops of 1 M K2CrO4
is added to 3 mL of 0.1 M BaCl2, and explain any changes
this is treated with a little 6 M HCl.
(d) The Thiocyano-iron(III) Complex Ion
The equilibrium equation is Fe3+ + SCN-1 <------> Fe(SCN)2+
Explain fully, in terms of Le Châtelier's principle,
the effect on those equilibria of adding (1) 0.2 M Fe(NO3)3,
(2) 0.1 M KSCN, (3) 6 M NaOH.
2. A Quantitative Study of the Equilibrium Between Iron(III) Ion and Thiocyanate Ion
For the equilibrium Fe3+ + SCN-1 <------> Fe(SCN)2+
Equilibrium [ ]'s