Experiment 6 Oral Report
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Transcript of Experiment 6 Oral Report
EXPERIMENT 6Colorimetric
Determination of pH
DEL MUNDO
LARIN
SEE
INTRODUCTION
Colorimetry• any technique by which an
unknown color is evaluated in terms of standard colors
• the technique may be visual, photoelectric, or indirect by means of spectrophotometry
http://www.answers.com/topic/colorimetry#ixzz1MK1Yv3Sl
pH Indicators• also called acid-base
indicators• pH indicators are usually weak
acids or weak bases that change their color depending on their dissociation (protonation) state
• pH indicators can be used to check pH of the solution
http://www.ph-meter.info/pH-measurements-indicators
Indicator Lower pH color
pH Range(transition
interval)
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
Buffer Solutions• A buffer solution is one in which
the pH of the solution is "resistant" to small additions of either a strong acid or strong base.
• Buffers consist of a weak acid and its conjugate base or vice versa, in relatively equal and "large" quantities.
http://www.chem.purdue.edu/gchelp/howtosolveit/equilibrium/buffers.htm
McIlvaine Buffer System • A citrate/phosphate buffer system
that can be volumetrically set for pH in a wide range (2.2 to 8)
http://www.biochemlab.cn/shiji/peizhi/20993.html
Henderson-Hasselbalch Equation• An equation expressing the pH of
a buffer solution as a function of the concentration of the weak acid or base and the salt components of the buffer.
http://medical-dictionary.thefreedictionary.com/Henderson-Hasselbalch+equation
Colorimetric Analysis• Uses the variation as a means of
determining the pH since the intensity of the color of a solution changes with its concentration or pH
http://www.inc.bme.hu/en/subjects/genchem/phdet2.pdf
http://www.inc.bme.hu/en/subjects/genchem/phdet2.pdf
• By comparing the intensity of the colour of a solution of unknown concentration (or pH) with the intensities of solutions of known concentrations (or pH), the concentration of an unknown solution may be determined
EXPERIMENT
Objective• To be able to determine the pH of
an unknown solution colorimetrically
• To be able to calculate the ionization constant of a weak acid
PART A: Preparation of Buffer Solutions
A set of McIlvaine buffers were accurately prepared in test tubes of uniform sizes labeled
according to their respective pH levels.
Five drops of the appropriate indicators to use for each pH level were added to each of
the buffer solutions.
Indicator: THYMOL BLUE (1.2 – 2.8)
pH 0.2M Na2HPO4 (mL) 0.1M Citric Acid (mL)
2.2 0.20 9.80
2.4 0.62 9.38
2.6 1.06 8.91
2.8 1.58 8.42
Indicator: BROMOPHENOL BLUE (3.0 – 4.6)
pH 0.2M Na2HPO4 (mL) 0.1M Citric Acid (mL)
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
4.6 4.67 5.33
Indicator: CHLOROPHENOL RED (4.8 – 6.4)
pH 0.2M Na2HPO4 (mL) 0.1M Citric Acid (mL)
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
Indicator: BROMOTHYMOL BLUE (6.0 – 7.6)
pH 0.2M Na2HPO4 (mL) 0.1M Citric Acid (mL)
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
Indicator: PHENOL RED (6.8 – 8.0)
pH 0.2M Na2HPO4 (mL) 0.1M Citric Acid (mL)
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
7.8 9.57 0.43
8.0 9.72 0.28
PART B: Colorimetric Determination of pH
The pH of the following solutions were approximated using pH papers, then applied them
with the appropriate indicator/s.
The pH of each solutions were then confirmed by comparing their colors to standards (from part A)
applied with the same indicator.
ResultsSolution Observed
pH
A 0.01M HOAc 5
B 1mL 0.1M HOAc + 1mL 0.1M NaOAc + 8mL H20
4.8
C 1mL 0.1M HOAc + 0.1mL 0.1M NaOAc + 8.9mL H2O
3.2
D 0.1mL 0.1M HOAc + 1mL 0.1M NaOAc + 8.9mL H2O
6.8
Discussion• In the preparation of the buffer
solutions in part A, it is important to use the appropriate indicators for each buffer solutions because in colorimetric determination of pH, the indicators in buffered solutions are most effective when it is within the specific pH ranges mentioned in the table. It is at these pH ranges that the indicators show a significant change in color.
• For test tubes B, C and D, using the Henderson-Hasselbach equation, it can be inferred that as the ratio of the molarity OAc- (from NaOAc) to that of HOAc increases, the pH also increases thus making the solution less acidic.
• pH = pKa + log [OAc-]
[HOAc]
• Because it has the highest ratio of OAc- to HOAc, the solution in test tube D is expected to be the least acidic while the solution in test tube C as the most acidic.
• Common ion effect can also account for these. Because of the presence of the common ion, OAc-, there will be a suppression in the ionization of the acid thus decreasing hydrogen ion concentration and increasing the pH.
• Therefore, our observations are correct!
Solution Observed pH
A 0.01M HOAc 5
B 1mL 0.1M HOAc + 1mL 0.1M NaOAc + 8mL H20
4.8
C 1mL 0.1M HOAc + 0.1mL 0.1M NaOAc + 8.9mL H2O
3.2
D 0.1mL 0.1M HOAc + 1mL 0.1M NaOAc + 8.9mL H2O
6.8
GUIDE QUESTIONS AND ANSWERS
1.) Calculate the ionization constant of acetic acid using colorimetric data.
pH of 0.01 M HOAc 5
pH = - log [H3O+]
5 = - log [H3O+]
[H3O+]= 10-5
[H3O+] of 0.01 M HOAc 1.00x10-5 M
[H3O+] of 0.01 M HOAc 1.00x10-5 M
To get Ka:
[H3O+] of 0.01 M HOAc 1.00x10-5 M
To get Ka:
HOAc → H+ + OAc-
Initial 0.01 0 0
Change - 1.00x10-5 + 1.00x10-5 + 1.00x10-5
Equilibrium 9.99x10-3 1.00x10-5 1.00x10-5
[H3O+] of 0.01 M HOAc 1.00x10-5 M
To get Ka:
HOAc → H+ + OAc-
Initial 0.01 0 0
Change - 1.00x10-5 + 1.00x10-5 + 1.00x10-5
Equilibrium 9.99x10-3 1.00x10-5 1.00x10-5
Ka = [H+][OAc-] = (1.00x10-5 )( 1.00x10-5 )= 1.00x10-8
[HOAc] 9.99x10-3
2.) Calculate the pH of three mixtures of HOAc and NaOAc (solutions A, B and C) using the Henderson-Hasselbalch equation and compare with the observed pH. (Use the value of ionization constant of HOAc at 25°C.) Support your answers with computations.
B 1 mL 0.1 M HOAc + 1 mL 0.1 M NaOAc+ 8 mL H2O
C 1 mL 0.1 M HOAc + 0.1 mL 0.1 M NaOAc+ 8.9 mL H2O
D 0.1 mL 0.1 M HOAc + 1 mL 0.1 M NaOAc+ 8.9 mL H2O
Ka of HOAc at 25°C = 1.8x10-5
pKa = -log Ka = - log (1.8x10-5) = 4.74
pH = pKa + log [conjugate base][acid]
The acid is HOAc and its conjugate base is OAc-.
For solution B: 1 mL 0.1 M HOAc + 1 mL 0.1 M NaOAc + 8 mL H2O
MA1 VA1 = MA2 VA2
(0.1M) (1mL) = [HOAc] (10mL)
[HOAc] = 0.01M
[NaOAc] = [OAc-]
MB1 VB1 = MB2 VB2
(0.1M) (1mL) = [OAc-](10mL)
[OAc-] = 0.01M
pH = pKa + log [conjugate base]
[acid]
= 4.74 + log 0.01 M
0.01 M
= 4.74 + 0
pH = 4.74
For solution B: 1 mL 0.1 M HOAc + 1 mL 0.1 M NaOAc + 8 mL H2O
For solution C: 1 mL 0.1 M HOAc + 0.1 mL 0.1 M NaOAc + 8.9 mL H2O
MA1 VA1 = MA2 VA2
(0.1M) (1mL) = [HOAc] (10mL)
[HOAc] = 0.01M
[NaOAc] = [OAc-]
MB1 VB1 = MB2 VB2
(0.1M) (0.1mL) = [OAc-](10mL)
[OAc-] = 0.001M
pH = pKa + log [conjugate base]
[acid]
= 4.74 + log 0.001 M
0.01 M
= 4.74 + (-1)
pH = 3.74
For solution C: 1 mL 0.1 M HOAc + 0.1 mL 0.1 M NaOAc + 8.9 mL H2O
For solution D: 0.1 mL 0.1 M HOAc + 1 mL 0.1 M NaOAc + 8.9 mL H2O
MA1 VA1 = MA2 VA2
(0.1M) (0.1mL) = [HOAc] (10mL)
[HOAc] = 0.001M
[NaOAc] = [OAc-]
MB1 VB1 = MB2 VB2
(0.1M) (1mL) = [OAc-](10mL)
[OAc-] = 0.01M
pH = pKa + log [conjugate base]
[acid]
= 4.74 + log 0.01 M
0.001 M
= 4.74 + 1
pH = 5.74
For solution D: 0.1 mL 0.1 M HOAc + 1 mL 0.1 M NaOAc + 8.9 mL H2O
SolutionpH
observed calculated
0.01 M HOAc 5 ----
1 mL 0.1 M HOAc + 1 mL 0.1 M NaOAc + 8 mL H2O (A) 4.8 4.74
1 mL 0.1 M HOAc + 0.1 mL NaOAc + 8.9 mL H2O (B) 3.2 3.74
0.1 mL 0.1 M HOAc + 1 mL 0.1 M NaOAc + 8.9 mL H2O (C) 6.8 5.74
A
B
C
D
Conclusion• The pH of a solution can be
approximated with the use of colorimetry and pH indicators. But it is also important to have to have proper knowledge on which indicator to be used on certain pH ranges and their color transitions for a successful colorimetric analysis.
• It is strongly advised to have accurate measurements for the preparation of buffer solutions to have an efficient standard and also for the solutions that will be used for colorimetric analysis.
Recommendation
References• Lemay, H., Brown, T., Bursten, B., & Burdge, J. (2004).
Chemistry: The Central Science. New Jersey: Pearson Education South Asia Pte Ltd.
• http://www.answers.com/topic/colorimetry#ixzz1MK1Yv3Sl• http://www.ph-meter.info/pH-measurements-indicators• http://www.chem.purdue.edu/gchelp/howtosolveit/
equilibrium/buffers.htm• http://www.biochemlab.cn/shiji/peizhi/20993.html• http://medical-dictionary.thefreedictionary.com/Henderson-
Hasselbalch+equation• http://www.inc.bme.hu/en/subjects/genchem/phdet2.pdf