EXPERIMENT 5 COMMON ION EFFECT DABUET, Nathalie E. GRINO, Manuel Angelo T.
AB2, Group 5, Julius Andrew Nunez May 2, 2013
I. ABSTRACT The experiment was conducted to observe a special case of Le Chatelier’s Principle; the Common Ion Effect, which, was tested through changes in ionization of acid and base solutions, the buffering effect, and changes in the solubility, all in the presence of common ions. The changes in ionization were observed through the measurement of the pH of three types of solutions: a strong acid solution, a weak acid solution and a strong base solution in two circumstances, one wherein water was added and another which instead had a substance which shares a common ion with the acid and base. The buffering effect was observed through the preparation of solutions where a strong acid or base was supplemented to observe its effect on the pH of the solution, which was measured beforehand. Lastly, through comparison with the data gathered from Experiment 4, the effect of common ions in the solubility of slightly soluble salts was determined. II. KEY WORDS Common Ion Effect, Buffering Effect, Le Chatelier’s Principle, Solubility III.
INTRODUCTION The common-ion effect is used to describe the effect on an equilibrium involving a substance that adds an ion that is a part of the equilibrium. Le Châtelier's Principle states that if equilibrium gets out of balance, the reaction will shift to restore the balance. If a common ion is added to a weak acid or weak base equilibrium, then the equilibrium will shift towards the reactants, in this case the weak acid or base. Adding a common ion prevents the weak acid or weak base from ionizing as much as it would without the added common ion. The common ion effect suppresses the ionization of a weak acid by adding more of an ion that is a product of this equilibrium. This results to the change in pH value of the solution because of the lessened value of hydrogen or hydroxide ions. In relation to strong acidic and basic solutions, this effect is not highly observed since it is difficult for an almost completely ionized system to divert back to the formation of its reactants.
This is the driving force behind buffer solutions. A buffer is an aqueous solution that has a highly stable pH. If you add acid or base to a buffered solution, its pH will not change significantly. Similarly, adding water to a buffer or allowing water to evaporate will not change the pH of a buffer. A buffer is made by mixing a large volume of a weak acid or weak base together with its conjugate. A weak acid and its conjugate base can remain in solution without neutralizing each other. The same is true for a weak base and its conjugate acid. There are two types of buffer solutions; acidic, made up of a weak acid and conjugate base (pH less than 7,) and alkaline, made up of a weak base and conjugate acid (pH greater than 7,). Buffer solutions act upon a supplementary base or acid solution to its system by eliminating any newly formed hydrogen ions or
hydroxide ions consequently neutralizing the added solution. In a buffer system, three types of chemical reaction exist: Auto ionization of water, Acid dissociation, and Hydrolysis. This unfastens the system to readily interchange protons amongst the acid/base, conjugate base/acid, hydrogen and hydroxide ions, and water. On one occasion a small amount of acid is added, it primarily ionizes to hydrogen ions, which are expended by the current conjugate base in the system. Whereas an added base ionizes to hydroxide ions and is consumed by the present acid in the solution through the suppression of the formation and addition of hydrogen as well as hydroxide ions, thus the pH value of the system is retained. IV.
METHODOLOGY The following apparatus were used in the experiment: 10mL test tubes, 100mL beakers, 50mL Erlenmeyer flask, a titration set-up, and a heating set-up. The reagents used were: Hydrochliric acid (HCl), Sodium Chloride (NaCl), Acetic acid (HOAc), Sodium Hydroxide (NaOH), Nitric acid (HNO3), Sodium Phosphate Monobaric Monohydrate (NaH2PO4), Ammonium Hydroxide (NH4OH), Sodium Nitrate (NaNO3), Ammonium Chloride (NH4Cl), Sodium Benzoate (NaC6H5COOH), and Benzoic acid (C 6H5COOH). The first part of the experiment, involved the preparation of acidic and basic solutions for pH testing with the use of pH paper. This was to observe the effect of common ions on the ionization of acids and bases. There were two types of solution; one in which the acid or base was diluted with water and another, wherein a solution
containing an ion common to that of the acid or base was added. The pH was then recorded for interpretation. The second part was for the observation of the buffering effect. The pH of distilled water was first determined. Then a drop of 6M HCl was added to 10 mL of the distilled water. The same procedure was done once more, whilst replacing HCl with NaOH. The pHs of the resulting solutions were recorded. Five different solutions were then prepared. These solutions were divided into two equal parts, one was added with a drop of HCl while the other, NaOH. The pHs of the ten solutions were determined. The following solutions were prepared:
Solution A. 10 mL 0.5 M HOAc + 10 mL 0.5 M NaOAc B. 10 mL 0.5 M HCl+ 10 mL 0.5 M NaCl C. 10 mL 0.5 M HNO3+ 10 mL 0.5 M NaNO 3 D. 10 mL 0.5 M NaH2PO4+ 10 mL 0.5 M NaH 2PO4 E. 10 mL 0.5 M NH 4OH+ 10 mL 0.5 M NH4Cl
The last part of the experiment was for the determination of the effect of common ions on the solubility of slightly soluble salts. 50 mL of distilled water was added with 0.5 g of Sodium Benzoate o and was heated to about 40 C. Benzoic acid crystals were then added until the solution was saturated. It was then cooled to room temperature by stirring and was filtered. Two drops of Phenolphthalein was added to 10 mL of the filtrate. It was then titrated with 0.01 M NaOH to a light pink end point. The collected data was noted for comparison to the data gathered from the previous experiment, Experiment 4: Ionic Equilibria. V.
10 mL H2O + 1 drop 6 M NaOH
14
Table 2. Part B- (i) pH results for H2O and H2O solutions
pH of the original solution
pH after addition HCl
pH after addition NaOH
Exp
Theo
A
5
5
5
B
B
B
1
1
1
B
NB
C
1
1
1
B
NB
D
7
6
7
B
B
E
8
1
9
NB
B
Solution
Conclusion
Table 3. Part B- (ii) Experimental and Theoretical results for Buffering Effect.
Volume of 0.01 M NaOH
16. 2 mL
Sollubility of Benzoic Acid in Water
0.013 M
Solubility of Benzoic Acid in Sodium Benzoate Solution
0.0162 M
Table 4. Part C- Results for Effect of Common Ion on Solubility of slightly soluble salts
COMPUTATIONS 2
+
-
S = Ksp = [H3O ][A ] + 1/2
- 1/2
S = [H3O ] [A ] +
-
[H3O ] = [A ] = ([0.01]16.2mL)/(10mL)
RESULTS
= [0.0162]
Solution 10 mL 0.1 M HCl + 2 mL H 2O 10 mL 0.1 M HCl + 2 mL 0.1 M NaCl 10 mL 0.1 M HOAc + 2 mL H 2O 10 mL 0.1 M HOAc + 2 mL 0.1 M NaOAc 10 mL 0.1 M NaOH + 2 mL H 2O 10 mL 0.1 M NaOH + 2 mL 0.1 NaCl
pH 1 1 3 4 14 10
Table 1. Part A- Experimental pH of the effect on the ionization of acids/bases
Solution
pH
10 mL H2O
7
10 mL H2O + 1 drop 6 M HCl
2
2x½
S = [0.0162] VI.
= [0.0162]
DISCUSSION The results in Part A, which involved the ionization of acids and bases showed minimal changes in the pH when water was added in all three electrolyte solutions. This was because of water’s neutral nature thus causing infinitesimal deviations on the ionization of the solutions. The solutions where a substance, which share a common ion was added, however, cause some deviations for two of the three set-ups’ pH values. There was no observable change on the pH of HCl solutions as HCl is a strong electrolyte and therefore, dissociates completely. The addition of + Cl does not affect the concentration of H3O ions,
as it is a weak base. The same should have also been the case for the NaOH solutions theoretically since it has a weak conjugate acid. There may have been some errors on the preparation or evaluation of the pH of the sixth solution thus causing a deviation from the expected result. However, it was expected that the pH of the solutions containing HOAc would exhibit a difference. This is because HOAc is a weak acid and does not dissociate entirely. Its conjugate base is strong and had therefore affected the pH of the solution. In Part B, wherein the buffering effect was observed, there were certain inconsistencies between the experimental and theoretical results. A buffer by definition is a solution that contains both an acidic and basic species which do not neutralize each other but instead neutralize excess + OH an H and are therefore able to co-exist. These species resist changes in the pH of the solution when small amounts of an acid or base are added, as they neutralize the additional acid or base. The addition of a common ion causes a shift in equilibrium and sets up the required condition for buffering behavior. The inconsistencies may have risen from erroneous preparations of the solutions or misinterpretation of the pH. Too much of the acid or base may have also been added which caused a relatively sizeable effect. Lastly, in Part C, which involved the comparison of the data collected from the previous experiment, an unexpected result had occurred. Theoretically, the addition of a common ion is supposed to have decreased the solubility of the slightly soluble salt used. The addition of a common ion causes a backward shift in the reaction and therefore lowers the solubility factor of the salt as less amounts of it dissociates. In the experiment, however, when the solution was titrated, a larger volume of the titrant was consumed which foretells greater dissociation, and in turn, greater solubility. An error may have occurred in Experiment 4. The solution may have not been completely saturated when it was prepared and hence causing a lower molarity for the Benzoic acid. CONCLUSIONS AND RECOMMENDATIONS The Common Ion Effect affects the ionization of weak electrolytes as the addition of a + strong conjugate decreases the amount of H or OH ions resulting to a backward shift of the reaction. Strong electrolytes are unaffected as they have weak conjugates, which do not have much of a reaction with the ions in the solution.
The addition of common ions to an electrolyte solution hinders the formation of new ions thus resisting alterations in the pH of the + solutions. This is due to the consumption of H or OH ions in the formation of the original reactants. The addition of a common ion decreases the solubility of a salt as its presence causes a backward shift, which hinders the salt’s dissociation. For future experimentations, greater accuracy in the measurement of the pH values is recommended. Moreover the students must also make sure that the salt solution is saturated for consistent results. The amounts of acids and bases supplemented to buffer solutions must also be held constant and carefully regulated to ensure uniformity. Additionally, the reagents must be handled with utmost care in order to avoid contamination.
VIII.
REFERENCES
Handout on buffer solutions. (2001). Retrieved May 1, 2013 from http://chem.cmu.edu/courses/09106/notes/buffers.pdf Helmenstine, A. M. (n.d.). Buffers: Chemistry of Buffer. Retrieved May 1, 2013, from About.com: http://chemistry.about.com/od/acidsbase1/a/buffer s.htm Tung, E and Danai, M. (n.d.). Common Ion Effect. Retrieved May 1, 2013, from UC Davis ChemWiki: http://chemwiki.ucdavis.edu/Physical_Chemistry/A cids_and_Bases/Solubilty/Common_Ion_Effect
I hereby certify that I have given substantial contribution to this report.
VII.
NATHALIE ELONA DABUET
MANUEL ANGELO TERANIA GRINO
