Colorimetric Determination of pH

Experiment #6: Colorimetric Determination of pH
Almira, Faerie Carleen Lucile L.
Gallardo, Charlotte O.
Group #6, Chemistry 18.1, MHEG1, Ma’am Arlou Angeles
September 23, 2013

I. Abstract
The acidity of the four unknown solutions were determined with the use of colorimetry using McIlvaine’s buffer solutions varying in proportion of its constituents (disodium phosphate and citric acid). These buffer solutions were subjected to the addition of corresponding pH indicators and the variation of colors depending on its pH level was used as standards. At the end of the experiment, the colors of the unknowns were compared with the standard buffer solutions and then calculated to obtain an estimation of their pH or acidity. Experimental results show that Solution A, B, C, and D, have pH levels of 4.9, 5.2, 4.2,and 4.9 respectively.
Keywords: McIlvaine buffers, colorimetry, pH indicators, Henderson-Hasselbalch equation

II. Introduction
Colorimetry is a quantitative chemical analysis with the use of standard colors and the determination of the spectral absorbance of a solution. This is essential in determining the acidity or basicity of a substance using acid-base indicators (also known as pH indicators). These indicators are usually weak acids or weak bases that, when added in solutions, helps detect changes in pH presented visually through color change. Common pH indicators are shown in Table 1.

Table 1. Common pH indicators used in Colorimetric determination of pH.
Indicator
Lower pH Color pH Range
Higher pH Color
Thymol blue
Red
1.2 - 2.8
Yellow
Bromophenol blue
Yellow
3.0 - 4.6
Purple
Chlorophenol red
Yellow
4.8 - 6.4
Violet
Bromothymol blue
Yellow
6.0 - 7.6
Blue
Phenol red
Yellow
6.8 - 8.4
Red

In this experiment, the colorimetric comparison was done using McIlvaine buffers which contain a citrate and a phosphate, using citric acid and Na2HPO4 stock solutions, to measure pH of an unknown in a wide range. Buffer solutions are solutions that contain an acid-base pair in reasonable concentrations. McIlvaine’s Buffer solutions can resist changes in pH because they contain both an acidic species to react with the OH- ion, and a basic species to react with the H3O+ ion. This is an example of a Common Ion Effect, a special case of Le Chatelier’s Principle that occurs when a given ion is added to an equilibrium mixture, and the position of equilibrium shifts away from forming more of it.
HA + H2O H3O+ + A-
Addition of more H3O+, shifts the equilibrium to the left forming more HA.

B- + H2O OH- + HB
Addition of more OH-, shift the equilibrium to the left forming more B-.

At the end of the experiment, the pH of the unknown solution was obtained through colorimetric analysis and using the Henderson-Hasselbach equation stating pH = pKa + log [conj. base]/[acid]. The ionization constant of a weak acid was also calculated using the equation: [H3O+][A-]
Ka = ____________ [HA]

III. Experimental
PART A.
The McIlvaine’s buffer solution was prepared using 0.2 M Na2HPO4 and 0.1 M citric acid with certain amounts to produce certain pH levels appropriate for each indicator:

Table 2: 5 drops of methyl orange were used to obtain standard colors of the prepared buffer solution in different proportions.

Methyl orange pH mL Na2HPO4 mL Citric Acid
2.8
1.58
8.42
3.0
2.05
7.95
3.2
2.47
7.53
3.4
2.85
7.15
3.6
3.22
6.78
3.8
3.55
6.45
4.0
3.25
6.15
4.2
4.14
5.86
4.4
4.41
5.59
Table 3: 5 drops of bromophenol blue were used to obtain standard colors of the prepared buffer solution in different proportions.

Bromophenol blue pH mL Na2HPO4 mL Citric Acid
3.0
2.05
7.95
3.2
2.47
7.53
3.4
2.85
7.15
3.6
3.22
6.78
3.8
3.55
6.45
4.0
3.86
6.15
4.2
4.14
5.86
4.4
4.41
5.59
4.6
4.67
5.33

Table 4: 5 drops of bromocresol green were used to obtain standard colors of the prepared buffer solution in different proportions.

Bromocresol green pH mL Na2HPO4 mL Citric Acid
3.8
3.55
6.45
4.0
3.25
6.15
4.2
4.14
5.86
4.4
4.41
5.59
4.6
4.67
5.33
4.8
4.93
5.07
5.0
5.15
4.85
5.2
5.20
4.80
5.4
5.58
4.42

Table 5: 5 drops of chlorophenol were used to obtain standard colors of the prepared buffer solution in different proportions.

Chlorophenol red pH mL Na2HPO4 mL Citric Acid
4.8
4.93
5.07
5.0
5.15
4.85
5.2
5.20
4.80
5.4
5.58
4.42
5.6
5.80
4.20
5.8
6.05
3.95
6.0
6.31
3.69
6.2
6.61
3.39
6.4
6.92
3.08

Table 6: 5 drops of bromothymol blue were used to obtain standard colors of the prepared buffer solution in different proportions.

Bromothymol blue pH mL Na2HPO4 mL Citric Acid
6.0
6.31
3.69
6.2
6.61
3.39
6.4
6.92
3.08
6.6
7.34
2.66
6.8
7.72
2.28
7.0
8.24
1.76
7.2
8.69
1.31
7.4
9.08
0.92
7.6
9.37
0.63

Table 7: 5 drops of phenol red were used to obtain standard colors of the prepared buffer solution in different proportions.

Phenol red pH mL Na2HPO4 mL Citric Acid
6.4
6.92
3.08
6.6
7.34
2.66
6.8
7.72
2.28

Tables 2-7 represents the indicator used with the corresponding pH of the Mcilvaine buffers. The indicators only covers approximate pH ranges.

PART B.
Following amounts of distilled H2O, 0.1 M CH3COOH, and 0.1 M NaCH3COO were mixed in four different test tubes (duplicated for two sets of indicators) and labeled as follows:

Table 8. Composition of the Unkown pH Solutions.

d H2O
0.1 M HOAc
0.1 M NaOAc
Indicator 1
Indicator 2
A
0 mL
10 mL
0 mL
Bromophenol blue
Methyl orange
B
8 mL
1 mL
1 mL
Chlorophenol red
Bromocresol green
C
8.9 mL
1 mL
0.1 mL
Bromophenol Blue
Methyl orange
D
8.9 mL
0.1 mL
1 mL
Bromothymol blue
Phenol red

Upon addition of two drops of designated indicators for every test tube, the resulting colors of the unknown were compared to that of the standard buffer solutions from part A to obtain the acidity for each. Each of the two results from unknown solutions A, B, C, and D were averaged to obtain estimated acidity.

IV. Results

Table 9. Observed and Calculated pH of Unknown pH Solutions.
Sol’n
Observed pH
Ave. pH
Calculated
pH
A
6.4
(bromophenol blue)
3.4
(methyl orange)
4.9
4.88
B
5.2
(chlorophenol red)
5.2
(bromocresol green)
5.2

4.74
C
4.0
(bromophenol blue)
4.4
(methyl orange)
4.2
3.74
D
4.8
(bromothymol blue)
5.0
(phenol red)
4.9

5,74

Above is a table showing the obtained and calculated pH values of Solutions A to D.

V. Discussion
Buffer solutions can be composed with a weak acid and its conjugate base pair and vice versa—a weak base and its conjugate acid. The solution in equilibrium is able to resist drastic changes in pH. When a strong acid is added to the buffer, the acid will be neutralized with by the conjugate base of the buffer. The equilibrium then would push forward because of the Le Chatelier’s principle and common ion effect thus the final pH level will not be drastically affected. The common ion in the experiment is OAc-.
McIlvaine’s Buffer System is composed of two stock solutions (0.2 M disodium phosphate and 0. 1 M citric acid) combined in such volumes to cover a range of pH 2.8- 7.6 (originally 2.2 to 8.0).
Two different indicators were used in each of the unknown pH solutions. Acid-base indicators should be used appropriately according range for pH change for better accuracy in determining pH of the unknown. Solutions B to D differ between the volumes of HOAc and NaOAc added. Solution B has the same HOAc and NaOAc volumes. Solution C has more volume of HOAc than NaOAc and the vice-versa for Solution D.

Henderson- Hasselbalch equation: [A-] pH= pKa + log _______ [HA]

The equation above is used to measure pH of buffered solutions.
The unknown pH of solutions has the chemical reaction below:

CH3COOH(aq) + H2O  H3O+(aq) + CH3COO-(aq)

The Henderson-Hasselbalch equation was used to determine the pH of solutions B to D. The equation relates the concentration of the weak acid or base to its corresponding salt. It is derived from the equilibrium constant expression.

[H3O+][CH3COO-]
Ka= ____________________ [CH3COOH]

[CH3COO-] pKa= pH - log _______________ [CH3COOH]

[CH3COO-] pH= pKa + log ______________ [CH3COOH]

Calculated pH (solutions B-D):
To compute MHOAc and MNaOAc :
M1V1=M2V2

Where:
M1 = initial concentration of HOAc/NaOAc
V1 = volume of the HOAc/NaOAc used
M2 = concentration of HOAc/NaOAc in final solution
V2 = volume of final solution (10 mL)

M1V1
MHOAc = _______ V2

M1V1
MNaOAc= _______ V2

Ka= 1.8 x 10-5

Solution B: (8 mL H2O, 1 mL 0.1 M CH3COOH, 1 mL 0.1 M NaCH3COO)
MHOAc= 0.01 M
MNaOAc= 0.01 M

[0.01 M] pH = - log (1.8 x 10-5) + log ___________ [0.01 M]

pH= 4.74

Solution C: (8.9 mL H2O, 1 mL 0.1 M CH3COOH, 0.1 mL 0.1 M NaCH3COO)
MHOAc= 0.01 M
MNaOAc= 0.001 M [0.001 M] pH = - log (1.8 x 10-5) + log ___________ [0.01 M] pH= 3.74

Solution D: (8.9 mL H2O, 0.1 mL 0.1 M CH3COOH, 1 mL 0.1 M NaCH3COO)
MHOAc= 0.001 M
MNaOAc= 0.01 M [0.01 M] pH = - log (1.8 x 10-5) + log ___________ [0.001 M] pH=5.74 The experimental ionization constant of acetic acid can be computed as follows:
Observed pH of 0.01 M HOAc: 4.9
[H3O+]:
pH= - log [H3O+]
4.9= - log [H3O+]
[H3O+]=10-4.9
[H3O+]= 1.26x10-5 M
Ka of HOAc= 1.8 x 10-5 [H+][OAc-]
Ka= _____________ [HOAc] HOAc  H+ + OAc- I 0.01 0 0
C -1.26x10-5 M 1.26x10-5 M 1.26x10-5 M
_________________________________________________________________________________________________________________________________________________________
E 0.01- 1.26x10-5 M 1.26x10-5 M 1.26x10-5 M

[1.26x10-5 M]2
Ka= ________________________ = 1.59 x 10-8 [0.01 - 1.26x10-5 M]

[H3O+] of 0.01 M HOAc: 1.26x10-5 M
Calculated Ka of HOAc: 1.59 x 10-8

The experimental result showed that increasing acidity of the solutions is in the order of solution B