1

ACID/BASE TITRATION

EXPERIMENT 2

INTRODUCTION

In PART I of this experiment, you will determine the concentration (molarity) of a sodium

hydroxide solution (NaOH) by titrating an acid, the potassium biphthalate (HOOCC6H4COOK).

The equivalence point will be determined by the change in color of the Phenolphthalein indicator.

Knowing that one mole of NaOH neutralizes one mole of HOOCC6H4COOK, the molarity of the

sodium hydroxide solution can be determined from the volume of base added to reach the

equivalence point, and the weigh of biphthalate acid used. The accurate determination of the NaOH

solution molarity is called standardization. After PART I, the NaOH solution will be called

Standard NaOH solution.

In PART II of this experiment, you will determine the molarity of an unknown weak acid by a pH-

titration with your own Standard NaOH solution. Measuring the pH of the solution and the volume

of NaOH added as the titration proceeds, a titration curve (graph) will be constructed. The

equivalence point, the pKa and Ka of the unknown acid can then be determined from the curve.

Mastering the techniques when using the pipet, buret and analytical balance will be crucial here.

THEORY

Acid/base titration is the process of mixing measured volumes of acid and base solutions in such a

manner that you can determine when, equivalent amounts (moles) of each are present. The purpose

of titration is to determine the concentration of one solution (the base in PART I of this

experiment), while the concentration of the other (the acid) is known to a high degree of accuracy.

The reaction of an acid HA with a base B is called neutralization, and is represented by equation

(1).

HA + B A + BH+ (1)

An acid/base neutralization reaction is simply a proton (H+) transfer. The acid acts as a proton

donor, the base as a proton acceptor.

The equivalence point of a titration is the point at which equivalent amounts (number of moles) of

acid and base have been mixed. In order to determine the equivalence point, a visual indicator is

added to the solution to be titrated. When properly selected, the indicator undergoes a sharp color

change slightly after the equivalence point. So slightly in fact that we consider the color change to

happen precisely at the equivalence point.

2

In PART I of this experiment, a precisely known (weighed) amount of acid HA is placed into an

erlenmeyer flask, along with a few drops of indicator HInd, water, and a magnetic stir bar. The flask

is then set under a buret, as shown in Figure 1. The buret is filled with base B solution. While the

base is slowly added to the erlenmeyer flask, reaction (1) occurs as long as HA is present in the

flask. When the last molecules of acid HA have reacted, the next drop of base B will react with the

indicator as shown by reaction (2), and the solution will turn pink.

HInd + B Ind + BH+ (2)

Colorless pink

Reaction (2) is a neutralization reaction where base B reacts with the acid HInd. To be appropriate,

the indicator must have the following characteristics:

the acid form HInd must be of a different color than the basic form Ind,

potent dye,

HInd must be a weaker acid than the titrated acid HA.

Figure 1 – Titration setup when using a color indicator.

In order to understand better the idea behind pH-titration (PART II), we will now focus on the

chemical species present in the solution during a titration, and their impact on the pH. In the

titration of a weak acid HA, the system (in the solution) is determined by the partial dissociation of

HA

HA + H2O ⇌ H

3O+ + A (Kc = Ka) (3)

and by

A + H2O

⇌ HA + OH (4)

3

Before the titration starts, the solution contains HA and H2O, the system at that point is completely

determined by equilibrium (3).

When the titration begins, the base added (OH) reacts with the acid HA according to the

neutralization reaction,

HA + OH H2O + A (5)

which is not an equilibrium but a complete forward reaction. The neutralization will obviously

decrease the amount of HA in the solution, and increase the amount of A. As more base is added,

HA transforms into A and equilibrium (4) becomes more important.

When half the initial amount of HA has been neutralized, the amount of HA and A in the solution

are equal, and equilibria (3) and (4) are equally important. At that point, Ka = [H3O+]. This is what

we call half-volume condition.

The equivalence point is reached when enough base has been added to react all the HA initially

present, leaving only A and H2O in the solution. The system at this point is completely determined

by equilibrium (4). The pH of the solution is therefore slightly basic. Figure 2 shows how the pH

typically changes during a titration.

Past the equivalence point, the excess OH added remain in solution, shifting equilibrium (4) to the

left, increasing drastically the pH of the solution.

20 15 10 5 0 0

1

2

3

4

5

6

7

8

9

10

11

12

13

14

NaOH added (mL)

pH

Figure 2 – Titration curve of a weak acid with NaOH

4

APPARATUS AND CHEMICALS

3 125 mL erlenmeyer 1 100 mL grad. cylinder NaOH solution

1 500 mL erlenmeyer 1 10 mL grad. cylinder Unknown weak acid solution

3 50 mL beaker 1 magnetic stir plate + bar Potassium biphthalate 99.5%

1 150 mL beaker 1 pH electrode + interface Phenolphthalein solution

1 10.00 mL pipet + bulb acetone

1 25.00 mL buret distilled water

PROCEDURE

PART I – Standardization of the NaOH solution.

Preparation of the NaOH solution. Obtain 10 mL of the stock NaOH solution directly into a

500mL erlenmeyer flask. Using a 100mL graduated cylinder, add 190mL of distilled water to

the same erlenmeyer (this will dilute the stock NaOH by a factor of ~20). Cap the flask and swirl

for a few seconds. Handle that solution with great care from then on.

Standardization. Rinse your buret three times with about 5 mL of your NaOH solution. Fill

the buret to near the 0.00 mL mark with the same solution. Make sure the tip of the buret is filled

and contains no air. Install the buret and stir plate as shown in Figure 1.

Precisely weigh approximately 0.18g of potassium biphthalate directly into a clean and dry 125

erlenmeyer flask. Add about 50 mL of distilled water, 4 drops of Phenolphthalein, and a stir bar.

Put the flask on the stir plate and set a gentle stir. Titrate until a faint pink color persists for at

least 30 seconds. Perform at least three of these titrations. Record all measurements in Table 1.

PART II – Titration of the unknown weak acid.

Prepare the computer for data collection by opening the file Experiment 24a from the folder

Chemistry with Vernier. The vertical axis has pH scaled from 0 to 14 pH units. The horizontal axis

has volume scaled from 0 to 25 mL.

Calibration of the pH sensor

First Calibration Point

Choose Calibrate from the Experiment menu and then click Perform Now .

For the first calibration point, rinse the pH sensor with distilled water, then place it into a

buffer of pH=4.00.

Type “4” in the edit box as the pH value.

Swirl the sensor, wait until the displayed voltage for Input 1 stabilizes, click Keep .

Second Calibration Point

Rinse the pH sensor with distilled water, and place it into a buffer of pH=7.00.

Type “7” in the edit box as the pH value for the second calibration point.

Swirl the sensor and wait until the displayed voltage for Input 1 stabilizes. Click Keep ,

then click OK .

5

Titration of the unknown acid

1. Pipet 10.00 mL of the unknown acid solution into a 150-mL beaker. Add 50 mL of distilled

water, or just enough to soak the pH electrode properly.

2. Place the beaker on a magnetic stirrer and add a stirring bar.

Figure 3 – Titration setup for pH-titration.

3. Use a utility clamp to suspend a pH Sensor on a ring stand as shown in Figure 3. Place the pH

Sensor in the solution and adjust its position so that it is not struck by the stirring bar.

4. Fill the buret with your standard NaOH solution. Make sure to adjust at the 0.00-mL level of

the buret.

5. Before adding NaOH titrant, click Collect and monitor pH for 5-10 seconds. Check that the

Meter window shows an acidic pH value. Once the displayed pH reading has stabilized, click Keep . In the edit box, type “0” (for 0 mL added). Press the ENTER key to store the first data

pair for this titration.

6. You are now ready to begin the titration.

a) Add the next increment of NaOH titrant (enough to raise the pH about 0.15 units). Let the

pH stabilizes for 5 seconds, then click Keep . In the edit box, type the current buret reading,

to the nearest 0.01 mL. Press ENTER. You have now saved the second data pair.

b) Continue adding NaOH solution in increments of 1 mL, or, enough to raise the pH by

about 0.15 units, whatever comes first. Enter the buret reading after each increment.

c) Stop when 25mL of NaOH solution have been added.

7. When you have finished collecting data, click Stop . Dispose of the beaker contents to the

sink.

6

Warning !

Proceeding further when data collection has not been stopped,

may end up in permanent loss of all collected data.

8. To print a copy of the Table: click on the frame of the Table Window, then choose Print Window

from the File menu, enter your name(s) and click OK. Enter the number of copies and click OK.

9. To print a copy of the Graph: click on the frame of the Graph Window, then choose Print Window

from the File menu, enter your name(s) and click OK. Enter the number of copies and click OK.

10. The equivalence-point is located exactly at the inflexion point of the titration curve. To

determine its position, examine the graph of the first derivative (pH/Vol) vs Volume. Click

the pH vertical-axis label of the graph, check the box for first derivative, uncheck the pH box,

then click OK. Reset the y axis by clicking on any y axis number and then select Autoscale.

Then print the Graph Window (step 9). Indicate with a mark (use a pen), the position of the

equivalence point on the titration curve.

At the end of the lab period:

- the pH electrode must be rinsed thoroughly with distilled water and returned to the storing

jar.

- the buret must be rinsed three times with distilled water, and left upside down on the stand

with the valve open.

7

CALCULATIONS

The key relation to determine the molarity of your NaOH solution is:

MAVA = MBVB (6)

where: MB : molarity of the NaOH solution

VB : volume of NaOH used to get to the equivalence point

MA : molarity of the biphthalate solution

VA : volume of the biphthalate solution used

In fact, equation (6) is simply another way to express the meaning of the equivalence point, where

nA = nB ( nA : number of moles of biphthalate acid, nB : number of moles of NaOH)

To calculate MB in PART I, you will therefore use equation (7):

nA = MBVB (7)

Knowing the molecular weight of potassium biphthalate (MW=204.23g/mol), nA can be calculated

from the weights in Table 1.

Determine the molarity of your unknown acid solution. First calculate the number of moles of

NaOH required to titrate the sample to the equivalence point (determined from the titration curve).

Then use the volume of acid sample to calculate the molarity of the acid.

As explained earlier, the pKa of your acid is calculated from the titration half-point, where the

number of moles of base added is equal to half the number of moles of total acid. Since at this

point, half of the original HA has been converted (to a very good approximation) to A, we have

[A] [HA] or [A]

[HA]

Therefore, this half-volume condition is a special situation where

K a

[ H 3 O ] [ A ]

[ HA ] [ H 3 O ]

(9)

Taking the negative log of each side we obtain:

pKa = pH (10)

1 (8)

8

Precision

The term precision describes the reproducibility of results. It can be defined as the agreement

between multiple measurements that have been made under the same conditions. The precision of a

set of measurements is related to the deviation of the data set from the average.

Trial # Data Deviation

from average

1 0.315 0.004 = ABS(0.319 0.315)

2 0.321 0.002 = ABS(0.319 0.321)

3 0.320 0.001 = ABS(0.319 0.320)

average 0.319 0.002

So the average is 0.319 with an average deviation of 0.002.

Or, in terms of percent average deviation (precision or relative deviation)

Percent Average Deviation = average

iationaveragedev x 100% =

319.0

002.0 x 100% = 0.6%

All deviation values (average and percent) must be reported with ONE significant figure only.

9

ACID/BASE TITRATION

Experiment 2 – Data

Name

First Last ID

Demonstrator

Section

Lab Date

Table 1 – Titration of weighed amounts of potassium biphthalate

Trial Mass of biphthalate Volume of NaOH sln

Initial Final added

# (g) (mL) (mL) (mL)

1

2

3

4

5

6

Table 2 – pH-titration of the unknown acid solution

Code number of unknown acid

Initial volume reading of NaOH solution in the buret

Before handing in this Data Sheet, please attach (staple):

1. Titration curve

2. First derivative curve

3. Data Table (from computer)

Data sheet, hand in before leaving the lab

10

ACID/BASE TITRATION

Experiment 2 – Lab report

Name

First Last ID

Demonstrator

Section

Lab Date

Table 3 – NaOH molarity from titration of weighed amounts of potassium biphthalate

Trial

mA* nA VB MB Deviation from

MB average

# (g) (mol) (L) (mol/L) (mol/L)

*mA : mass of potassium

biphthalate from Table 1 Average

Percent deviation

Table 4 – Concentration and strength of the unknown acid. (Code of unknown acid: _______)

MB† VB VA MA Ka pKa

(mol/L) (L) (L) (mol/L) (mol/L)

†MB from Table 3 (average)

Lab report, hand in within 24 hours

11

QUESTIONS

1. Show all the steps and calculations involved to complete the first row of Table 3 and 4.

2. Why is it important that the indicator be a potent dye ? Explain.

Lab report, hand in within 24 hours

12

3. At the beginning of a titration to standardize the NaOH solution, Student A adjusted very

carefully the initial burette volume to 0.00mL. But he did not notice an important air bubble in

the tip of the burette. At the end of the titration, the air bubble is gone. Explain the effect of

that mistake on the calculated molarity MB. (Will the experimental MB calculated by Student A

be higher or lower than the true MB value?)

Lab report, hand in within 24 hours

13

Questions and problems

1. What indicator is used in this experiment?

2. What is the chemical formula of the acid used to standardize the NaOH solution ?

3. During standardization, what’s the color of the solution in the erlenmeyer flask at the

equivalence point ?

4. Define equivalence point.

5. Write the equation for the reaction of NaOH with the indicator. (Use Hind as the chemical

formula of the indicator).

6. You used 15.44 mL of NaOH(0.1022M) to neutralize 10.00 mL of HA solution. Knowing that

Ka(HA)=2.0105

and Kb(A)=5.010

10, answer the following questions:

a) Write down the chemical equation for the neutralisation of HA by NaOH.

b) Calculate the pH at the beginning of the titration, before adding any NaOH

Answers

1. Phenolphthalein

2. HOOCC6H4COOK

3. Pink (it’s possible that the pink color fades away 30 seconds after equivalence has been

reached)

4. The equivalence point of a titration is the point at which equivalent amounts (number of

moles) of acid and base have been mixed.

5. HInd + NaOH NaInd + H2O or HInd + OH Ind + H2O.

6a. HA + OH H2O + A

6b. MA = MB•VB / VA = 0.1022•15.44 / 10.00 = 0.1578 M

[HA] = 0.1578 M

HA + H2O A + H3O

+

0.1578 --- 0 10-7

-x -x +x +x

0.1578-x --- x x+10-7

x <<0.1578

x >>10-7

x2/0.1578 = 2.010

5 x = 1.810

3

pH = Log(1.8103

) = 2.8