Colorimetric Determination

of pHBausa, D.M.

Uy, A.R.

objectives

1)Determine colorimetrically the pH of an unknown solution.

2)Calculate the ionization constant of a weak acid.

introduction

Colorimetry

The basis for colorimetric analysis is the variation in the intensity of the colour of a solution with changes in

concentration (or pH). The concentration of an unknown solution may be determined by comparing the

intensity of its color with the intensities of the standard solutions of known

concentrations (or pH).

pHThe term pH stands for

“potential” of “Hydrogen”.  It is the amount of hydrogen

ions in a particular solution. The more ions, the more acidic

the solution.  The fewer ions the more basic the solution.

pH= -log[H+]pOH= -log[OH-]

pH+pOH=14

Acid-base Indicators

Acid-base indicators are weak acids/bases that have different colors in their ionized and unionized forms. The color change in the indicator occurs over a range of hydrogen ion concentrations. This range is termed the color change interval and is expressed as a pH range.

Acid-base Indicators

The pH of the solution at its turning point is called the pKln and is the pH at which half of the indicator is in its acid form and the other

half in the form of its conjugate base.

Buffer solutions

This are solutions consisting of weak conjugate acid-base pairs that prevent drastic pH changes by

maintaining a steady acidity level even when a strong acid or strong

base is added through neutralization by the acidic & basic species.

McIlvaine Buffer

A buffer system which uses a citrate and a phosphate usually citric acid

and Na2HPO4 to volumetrically set for pH in a wide range

Henderson-Hasselbalch Equation

An equation used for measuring the pH of buffered solutions

Experimental

Part A.

Prepare the set of McIlave buffers by mixing the designated amount of reagents.

Label the test tubes according to the pH of its solution.

Add 5 drops of indicator to the test tube whose pH is in the range of the indicator.

Thymol blue (pH range: 1.2-2.8)

pH 0.2 M Na2HPO4

0.1 M Citric Acid

2.2 0.20 9.80

2.4 0.62 9.38

2.6 1.06 8.91

2.8 1.58 8.42

pK= 1.65

Bromophenol blue (pH range: 3.0-4.6)

pH 0.2 M Na2HPO4

0.1 M 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 3.25 6.15

4.2 4.14 5.86

4.4 4.41 5.59

4.6 4.67 5.33pK= 4.10

Chlorophenol red (pH range: 4.8-6.4)

pH 0.2 M Na2HPO4

0.1 M 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 pK= 6.25

Bromothymol blue (pH range: 6.0-7.6)

pH 0.2 M Na2HPO4

0.1 M 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 0.63pK= 7.30

Phenol red (pH range: 6.8-8.4)

pH 0.2 M Na2HPO4

0.1 M Citric Acid

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

pK= 8.00

Prepare the following solutions in 10-mL test tubes

0.01 M HOAc

[Solution A]1 mL 0.1 M

HOAc + 1 mL 0.1 M NaOAc + 8 mL

H20

[Solution B]1 mL 0.1 M

HOAc + 0.1 mL 0.1 M NaOAc +

8.9 mL H20

[Solution C]0.1 mL 0.1 M

HOAc + 1 mL 0.1 M NaOAc + 8.9

mL H20

Add 2 drops of the pH indicator designated to the solution.

Compare the color of the solution to the previously made McIlave buffers and note the observed pH.

Part b.

Results

A B C D

Part b.

ResultsPart b.SOLUTION OBSERVED pH CALCULATED

pH0.01 M HOAc 3.38 xxxxxxx

1ml 0.1 M HOAc + 1 ml 0.1 M NaOAc + 8 ml H20 4.8 4.74

1ml 0.1 M HOAc + 0.1 ml 0.1 M NaOAc + 8.9 ml

H20

4.0 3.74

0.1ml 0.1 M HOAc + 1 ml 0.1 M NaOAc + 8.9 ml

H20

5.8 5.74

As the ratio of the [NaOAc]/[HOAc] increased, the pH level also increased, making the solution less acidic.

Solution C was the least acidic because it had the highest ratio of NaOAc to HOAc, while solution B was

the most acidic because it had the lowest ratio of NaOAc to HOAc.

Part b.

Solution:

HOAc + H2O <—> H3O + Oac

NaOAc <—> Na + OAc

Part b.

Solution:

HOAc + H2O <—> H3O + OAc

NaOAc <—> Na + OAc

The presence of a OAc caused a common ion effect. The production of more OAc caused a reverse shift which resulted to a decrease in the concentration of H3O ions. This decrease

caused the increase in pH level.

Results1. Calculate the Ionizaiton Constant of Acetic Acid.

pH of 0.01 M HOAc 3.38[H3O+] of 0.01 M HOAc 4.17x 10^-4

Calculated Ka of HOAc 1.8 x 10^ -5

pH= -log [H3O +]

3.38=-log [H3O +]

[H3O +]= 4.17 x 10-4

Results1. Calculate the Ionizaiton Constant of Acetic Acid.

Ka= (4.17 x 10-4)2

0.00958Ka = 1.8x 10 -5

Results2. Calculate the pH of the 3 mistures of HOAc & NaOAc.

Henderson Hasselbach equation:

Ka= 1.8 x 10^-5pKa= -log Ka = 4.74

Results

SOLUTION A

2. Calculate the pH of the 3 mistures of HOAc & NaOAc.

Results2. Calculate the pH of the 3 mistures of HOAc & NaOAc.

SOLUTION B

Results2. Calculate the pH of the 3 mistures of HOAc & NaOAc.

SOLUTION C

Conclusion

The pH of any solution may be determined through colorimetry and pH indicators though the pH meter will give the most accurate pH. The pH range of specific

indicators must be known before they are used to be accurate in the colometric

analysis.

Conclusion

Buffer solutions are able to resist drastic changes in pH but a small change may be observed

because of common ion effect. The ratio of strong to weak electrolyte affects the resulting pH

because it determines the decrease or increase in H+ ions in the solution. The Henderson-

Hasselbalch equation is used in determining the pH of a buffer solution.

Recommendation

the buffer solutions should be correctly and accurately prepared

to achieve the accurate pH needed in the colorimetric

determination of pH

Sources:

• http://chemistry.about.com/od/acidsbases/a/Acid-Base-Indicators.htm

• http://www.frequencyrising.com/pH.htm• http://www.ch.ic.ac.uk/vchemlib/course/indi/

indicator.html• http://www.chembuddy.com/?left=pH-

calculation&right=pH-buffers-henderson-hasselbalch• http://www.inc.bme.hu/en/subjects/genchem/

phdet2.pdf• http://www.docbrown.info/page07/equilibria6a.htm