Transcript
Page 1: Experiment 6 Oral Report

EXPERIMENT 6Colorimetric

Determination of pH

DEL MUNDO

LARIN

SEE

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INTRODUCTION

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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

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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

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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

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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

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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

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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

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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

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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

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EXPERIMENT

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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

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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.

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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

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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

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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

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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

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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

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PART B: Colorimetric Determination of pH

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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.

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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

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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.

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• 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]

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• 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.

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• 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

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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

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[H3O+] of 0.01 M HOAc 1.00x10-5 M

To get Ka:

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[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

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[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

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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

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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-.

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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

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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

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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

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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

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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

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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

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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

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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.

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• 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

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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


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