PH MEASUREMENT AND BUFFER PREPARATION

Anna Katrina M. Donato, Abigail O. Dy, Keith Brian T. Enriquez, Debbie Marie R. Fermin

Group 3 2E- Medical Technology Biochemistry Laboratory

ABSTRACT

A 500 ml Phosphate buffer solution of pH 7.0 with a molar concentration of 5M was prepared using 10.67057611 g of Primary sodium phosphate monohydrate (NaH2PO4.H2O) and 12.77743171 g of Secondary sodium phosphate heptahydrate (Na2HPO4.7H2O). The pH of the buffer was adjusted to the desired value by adding in either portions of 1.0M HCl or 1.0 M NaOH while being monitored by a pH meter. The buffers prepared with varying pH were subjected in Colorimetric determination using different acid-base indicators. The sample, distilled water with pH 3, was also subjected to Colorimetric determination giving the result colors of yellow-orange for Thymol blue, light green for Bromphenol blue, blue for Bromcresol green, yellow for Bromcresol purple, yellow for Phenol red, pink for Methyl red, orange for Methyl orange, and colorless for Phenolphthalein.

INTRODUCTION

The term pH refers to a measure of the hydrogen ion concentration of a solution. Solutions with a high concentration of hydrogen ions have a low pH and solutions with low concentrations of H+ ions have a high pH. Therefore, pH is also used as a measure of the acidity or basicity of a solution. Mathematically, pH is expressed as the negative log in base 10 of the aquated hydrogen ion concentration.

pH = -log [H+]

One way of measuring pH is by using a device called pH meter. A pH meter consists of a pair of electrodes connected to a meter capable of measuring small voltages, on the order of millivolts. A voltage, which varies with the pH, is generated when the electrodes are placed in a solution. This voltage is read by the meter, which is calibrated to give the pH.

A buffer is a solution which contains a weak conjugate acid-base pair that can resist drastic changes in pH upon the addition of small amounts of a strong acid or base. A buffer resists changes in pH because it contains both an acidic species to neutralize OH - ions and a basic one to neutralize H+ ions. It is a requirement though that the components of a buffer must not consume each other, that’s why buffers are often prepared by mixing a weak acid or a weak base with a salt of that acid or base.

EXPERIMENTAL

A. Compounds tested (or samples used)

Distilled water, Primary sodium phosphate monohydrate (NaH2PO4.H2O), Secondary sodium phosphate heptahydrate (Na2HPO4.7H2O), Acid- base indicators ( Thymol blue, Bromphenol blue, Bromcresol green, Bromcresol purple, Phenol red, Methyl red, Methyl orange, Phenophthalein)

B. Procedure

1. Preparation of Reagents

250 ml of 0.5 M HCl was prepared from 1.0 M HCl. The container was labeled properly.

(1.0 M HCL)(500 ml) x (0.5 M HCL) (x ml)

X= 250 ml HCl

Figure1.Calculation of volume of 0.5 M HCl

2. Preparation of buffer

The buffer solution was prepared using the following guidelines:

Table 1. Guideline for buffer preparation

Volume(L)

Concentration(M)

Buffer Solution Desired pH

0.500 0.5 Phosphate 7.0

Primary sodium phosphate monohydrate (NaH2PO4.H2O) and Secondary sodium phosphate heptahydrate (Na2HPO4.7H2O) were used in preparing the buffer. The container was labeled properly.

Given:

pH= 7 Weak acid= HPO4-2

0.5 M Conjugate base= H2PO4-

250 ml pKa= 7.21

Handerson-Haselbach:

pH= pKa + log [WA]/ [CB]

pH = 7.21 + log [HPO4-2]/ [H2PO4-]

log-1 (7-7.21)= [HPO4-2]/ [H2PO4-]

0.616595001/ 1= [HPO4-2]/ [H2PO4-]

Total theoretical moles of buffer= 1.616595001

Total actual moles of buffer= (0.5M)(0.25L)= 0.125 moles

Actual moles [HPO4-2]: (0.616595001/1.616595001) = ( x/ 0.125)

= 0.047676984 moles HPO4-2

Actual moles [H2PO4-]: (1/ 1.616595001)= (x/ 0.125)

= 0.077323015 moles H2PO4-

Grams of HPO4-2:

Actual moles x MW = 0.047676984 moles x 268 g/mol

= 12.77743171 g HPO4-2

Grams H2PO4-:

Actual moles x MW= 0.077323015 moles x 138 g/mol

= 10.67057611 g H2PO4-

Figure 2.Computation for buffer preparation

1. Electrometric Determination of pH

The pH meter is calibrated. The pH of 20 ml portion of distilled water was measured and the [H+] was calculated. The pH of the buffer solution was adjusted to the desired pH by adding in portions of either 1.0 M HCl or 1.0 M NaOH while being monitored by the pH meter.

2. Colorimetric Determination of pH

a. Preparation of Color Standards Using the Buffer Solutions

Six vials/ test tubes were prepared and labeled with the pH of the buffer and acid-base indicator to be added. 5 ml of a buffer of a certain pH and 2 drops of an acid-base indicator was added into the vial. The mixture was shaken and the resulting color was noted. This procedure was performed on all the buffers prepared. Another set of 6 vials/ test tubes were prepared. The procedure was repeated but another acid-base indicator was used.

+2 drops of acid-base indicator to a spot

+10 drops Phosphate buffer to each spot

mix note color

Figure 3. Colorimetric determination using spot plate

Table 2. Acid-base indicators used in Colorimetric determination of pH

Acid-base IndicatorsThymol blue

Bromphenol blueBromcresol greenBromcresol purple

Phenol redMethyl red

Methyl orangePhenolphthalein

b. Determination of the pH of samples

Two drops of an acid-base indicator was added to 5 ml of distilled water. The mixture was shaken and the resulting color was noted. This procedure was repeated using each acid-base indicator.

RESULTS AND DISCUSSION

1. Electrometric Determination of pH

The pH meter showed accurate readings of the pH of the buffer. It showed fluctuations in readings with the slightest addition of HCl and NaOH. It displayed sensitivity to a small amount of [H+] and [OH-] ions. The pH meter therefore, is more accurate in reading pH levels compared to a pH paper.

2. Colorimetric Determination of pHTable 3. Results of Calorimetric Determination of pH

Acid-Base indicator

2.0 3.0

pH

7.0 7.5 8.0 12.0

Distilled Water

Thymol blue red yellow-orange

yellow yellow yellow blue yellow-orange

Bromophenol blue

yellow yellow-green

blue blue blue blue-violet

light green

Bromcresol green

yellow dark yellow

blue blue blue-green

blue-green

blue

Bromcresol purple

yellow yellow violet violet blue-violet

violet yellow

Phenol red yellow yellow yellow-orange

red red red violet

yellow

Bogen indicator ------ ------ ------ ------ ------ ------ ------

Methyl red pink pink yellow yellow yellow pink pink

Methyl orange red red-orange

yellow-orange

orange orange orange orange

Phenolphthalein

colorless colorless colorless colorless light-pink

red-violet

colorless

pH 3.0

For pH 2.0, 3.0, 7.0, 7.5, 8.0, 12.0, and distilled water, the following are the color results for acid-base indicator Thymol blue: red, yellow-orange, yellow, yellow, yellow, blue, and yellow-orange.

For pH 2.0, 3.0, 7.0, 7.5, 8.0, 12.0, and distilled water, the following are the color results for acid-base indicator Bromophenol blue: yellow, yellow-green, blue, blue, blue, blue-violet, and light green.

For pH 2.0, 3.0, 7.0, 7.5, 8.0, 12.0, and distilled water, the following are the color results for acid-base indicator Bromcresol green: yellow, dark-yellow, blue, blue, blue-green, blue-green, and blue.

For pH 2.0, 3.0, 7.0, 7.5, 8.0, 12.0, and distilled water, the following are the color results for acid-base indicator Bromcresol purple: yellow, yellow, violet, violet, blue-violet, violet, and yellow.

For pH 2.0, 3.0, 7.0, 7.5, 8.0, 12.0, and distilled water, the following are the color results for acid-base indicator Phenol red: yellow, yellow, yellow-orange, red, red, red-violet, and yellow.

For pH 2.0, 3.0, 7.0, 7.5, 8.0, 12.0, and distilled water, the following are the color results for acid-base indicator Methyl red: pink, pink, yellow, yellow, yellow, pink, and pink.

For pH 2.0, 3.0, 7.0, 7.5, 8.0, 12.0, and distilled water, the following are the color results for acid-base indicator Methyl orange: red, red-orange, yellow-orange, orange, orange, orange, and orange.

For pH 2.0, 3.0, 7.0, 7.5, 8.0, 12.0, and distilled water, the following are the color results for acid-base indicator Phenolphthalein: colorless, colorless, colorless, colorless, light-pink, red-violet, and colorless.

Colorimetric determination of pH showed the varying color changes an acid-base indicator undergoes when added to a solution of certain pH. This property of an acid-base indicator can therefore be used to identify different substances by narrowing their pH range. For example: Using Bromophenol blue as an acid-base indicator, a solution turned yellow-green. By such observation, one can say that the pH of the solution is 3.0. This can help in the identification of a substance since different substances exhibit different pH levels. Acid-base indicators can also be used to narrow down the pH range of a substance. For example: A resulting color of blue-violet using acid-base indicator Bromcresol green indicates a pH>8.0, and a resulting color of violet in

acid-base indicator Bromcresol purple indicates a pH<7.5. Therefore, we can estimate that the pH of the substance must be between 7.5 and 8.0.

Acid-base indicators also show molecular characteristics of a substance. Color changes in molecules can be caused by changes in electron confinement. More confinement makes the light absorbed bluer (darker), and less makes it redder (lighter).

REFERENCES

From books:

[1] Bursten, B.E., Brown, T.L., LeMay, H.E.(2004). Chemistry: The Central Science. 9th ed. Singapore: Pearson Education Inc.

[2] Campbell, M.K., Farell, S.O.(2009). Biochemistry. 6th ed. Philippines: Cengage Learning Asia Pte. Ltd.

From the internet:

[1] Biology online

http://www.biology-online.org/dictionary/Buffer 1/10/ 10

[2] Brooklyn academic

http://academic.brooklyn.cuny.edu/biology/bio4fv/ page/ph_def.htm 01/10/10

[3] Harper College

http://www.harpercollege.edu/tm-ps/chm/100/dgodambe/thedisk/ph/abind.htm 01/12/10