1

INTRODUCTION

Titration is a process where a solution of known concentration is reaction with a

solution with unknown concentration in order to know more about the unknown solution

(Daintith, 2008). During this process, a burette is used together with a retort stand to clamp

the burette so that it will be in a static position to avoid any errors. The solution with a known

concentration will be filled in the burette and the solution with unknown concentration will

usually placed in a small beaker. The solution with unknown concentration will be positioned

at the bottom of the burette and will be titrated with the solution in the burette. (Hoogendijk,

1999).

Acid-base titration is one of the widely used methods. In this experiment, titration was used

to determine the value of acid ionization constant, Ka of an unknown acid solution. The

prediction of Ka value plays a vital role in chemistry especially in life and material sciences,

pharmaceutical industry and other R&D oriented enterprises (BogusBaw et al., 2015). The

unknown acid solution was titrated with 0.1 M NaOH solution. One is placed in a burette and

the other one which is to be titrated is placed in a beaker and is placed below the mouth of the

burette (“Definition of Titration”, 2014). After getting all the data needed, titration curves

were plotted. Titration curve is a graph of measured pH values obtained from pH meter

readings versus volume of titrant being added. Figure 1.1 and 1.2 shows a general acid-base

titration set up and titration Acid-base titration curve of weak acid treated with a strong base

respectively.

Figure 1.1: General acid-base titration set up.

2

Figure 1.2: Acid-base titration curve of weak acid treated with a strong base.

It is important to know the equivalence point of every titration curves as it will be

used in the calculation of getting the value of Ka. Equivalence point occurs when the moles

of acid in the solution equals to the moles of base added in the solution. It can be plotted from

the region of the curve where large and noticeable change in pH with a relatively small

change in volume of titrant occurs.

OBJECTIVES

To determine the Ka value of a weak acid.

To identify the name of the weak acid used in this experiment.

3

THEORY

pH : A measure of Acidity

pH scale is basically a logarithmic scale for indicating a the acidity or alkalinity of a solution

(Daintith et al., 2010). The term pH is originally come from a combination of “p” for power

and “H” for the symbol of the element Hydrogen (anonymous, 2013). In this experiment, the

pH of the unknown acid solution was taken until it has reached equivalence point. The pH of

the solution is defined as the negative logarithm of the hydrogen ion concentration (Chang R.

& Goldsby K.A., 2013). Given the expression for pH :

pH = -log [H3O

+] (Equation 1.1)

[H3O

+] = 10

-pH (Equation 1.2)

The concentration of hydronium ion in a solution can be defined by reversing the

mathematical operation based on Equation 1.2. Different acid and bases has different range of

pH value. Acidic solutions have pH less than seven while basic solutions pH values are more

than seven. pH at 7 is for neutral solutions (Chang R. & Goldsby K.A., 2013). Acids and

bases are categorized into strong and weak. Strong acids will have small values of pH while

strong bases will have high values of pH and vice versa. For example , hydrochloric acid

which is a strong acid has a pH value of 1.1. This acid has a high concentrations of hydrogen

ions. Thus , pH decreases as [H+] increases (Daintith et al., 2010).

Acid ionization constant , Ka

Acid ionization constant is basically the equilibrium constant for the ionization of an acid.

We can tell how strong is the acid by determining the Ka value. The larger the Ka, the higher

the concentrations of hydrogen ions, thus the stronger the acid (Chang R. & Goldsby K.A.,

2013). The equilibrium expression for acid ionization are :

Ka =

(Equation 1.3)

4

Ka = [H3O+] (Equation 1.4)

pKa = -log Ka (Equation 1.5)

Ka = 10-pka

(Equation 1.6)

When the solution has achieved the equilibrium point, that is [HA] and [A-] will be the same,

so these two can be cancelled off giving the expression in Equation 1.4. So, the pKa is equal

to the pH. (anonymous, 2004). The pKa can be obtained from the half equivalence point of

the titration.

APPARATUS

Equipment

A 250 mL volumetric flask, 50 mL beakers, a burette and its clamp, a stopcock, a

spatula, 10 mL volumetric cylinder , a pH meter.

Materials

0.1 M sodium hydroxide solution, 90 mL of 0.1 M unknown acid solutions, distilled

water.

Figure 1.3: Apparatus set up for titration method.

0.1 M NaOH

Unknown acid

5

PROCEDURE

PART ONE: Titration Method

1. The burette was rinsed with distilled water.

2. 0.1 M of sodium hydroxide solution (NaOH) was used as a titrant. Small amount of

sodium hydroxide solution was poured into the burette and drained into a waste

beaker to remove the presence of air bubbles from the tip of the burette and stopcock.

3. The burette is then fully filled with the sodium hydroxide solution. The sodium

hydroxide solution was poured up to to 0.0 mL line or below.

4. An unknown acid with 0.1M and 10 mL was placed in a beaker.

5. The initial burette reading was recorded.

6. The initial pH of the unknown acid was recorded. The unknown acid in the beaker is

place at the bottom of the tip of the burette and the burette is adjusted at the right

position so that the acid solution can be titrated correctly.

7. The acid solution was titrated with the NaOH solution and for every 1 mL drop of

NaOH solution, the pH value of the unknown acid was recorded.

8. The NaOH solution was added until the pH value has reached basic region and has

leveled off.

9. The titration of the acid sample in the beaker was discarded.

10. The titration process was repeated for another twice.

11. A graph of pH value and the volume of NaOH solution were plotted for each titration.

The equivalence point from each graph was determined.

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PART TWO: pH value method

1. An unknown acid acid of 20 mL was placed in a beaker.

2. The pH value for the solution was observed until it has reached a constant value

and was recorded.

3. The steps were repeated for another twice.

RESULT

PART 1

Table 1.1: Determination of the Ka value of a weak acid by titration with NaOH

First Trial Second Trial Third Trial Volume

(mL)

pH Volume

(mL)

pH Volume

(mL)

pH

0 4.47 0 4.33 0 4.32

1 4.74 1 4.56 1 4.54

2 4.90 2 4.78 2 4.73

3 5.03 3 4.96 3 4.93

4 5.21 4 5.13 4 5.09

5 5.35 5 5.27 5 5.27

6 5.55 6 5.44 6 5.43

7 5.71 7 5.62 7 5.58

8 5.94 8 5.86 8 5.85

9 6.32 9 6.25 9 6.22

10 10.79 10 10.64 10 10.60

10.1 10.82 10.1 10.68 10.1 10.64

10.2 10.84 10.2 10.75 10.2 10.65

10.3 10.85 10.3 10.76 10.3 10.69

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PART TWO

Table 1.2: Identifying Ka weak acid by pH value

Volume (mL) pH

20 4.31

20 4.28

20 4.27

PART ONE: TITRATION CURVES

Figure 1.4: Titration curve for the first titration.

0

2

4

6

8

10

12

0 5 10 15

pH

Volume of NaOH (mL)

pH

EQUIVALENCE POINT

HALF EQUIVALENCE POINT

9.51

5.38

8

Figure 1.5: Titration curve for the second titration.

Figure 1.4: Titration curve for the third titration.

0

2

4

6

8

10

12

0 5 10 15

pH

Volume of NaOH (mL)

pH

HALF EQUIVALENCE

EQUIVALENCE

0

2

4

6

8

10

12

0 5 10 15

pH

Volume of NaOH (mL)

pH

EQUIVALENCE POINT

9.51

5.39

9.65

5.41

HALF EQUIVALENCE POINT

9

SAMPLE CALCULATIONS

PART ONE: Determination of the Ka value of a weak acid by titration with NaOH

Titration 1

Equivalence point

6.32 + 10.79 = 8.6

2

Volume = 9.51

Half equivalence point

9.51= 4.80

2

pH= 5.40 = pka

Ionization constant expression, Ka

pka = -log Ka

Ka = 10-pka

= 10-5.40

= 3.98 x 10-6

Table 1.3: Tabulated data for Part 1 calculation.

Titration 1 Titration 2 Titration 3

Equivalence point 9.51 9.51 9.65

Half equivalence

point (mL)

4.80 4.76 4.83

pKa at half

equivalence point

5.40 5.39 5.41

Ka value 3.98 x 10-6

4.07 x 10-6

3.89 x 10-6

Average Ka value 3.98 x 10-6

10

Based on Appendix A; when Ka = 3.90 x 10–6

, the corresponding acid is Potassium Hydrogen

Phthalate (KHP). Thus the unknown acid is KHP because the average Ka value is 3.98 x 10-6

which is approximately 3.90 x 10–6

.

Percentage error,

= Experimental value – theory x 100%

Theory

= – –

– = 2.05 %

PART TWO: Identifying Ka of weak acid by pH value by using ICE method

First Trial

pH value of unknown acid solution = 4.31

pH = -log [H3O

+]

[H3O

+] = 10

-pH

= 10-4.31

= 4.90 x 10-5

M

Ka =

When the solution has achieved the equilibrium point, [HA] and [A-] will be the same, so

these two can be cancelled off. Thus, the expression for calculating Ka is:

Ka =

Ka = [H3O+]

Ka = 4.90 x 10-5

11

Average Ka value

= 4.90 x 10-5

+ 5.25 x 10-5

+ 5.37 x 10-5

3

= 5.17 x 10-5

Table 1.4: Tabulated data for Part 2 calculation.

pH value Ka value

Trial 1 4.31 4.90 x 10-5

Trial 2 4.28 5.25 x 10-5

Trial 3 4.27 5.37 x 10-5

Average Ka value 5.17 x 10-5

According to the Ka value from Appendix A-1, when Ka is 6.4 x 10-5

the acid will be oxalic

acid. Thus, the unknown acid from this experiment is oxalic acid as the calculated average

value of Ka is close to 6.4 x 10-5

which is 5.17 x 10-5.

By taking the theoretical value of Ka in

KHP =

Percentage error,

= Experimental value – theory x 100%

Theory

= –

– = 1225.6%

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DISCUSSIONS

The first part of the experiment is the titration of the unknown acid solution with 0.1

M NaOH. The result was tabulated in a Table 1.1. Based from the table, the titrations were

done three times. The value of equivalence point for each titration was first obtained before

proceeding to the calculation of the ionization constant expression, Ka values. Ka is the

equilibrium constant for the ionization of an acid (Chang R. & Goldsby K.A., 2013). The acid

ionization constant expression of Ka:

pka = -log Ka

Ka = 10-pka

Equivalence point can be obtained by dividing the two points of pH where a sudden change

of pH occurred. From the equivalence point, the volume at equivalence point can be plotted

from the graph. Then, by dividing the volume into half, we get the pH value at half

equivalence point. The pH at half equivalence point is equal to pka. From the equations, the

value of Ka was determined. The average value of Ka from the three titrations is 3.98 x 10-6

.

Based on Appendix A; when Ka = 3.90 x 10–6

, the corresponding acid is Potassium Hydrogen

Phthalate (KHP). Thus the unknown acid is KHP because the average Ka value is 3.98 x 10-6

which is approximately 3.90 x 10–6

.

In part two of the experiment, there was no titration has been done. The pH value of

the unknown acid solution was taken and repeated three times. The Ka value was calculated

by using these expressions:

pH = -log [H3O

+]

[H3O

+] = 10

-pH

Ka =

After getting three values of Ka, the average Ka value obtained is 5.17 x 10-5

.

According to the Ka value from Appendix A-1, when Ka is 6.4 x 10-5

the acid will be oxalic

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acid. Thus, the unknown acid from this experiment is oxalic acid as the calculated average

value of Ka is close to 6.4 x 10-5

which is 5.17 x 10-5.

The difference value of Ka calculated

as compared to the first part of the experiment was due to some errors. One of the errors

might be the use of pH electrode. The pH electrode doesn’t always give accurate readings as

we can see the rapid changes of the pH values. When the pH electrode was immersed in the

solution, the tip of the electrode might have touched the surface or wall of the beaker which

would affect the pH readings. However, based on the percentage error calculated from both

parts of the experiment, the Ka value from part one of the experiment is more accurate as the

percentage error calculated is only 2.05% rather 1224.5% in part two. The error in part two is

too large. Thus, titration method is more preferable in order to determine the Ka of an

unknown acid. Some precautions should be practised in order to get an accurate value of Ka.

14

CONCLUSIONS

Based from the obtained results, it can be concluded that the determination of Ka

value of an unknown acid solution can be achieved by using titration method. From the

calculations based on the first method, the value of the Ka of the unknown acid solution is

3.98 x 10-6

which is approximately close to 3.90 x 10–6

. According to Appendix A, when

Ka=3.90 x 10–6

the corresponding acid is potassium hydrogen phthalate (KHP). Thus, the

unknown acid used in this experiment is potassium hydrogen phthalate. Compared to the

second method, the value of Ka is 5.17 x 10-5

. Theoretically when Ka is 6.4 x 10-5

, the

corresponding acid will be oxalic acid. However, the value of Ka based on the first method is

more accurate because it has a less value of percentage error which is only 2.05%. So, it is

concluded that the unknown acid used in this experiment is hydrogen phthalate with Ka value

of 3.98 x 10-6

.

RECCOMENDATIONS

There are a few recommendations that should be taken into account in order to obtain

the results more accurately. Some of the recommendations are:

Use dried standard apparatus before weighing and diluting because the standard

solution of NaOH should be a hundred percent pure and stable at room temperature.

When taking the pH of the solution, make sure that the tip of the electrode does not

comes in contact with the wall or bottom part of the beaker to avoid parallax error.

During titration, the eye must be perpendicular to the meniscus of the sodium

hydroxide solution to get accurate reading.

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REFERENCES

BogusBaw Pilarski,Roman Kaliszan, DariuszWyrzykowski. (2015). General Analytical Procedure for

Determination of. Journal of Analytical Methods in Chemistry, 1-9.

Daintith, J. (2008). A dictionary of Chemistry. Oxford University Press.

Determination of the Identity of an Unknown Weak Acid. (2004). General Chemistry Laboratory, 1-7.

Hoogendijk, R. (1999). A compact titration configuration for process analytical applications. Analytica

Chimica Acta, 211-217.

Raymond Chang, Kenneth A. Goldsby. (2013). Chemistry . New York: McGraw-Hill.