Common-Ion Effect and Buffers Shela Marie L. Algodon
National Institute of Molecular Biology and Biotechnology, University of the Philippines, Diliman, Quezon City 1101 Philippines Experimental Details
Four different volumes of the following solutions were prepared in separate 50-mL beakers: Solution 1: 30 mL 0.10 M HOAc Solution 2: 15 mL 0.20 M HOAc ± 15 mL 0.20 M NaOAc Solution 3: 30 mL 0.10 M NH3 Solution 4: 15 mL 0.20 M NH3 ± 15 mL 0.20 M NH4Cl The pH of each solution was estimated using indicators. One drop of methyl orange served as indicator for Solutions 1 and 2. The colors of the two solutions were compared afterwards. On the other hand, one drop of phenolphthalein indicator was used as an indicator indicator for Solutions Solutions 3 and 4. The colors colors of the solutions were also compared. Using a properly calibrated pH meter, a more accurate pH reading of each solution was obtained. After determining the pH of each solution using a pH meter, each solution was divided into three equal portions. One drop of 1.0 M HCl was added to the first portion. To the second portion, one drop of 1.0 M NaOH solution was added. The third portion served as the control and the colors of the other two portions were compared with it. Again, the pH of each solution was estimated based on their colors, and a more accurate pH was obtained using the pH meter. The effect of common ions on the solutions and the effect of adding other reactants
in a buffer solution in this experiment are thus observed. esults R esults
and Discussion R eport eport
A buffer solution is a solution consisting of a conjugate acid±base pair where both the concentrations of the acid and the base are reasonable. Buffer systems resist changes in the pH upon addition of small amounts of strong acid or strong base. The acid in the buffer solution reacts with the strong bases added in the solution. Similarly, the base in the buffer solution reacts with the strong acids added added in the t he 2 solution . One factor affecting the processes involved in the buffer solution is the common ion effect, a special case of LeChatelier¶s Principle. Many solutions exhibit this kind of behaviour. However, the most frequently encountered types of solutions are (1) a solution of a weak acid plus a soluble ionic salt of the weak acid and (2) a solution of a weak base plus 1, 3 a soluble ionic salt of the weak base . TABLE 1: pH Using Visual Indicators and pH Meter
Color of Solution Solution
+ Methyl Orange
+ Phenolphthalein
pH Reading
1
Salmon pink
3.34
2
Yellow
4.82
3
Red
10.02
4
Very light pink
8.85
1
TABLE 2: Effect of Strong Acid and Strong Base on Buffers
Estimated pH Range Soluti on
+ Meth yl Oran ge
pH Met er
Calcula ted
No chan ge
3.28
2.27
Light er salm on pink
3.52
3.47
c
Norm al
3.34
2.86
a
No chan ge
4.43
4.70
b No chan ge
4.79
4.79
Norm al
4.82
4.74
9.33
10.53
a 1
b
2
c a
Very light pink
3
+ Phenolphtha lein
pH
b
Light pink
9.77
11.73
c
Normal
10.0 2
11.12
a
Lighter than normal
8.25
9.23
4
b c
No change Normal
8.85
8.63
9.30
9.26
In Solution 1, the solution with 30 ml of 0.10 M CH3COOH, the pH was obtained just by getting the concentration of the H + ion in the solution and the formula for pH: pH=-log[H+]. When the solution was divided into three equal portions, where 1 drop of 1.0 M HCl was added
to the first portion, 1 drop 1.0 M NaOH was added to the second portion, and the third portion served as the control solution, there were changes in the pH of the first two portions. It is because there is initial concentration for the H+ ion in portion 1, Solution 1a. This initial concentration attributes to the extent the forward reaction will proceed, as had been stated in the LeChatelier¶s Principle. The reason is the same for Solution 1b, the portion with 1 drop of 1.0 M + NaOH. Because the NaOH dissociates into Na and OH ions completely, the OH ions react with the CH3COOH to form CH3COO and H2O, giving the CH3COO an initial concentration before the dissociation of the CH 3COOH alone. The pH of Solution 2, the solution with 15 mL 0.20 M CH3COOH and 15 mL 0.20 M NaCH3COO, has a different pH from Solution 1 even if the solution also has the same concentration of CH 3COOH because CH3COO-, the common ion, has an initial concentration. This solution is a buffer solution. The NaCH3COO, a soluble ionic salt of CH 3COOH, + completely dissociates into Na and CH3COO . On the other hand, CH 3COOH only partially ionizes. When it was divided into three portions, the same reaction as that of Solution 1a and Solution 1b happens with the Solution 2a and 2b. Since there is a common ion CH 3COO-, the change in the pH of the solution will be minimal compared to Solutions 1a and 1b since CH3COO- will react with the added strong acid, and CH3COOH will react with the added strong base. The same principle applies for both Solutions 3 and 4 and their portions. The pH of the solution with 30 mL 0.10 M NH 3, Solution 3, was obtained by calculating the amount of OH ions and using the equation for pOH: pOH=-log[OH ]. Since pH is what is asked, the value for pOH is subtracted from 14 based on the formula pH + pOH=14. The solution is divided into three portions. When 1 drop of 1.0 2
M HCl was added to the first portion, the H + from the completely dissociating HCl reacts with NH3 to form NH4+. Thus, there is an initial + concentration for NH4 . On the other hand, when NaOH was added to the second portion, the + NaOH dissociates completely into Na and OH , giving an initial concentration of OH ions even before NH3 dissociates. When the HCl and NaOH was added to solution 4, the solution with 15 mL 0.20 M NH 3 and 15 mL 0.20 M NH 4Cl, there is not much change in the pH since it is a buffer solution. Strong acids and strong bases react with the existing bases and acids, respectively, as is the case for this solution. The H + ions from the added HCl reacts with the NH 3 producing more NH4+ in return. Similarly, the OH - ions from the added NaOH reacts with NH 4+ to produce the compound NH3. As can be seen from Table 2, the values obtained from using a pH meter to determine the pH of the solution is different from the values theoretical values obtained. This is in account for the human errors and the accuracy of the
measuring equipment which doesn¶t always give exact amounts. However, it should be noted that the trend in changes of the pH for both the experiment and theoretical values of the pH are the same. It is recommended to use more accurate measuring devices so as to get more reliable data, and avoid errors. Moreover, it should always be remembered to wait for the stable sign in the pH meter before recording data.
R eferences
(1) Ebbing, D. D., Gammon, S. D.(2009) General Chemistry, 9 th ed . USA: Houghton Mifflin Company (2) Pauling, L. (1970) General Chemistry. US: Dover Publications (3) Whitten, K. W., et al. (2004) th General Chemistry, 7 ed. US: Brooks/Cole
3
Appendices
Appendix A pH of Solution 1: 30 mL 0.10 M HOAc
CH3COOH Initial
0.10
Change
-
Equilibrium
x
0.10 - x
CH 3COO
H
0
0
+x
+x
x
x
«-CH3 COO »½ «- H »½ K a !
?CH COOH A 3
1.8 v 105 !
x
xx 0.10 x
! 1.332670973v 103
pH ! log «- H »½ pH ! log «-1.332670973v 103 »½ pH ! 2.875277062
4
Appendix B pH of Solution 2: 15 mL 0.20 M HOAc ± 15 mL 0.20 M NaOAc
¨ ?baseA¸ © ?acid A¹¹ ª º
pH ! pKa log ©
¨ ?0.20A¸ © ?0.20A¹¹ ª º
pH ! log(1.8 v 105 ) log © pH ! 4.744727495
5
Appendix C pH of Solution 3: 30 mL 0.10 M NH3
NH 3
Initial
0.10
Change
-
Equilibrium
x
0.10 ± x
H 2 O
p n
NH 4
OH
0
0
+x
+x
x
X
«- NH 4 »½ «- OH »½ K b !
? NH A 3
1.8 v 105 !
x
xx 0.10 x
! 1.332670973v 103
pOH ! log «-OH »½ pOH ! log «-1.332670973v 103 »½ pOH ! 2.875277062
pH pOH ! 14 pH 2.875277062 ! 14 pH ! 11.12472294
6
Appendix D pH of Solution 4: 15 mL 0.20 M NH3 ± 15 mL 0.20 M NH4Cl
¨ ?acid A¸ © ?baseA¹¹ ª º
pOH ! pKb log ©
¨ ?0.20A¸ © ?0.20A¹¹ ª º
pOH ! log(1.8 v 105 ) log © pOH ! 4.744727495
pH pOH ! 14 pH 4.744727495 ! 14 pH ! 9.255272505
7
