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EXPT 1

By admcaraos Feb 28, 2015 1832 Words
pH MEASUREMENT AND BUFFER PREPARATION

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2A—Pharmacy, Faculty of Pharmacy, University of Santo Tomas

Canare, Ann Shery; Caraos, Aeraille Diane; Carillo, Pauline Mari; Carreon, Leslie Ann; Co, Alexandra Kay

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ABSTRACT. The regulation of internal pH is a physiological function of major importance for all organisms.

The quantity pH is intended to be a measure of the activity of hydrogen ions in solution. Buffers, on the other hand, are solutions containing a substantial concentration of both members of a weak conjugate acid-base pair; and it can resist moderate pH changes. Biochemists and other life scientists also need to control pH in their experiments and frequently prepare aqueous buffer solutions for enzyme assays, extraction solvents, incubation media, etc. The relationship between pH and the ratio of the concentrations of the buffer components is given by the Henderson-Hasselbalch equation. In the experiment, electrometric determination and colorimetric determination of pH was performed. In electrometric determination, a buffer solution was prepared by adjusting the pH of the sample (i.e. Dove shampoo) to 5. In colorimetric determination, different acid-base indicator were added to different buffer solutions, and the pH where the change in colour were seen were noted. The data shown later would show a different change in colour for every pH, which indicates that the hydrogen or hydroxide ions change in an aqueous solution due to the pH indicators.

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!I.
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INTRODUCTION

pH is often heard and always computed
every Chemistry class but is it only important for
Chemistry students? Of course not. pH is important
and will always be important to everyone. From the
soil to the food and most importantly our body is in
relation to pH. One example of the importance of pH
is our blood, which is slightly basic, and if blood pH
is not within the normal values, certain complications
may arise or worse the body dies. Another one is the
buffer system within our stomach. It resist drastic
changes in pH within the stomach especially when
we are hungry. That is why everybody needs to have
a right amount of knowledge about pH.

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In this experiment, the students were tasked
to prepare different buffer solutions. With the use of
different liquid indicators, pH of the buffers and
samples will be determined colorimetrically. It will
also be measured electrometrically using the pH
meter.

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II. MATERIALS AND METHODOLOGY
! 1. Preparation of Reagents

First, the amounts needed to prepare the
reagents were computed using the dilution
factor as well as the formula for getting the
molar concentration. For 500 mL of 6.0M

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HCl, 245.30 mL was diluted while for 500
ml of 6.0 M NaOH, 120 g were used.
2. Buffer Preparation
Acetate Buffer was assigned to the group.
To prepare the buffer solution, the amounts
of solute needed were computed, which is
1.43 mL ofCH3COOH and 0.67g of NaOH
pellets were combined and diluted with
distilled water. The buffer solution was then
placed in a 250 mL volumetric flask and
was labeled.
3. Electrometric Determination of pH
The pH meter was calibrated first at pH 4, 7
and 10. For the first part, the pH of the
prepared buffered solutions was measured,
then the distilled water then the assigned
sample which is for our group, we chose
shampoo to be our sample. For each of the
pH measured, the [H+] were calculated
using the formula 10-pH. Then, the pH of the
prepared buffer solution was measured. The
buffer solution was then manipulated to pH
5.0 which is the desired pH using the 6.0M
NaOH solution to make it more basic and
6.0M HCl solution to make it more acidic.
4. Colorimetric Determination of pH
Each buffer solution having a different pH,
from pH 2-11 including pH 7.5 were tested

using all the acid-base indicators which are:
thymol blue, bromophenol blue,
bromocresol green, bromecresol purple,
phenol red, methyl red, methyl orange and
phenolphthalein. Eleven test tubes were
prepared for every indicator and were
labeled with pH 2-11. In each test tube, 10
mL of the sample solution was used. Then,
2-3 drops of the indicator was added in
each. The results were then observed to
determine the pH range of the indicator.

The electrometric method is generally accepted as the
most accurate method employed for the
determination of pH measurements. The pH of the
samples and buffer were determined
electrometrically by using a pH meter that measures
the electrical potential produced by the solution,
compares it with the voltage of a known solution, and
uses the potential difference to deduce the pH. The
pH of the acetate buffer was adjusted to its desired
value by addition of small amounts of 1.0 M HCl or
1.0 M NaOH. Presence of hydrogen and hydroxide
ions, cause a decrease and increase in the pH
readings, respectively. The hydrogen concentrations
of the samples and buffer were calculated by taking
the antilog of their corresponding pH value.

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!III. RESULTS AND DISCUSSIONS
!Table 1. Results of the Electrometric Determination

!Conclusion:

The electrometric method of
determining the pH should be used where any
considerable accuracy is required.

of pH

Sample

pH

[H

Distilled Water

7.50

3.16 x 10

Dove Intense
Repair
Shampoo

5.40

3.98 x 10

5.00

1.0 x 10

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Buffer

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Acetate Buffer

Table 2. Results of Colorimetric Determination Using a Range of Acid-Base Indicators

Acid – Base
Indicator

pH
2.0

3.0

4.0

5.0

6.0

7.0

7.5

8.0

9.0

10.0

11.0

Thymol

Peach

Dull
Yellow

Dull
Yellow

Light
Yellow

Light
Yellow

Light
Yellow

Light
Yellow

Yellow

Blue
Green

Dull
Blue

Light
Blue

Bromophenol
blue

Yellow

Light
Green

Light
Blue

Bluish
purple

Light
purple

Light
purple

Blue

Indigo

Dark
purple

Dark
purple

Dark
Indigo

Bromocreson
green

Light
Yellow

Light
Green

Light
Green

Light
green

Light
green
blue

Light
Blue

Light
Blue

Blue

Blue

Blue

Blue

Bromocresol
purple

Yellow

Yellow

Yellow

Yellow

Dark
Yellow

Violet

Violet

Violet

Violet

Violet

Violet

Phenol red

Yellow

Yellow

Yellow

Yellow

Dark
Orange

Dark
Orange

Red
Orange

Cherry
Red

Dark
Pink

Magenta

Dark
Magenta

Methyl red

Light
pink

Pink

Pink

Pink

Dull
Yellow

Yellow

Yellow

Yellow

Yellow

Yellow

Yellow

Methyl orange

Red
Orange

Dark
Orange

Dark
Orange

Dark
Orange

Orange

Orange

Orange

Light
orange

Light
Orange

Light
Orange

Phenolphthalei
n

Colorle
ss

Colorle
ss

Colorle
ss

Colorle
ss

Colorle
ss

Colorle
ss

Colorle
ss

Light
Pink

Hot
pink

Hot
pink

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Colorimetric determination of pH showed the
varying color changes acid-base indicator undergoes
when added to a solution of a certain pH. This is due
to the fact that acid-base indicators also known as pH
indicators is either a weak acid or a weak base that
exhibits a color change as the concentration of
hydrogen (H+) or hydroxide (OH-) ions changes in an
aqueous solution. Certain organic substances change
color in dilute solution when the hydrogen ion
concentration reaches a particular value. [6]
Let’s take methyl orange as an example, the
equilibrium in a solution of the acid-base indicator
methyl orange, a weak acid, can be represented by
the equation

Light
Orang
e
Pink

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change in colour is visible for any further increase in
the hydrogen ion concentration.
Addition of a base to the system reduces the
hydrogen ion concentration and shifts the equilibrium
toward the yellow form. At a pH of 4.4 about 90% of
the indicator is in the yellow ionic form, and a further
decrease in the hydrogen ion concentration does not
produce a visible colour change. The pH range
between 3.1 (red) and 4.4 (yellow) is the colourchange interval of methyl orange; the pronounced colour change takes place between these pH values.
In general, the colour-change interval of an indicator
is the pH range, where pronounced colour change
takes place; the borders of this interval can be
estimated by pKa-1 and pKa+1.

!The basis for what the chemist calls colorimetric
The anion of methyl orange is yellow, and the nonionized form is red. If acid is added to the solution, the increase in the hydrogen ion concentration shifts
the equilibrium toward the red form in accordance
with the law of mass action.

The indicator's colour is the visible result of the ratio
of the concentrations of the two species A− and HA.
For methyl orange:

When [H+] has the same numerical value as Ka, the
ratio of [A-] to [HA] is equal to 1, meaning that 50%
of the indicator is present in the red acid form and
50% in the yellow ionic form, and the solution
appears orange in colour. When the hydrogen ion
concentration increases to a pH of 3.1, about 90% of
the indicator is present in the red form and 10% in
the yellow form, and the solution turns red. No

analysis is the variation in the intensity of the colour
of a solution with changes in concentration (or pH).
The colour may be due to an inherent property of the
constituent itself (e.g. MnO4 − is purple) or it may be
due to the formation of a coloured compound as the
result of the addition of a suitable reagent (e.g.
indicator). 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.

!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.

!As shown in the table above, we see the color change
of the varying pH, and if we compare it to the

standard it is in range with the results. Like in
phenolphthalein, colorless in acidic pH, while turning
pink in basic pH.

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Fig. 1. Acid-base indicators
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!IV. REFERENCES
!!
Kahn, George, (april 29, 1926). The Comparison of

the Electrometric and Colorimetric Methods for
Determination of the pH of Gastric Contents. . (e.g.
2), pp.75-80

!Stern, R. (n.d.). Buffer Preparation. Retrieved from
http://www.unf.edu/~rstern/BufferLabspr07.pdf
!Buffer Solutions. (n.d.). Retrieved from http://
www.chem.uci.edu/~lawm/4-10-13.pdf
!Measurement of pH: Dsefinition, Standards, and

Procedure. (2002). Retrieved from http://
w w w. i u p a c . o rg / p u b l i c a t i o n s / p a c / 2 0 0 2 / p d f / 7411x2169.pdf

!Preparation

of buffer solution & Colorimetric
Determination of pH. Retrieved from http://
iris.inc.bme.hu/en/subjects/genchem/phdet2.pdf

!
!

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