Titration Curves of Strong and Weak Acids and Bases
Experiments was carried out to titrate strong and weak acids and bases, to understand the differences in titration curves when strong and weak acids are titrated against strong and weak bases. Besides, selection of appropriate indicator for titration was also another purpose of these experiments.
In aqueous solutions, always exist equilibrium between H+ ions, OH- ions and H2O as given by the equation: H2O(l) H+(aq) + OH-(aq); Kw = [H+][OH-]; Kw = 1×10-14 at 25oC.
If [H+] = [OH-] the solution is neutral
If [H+] > [OH-] the solution is acidic
If [H+] < [OH-] the solution is basic
There is another scale called pH scale to express [H+] of in an aqueous solution. This scale is defined as pH = -log [H+]. In terms of [H+], [OH-] and pH; the acidic, basic and neutral solutions at 25oC are defined as
Solution [H+] /M [OH-] / M pH
Acidic > 1×10-7 < 1×10-7 < 7.0
Neutral = 1×10-7 = 1×10-7 = 7.0
Basic < 1×10-7 > 1×10-7 > 7.0
The pH value of an acidic solution is < 7.0 because acids release H+ into aqueous solution and pH value of an alkaline solution is more than 7.0 because bases release [OH-] into aqueous solution. Strong acids like HCl dissociate completely into aqueous solution:
HCl(aq) H+(aq) + Cl-(aq)
Weak acids only partially dissociate in aqueous solution like acetic acid
HC2H3O2(aq) H+(aq) + C2H3O2-(aq)
Strong bases like NaOH completely dissociate in aqueous solution as per following reaction
NaOH(aq) Na+(aq) + OH-(aq)
Weak bases like NH4OH only partially dissociates in aqueous solution as per following equation
NH4OH(aq) NH4+(aq) + OH-(aq)
When an acid is added to an alkaline solution, a salt and water forms. This reaction known as neutralization reaction forms the basis of titration, which is nothing but measuring unknown concentration of an alkaline solution with known concentration of an acidic solution and vice versa. Some examples of acid base reaction are listed below:
H+(aq) + Cl-(aq) + Na+(aq) + OH-(aq) Na+(aq) + Cl-(aq) + H2O(l)
HC2H3O2(aq) + Cl-(aq) + Na+(aq) + OH-(aq) Na+(aq) + C2H3O2-(aq) + H2O(l)
H+(aq) + Cl-(aq) + NH4OH(aq) NH4+(aq) + Cl-(aq) + H2O(l)
HC2H3O2(aq) + NH4OH(aq) NH4+(aq) + C2H3O2-(aq) + H2O(l)
As neutralization reaction proceeds, more and more of salt and water forms and pH of the acidic solution starts to move towards 7.0.
Before the equivalence point, pH of the solution is determined by the concentration of acid. For same concentration of acid, a strong acid like HCl gives lower pH than a weak acid like acetic acid. At the equivalence point, which is the point at which sufficient base has been added to the acidic solution to completely neutralize it, pH of the solution is due to the resulting salt. At equivalence point, the salt of strong acid and strong base gives a neutral solution; the salt of strong acid and a weak base produces slightly acidic solution (pH <7.0); the salt of weak acid and a strong base produces slightly basic solution (pH > 7.0) and salt of weak acid and weak base produces a solution pH of which depends on relative strength of the acid and the base.
Beyond the equivalence point, pH of the solution depends on the concentration of the base being added. For same concentration of the base, the strong base like NaOH gives higher pH than a weak base like ammonium hydroxide.
Titration curves are useful in determining concentration of acidic or basic solutions. Control of pH is very important in biological systems, because biochemical reactions are pH sensitive. Human blood pH remains in 7.3 – 7.5 range and a change of more than 0.4 in pH is fatally dangerous.
In these experiments titration of strong and weak acids – hydrochloric and acetic acids – has been carried out using strong and weak bases – sodium hydroxide and ammonium hydroxide and the resulting titration curves have been obtained and analyzed.
IBM compatible computer, Serial box interface, Vernier pH amplifier and pH electrode, Magnetic stirrer, magnetic stirring bar, 250 mL beaker, phenolphthalein indicator, wash bottle, 0.10M NaOH, 0.10M NH4OH, 0.10M HCl, 0.10M HC2H3O2, 50 mL buret, ring stand, two utility clamps, distilled water.
1. Lab coat, gloves and safety glasses were used as safety measure during the experiment as the experiment involved handling of hazards chemicals.
2. Approximately 8 mL of 0.10M HCl solution was put into 250 mL beaker and 100 mL distilled water was added into this. 2-3 drops of phenolphthalein acid-base indicator was also added into the solution.
3. The beaker was kept over a magnetic stirrer, the magnetic stirring bar was put into it and the stirrer was turned on and adjusted for slow stirring of the solution.
4. Using a utility clamp a pH electrode was suspended on a ring and dipped into the hydrochloric acid solution.
5. A 50 mL buret was taken and rinsed with a few mL of 0.10M NaOH solution. The buret was attached to the ring stand using a utility clamp. It was filled with 0.10M NaOH upto a little above the 0.0 level and a small amount of NaOH was drained until the bottom of meniscus came at the 0.0 level.
6. Computer was prepared for acquiring the data during the experiment. The vertical axis has pH scale from 0 to 14 units. The horizontal axis has time scale from 0 to 250.
7. The titration curve was prepared by clicking “collect” and instantly opening the buret stopcock to provide a dipping rate of 1-2 drop per second.
8. The time to change in the color of phenolphthalein indicator was noted and flow of NaOH was stopped by turning the buret stopcock after 250 seconds when data collection ends.
9. The graph title was given and a printout of the graph was taken
10. The experiment was repeated for the other acid base pairs and thus a total of four titration curves were produced.
5. Results and Analysis:
All the four titration curves are presented below:
Figure 1 shows titration curve of HCl with NaOH, in the initial stage of titration the pH changes very slowly in 2 – 3.5 range. When equivalence point is approached, the pH changes very rapidly from 3.5 to 11 with just a few drops of the NaOH added and then stabilizes around 12.0.
Figure 3 shows titration curve of acetic acid with NaOH. In this case the pH starts from more than 4 and rises gradually to 6.0, then suddenly it shoots to 11.0 near the equivalence point and finally stabilizes near 12.0.
Figure 3 shows titration of HCl with ammonium hydroxide. In this case the change of pH is very gradual even near the equivalence point. This can be used as an acidic buffer solution. The transition region is very gradual.
Figure 4 shows titration curve of acetic acid with ammonium hydroxide. The change of pH is gradual and the transition region is sharp but very small near the equivalence point.
a) The experiments show that indeed it is very difficult to control pH near 7.0, as pH changes very rapidly.
b) Comparing figures 3 and 4, one can conclude that it is better to use a strong acid to neutralize a basic solution to keep the pH near 7.0. It is because the change in pH is gradual with strong acid (figure 3) and steep with weak acid (figure 4).