EFFECT OF TEMPERATURE ON THE ACID ... - …senaychem.com/wp-content/uploads/2010/09/final-Acid-Base-Oğul... · constant for a chemical reaction known as dissociation in the context

TED ANKARA COLLEGE FOUNDATION HIGH SCHOOL

EFFECT OF TEMPERATURE ON THE ACID

DISSOCIATION CONSTANT OF ACETIC ACID

Assessment Criteria: Full Investigation (D, DCP, CE)

Session: May 2013

Candidate Name: Oğul Ersin ÜNER

Candidate Number: D1129066

Centre Name: TED Ankara College High School, Ankara, TURKEY

Date of Experiment: 10.04.2012

Instructor: Ms. Ayse Senay

OĞUL ERSİN ÜNER D1129066

2

BACKGROUND INFORMATION

What is Titration?

Titration determines the molarity of an unknown concentration of a solution. During the

process, the concentration of the solution is found by the addition of another substance which has a

known concentration. “The addition is stopped when the endpoint is reached.” 1

The colour of the

solution changes due to the presence of an indicator, a colourless solution which specifically changes

colour according to the pH of the medium.

In this titration experiment, the weak acid acetic acid, CH3COOH(aq), is titrated with the strong

base sodium hydroxide, NaOH(aq), in order to find the concentration of the acetic acid of unknown

concentration. The colour is changed by using the indicator phenolphthalein, which “is colourless in

acid solution and pink in alkaline solution”2.

What is Ka?

“An acid dissociation constant, Ka, (also known as acidity constant, or acid-ionization

constant) is a quantitative measure of the strength of an acid in solution. It is the equilibrium

constant for a chemical reaction known as dissociation in the context of acid-base reactions.”3 It can

only be written for weak acids, such as for acetic acid, which partially ionize in water. It can be

expressed as:

3

where:

Ka is the acid dissociation constant of the acid HA

[A-] denotes the concentration of the conjugate base of the acid HA

[H+] denotes the concentration of hydrogen ions.

[HA] denotes the concentration of the acid HA

What is Temperature and how does it affect the Ka of acids?

“Temperature is a physical property of matter that quantitatively expresses the common

notions of hot and cold. Objects of low temperature are cold, while various degrees of higher

1 Britannica. n.p., n.d. Web. 3 May. 2012. URL: < http://www.britannica.com////>

2 Atkins, P.W. (1978). “Physical chemistry”, Oxford University Press. Print. 3 Rossotti, F.J.C., Rossotti, H. (1961). “The Determination of Stability Constants”. McGraw–Hill. Chapter 2. Print.

http://www.britannica.com/EBchecked/topic/190913/equivalence-point


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temperatures are referred to as warm or hot. Heat spontaneously flows from bodies of a higher

temperature to bodies of lower temperature, at a rate that increases with the temperature difference and

the thermal conductivity. No heat will be exchanged between bodies of the same temperature; such

bodies are said to be in "thermal equilibrium".”4 Temperature has a very important effect on the

dissociation constant. Usually, the dissociation constant is calculated for that at 25.0oC.

The influence of temperature on the acid dissociation constant of acetic acid can be calculated

using the Van’t Hoff equation.5 “The Van’t Hoff equation states amongst others that dissociation

constants depend on temperature and the difference in the 'heat of formation' between the ion or

compound that dissociates and the ion or compound it dissociates into.”5

5

where:

K1 is the equilibrium constant at absolute temperature T1,

K2 is the equilibrium constant at absolute temperature T2,

ΔHΘ is the standard enthalpy change of the process and

R is the gas constant (8.3144621 J K−1

mol−1

)

The dissociation constant will be found by determining the molarity of acetic acid by titration

and then finding the hydrogen and acetate ion concentrations in the acetic acid solution. Acetic acid is

a weak acid, so the partial ionization of the acid and the equilibrium will be examined first. The pH

will then be measured and the acid dissociation constant will be calculated. After comparing the Ka

values of the temperatures 20.0°C, 35.0°C and 50.0°C, which are 293 K, 308 K and 323 K

respectively, the relationship will be found out.

Factors that affect the Ka of an acid:

The factors which affect the rate of reaction are listed below:

• Temperature

• Nature of Reactant

• Nature of Solvent

• Pressure

4 Thermometer and temperature. n.p., n.d. Web. 29 April. 2012. URL: <http://hyperphysics.phy-astr.gsu.edu/hbase/thermo/temper.html> 5 Prant, J.A. (1989). “Chemical Components”. Cambridge University Press. Print.


4

EFFECT OF TEMPERATURE ON THE ACID DISSOCIATION CONSTANT OF

ACETIC ACID (ETHANOIC ACID)

Research Question:

What is the effect of increasing the temperature of 30.0 mL CH3COOH(aq) solution of an

unknown concentration to 20.0°C, 35.0°C and 50.0°C, on the acid dissociation constant of

CH3COOH(aq), whose concentration is found by titrating 30.0 mL CH3COOH(aq) with 1 M NaOH(aq)

and its acid dissociation constant found by changing the temperature of the CH3COOH(aq) solution and

measuring the pH with a pH meter at that temperature when room pressure, 1067.0 hPa, type of acid,

acetic acid, volume of acetic acid used in each trial, 30.0 mL, molarity of sodium hydroxide solution

used in each trial, 1M and the number of phenolphthalein solution drops, three drops, in each trial are

kept constant?

Purpose:

The aim of this investigation is to determine the effect of increasing the temperature of 30.0

mL CH3COOH(aq) solution of an unknown concentration to 20.0°C, 35.0°C and 50.0°C, on the acid

dissociation constant of CH3COOH(aq), whose concentration is found by titrating 30.0 mL

CH3COOH(aq) with 1 M NaOH(aq) and its acid dissociation constant found by changing the temperature

of the CH3COOH(aq) solution and measuring the pH with a pH meter at that temperature when room

pressure, 1067.0 hPa, type of acid, acetic acid, volume of acetic acid used in each trial, 30.0 mL and

the number of phenolphthalein solution drops, three drops, in each trial are kept constant.

Variables:

Dependent: Acid dissociation constant of CH3COOH(aq) (acetic/ethanoic acid) at 20.0°C, 35.0°C and

50.0°C

Independent: Temperature of CH3COOH(aq), which is changed by heating the CH3COOH(aq) solution

to 20.0°C, 35.0°C and 50.0°C from 19.5°C (room temperature)

Controlled:

Room temperature, 19.5 ± 0.2°C.

Room pressure, 1067.0 ± 0.2 hPa.

Temperature of NaOH(aq) , 19.5 ± 0.2°C, which should be the temperature of the laboratory

http://www.blurtit.com/q614258.html












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Temperature of the beaker and erlenmeyer flasks, 19.5 ± 0.2°C, which should be the temperature

of the laboratory

Volume of CH3COOH(aq) , 30.0 ± 0.2 mL in each trial

Volume of phenolphthalein solution used in all trials, 3 drops in each trial

Molarity of NaOH(aq) , 1 M in each trial.

Molarity of the unknown concentration of CH3COOH(aq) in each titration.

Type of the indicator solution used in all trials, phenolphthalein

Using the same stock of phenolphthalein solution for all titrations so that the molarities of the

phenolphthalein solutions will be constant throughout all trials.

Using the same stock of NaOH(aq) for all titrations so that the uncertainties in the molarities of

NaOH(aq) will be constant throughout all trials.

Time when the titration was stopped, when the pink colour of CH3COOH(aq) is observed and

remains for 20 seconds.

Using the same stock of unknown concentration of CH3COOH(aq) for each titration so that the

uncertainties in the molarities of CH3COOH(aq) will be constant in all trials.

Using the same pH meter in each trial.

The renewal of the batteries and/or charging of the pH meter before the experiment.

The calibration of the pH meter before the experiment.

Using the same equipments (the same graduated cylinder, erlenmeyer flask, burette, thermometer,

barometer etc.) in all trials so that the uncertainties are kept constant in all trials.

MATERIALS:

450.0 mL CH3COOH(aq) solution of unknown molarity, provided by the instructor

1 M, 450.0 mL NaOH(aq) solution

45 drops of phenolphthalein solution of the same concentration

25.0 mL graduated cylinder x1

50.00 mL burette (± 0.02 mL) x1

100. mL erlenmeyer flask (± 4 mL) x1

600. mL beaker (± 20. mL) x1

Thermometer with a range of -10.0 °C to 110.0 °C (± 0.2°C) x2

Barometer with a range of 870 hPa to 1084 hPa (±.02 hPa) x1

3.0 mL dropper pipette (± 0.1 mL) x1

Funnel x1

Ring stand x2



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Test tube clamp x1

Burette clamp x1

Burette card with dimensions 3cm x 5cm, white card, painted half black. Put on the back of the

burette for reducing the parallax errors while reading the volume x1

Bunsen burner x1

Matchbox x1

pH meter x1

Chronometer x1

Wash bottle full of distilled water x1

Some paper towel

Safety goggles x1

A water bath will be prepared by adding 300 mL tap water to the 600. mL beaker (Figure 2).

SAFETY PRECAUTIONS:

Wear approved eye protection such as safety goggles in the laboratory at all times.

Sodium hydroxide is “corrosive, irritant and caustic. Do not handle the liquid with your hands. In

the event of skin contact, wash well with water. If the skin is irritated or broken, seek professional

medical treatment. ”6

Concentrated acetic acid “can be irritant and corrosive to skin, and it can sometimes be flammable

depending on its concentration”7. Although the concentration given by the instructor for this

investigation is not a great value, dilute acetic acid is also irritant. Therefore, “it should be handled

with appropriate care, since it can cause skin burns and permanent eye damage.”7

Phenolphthalein is “moderately toxic and irritant. Avoid eye and skin contact.”8

The titration of acetic acid with sodium hydroxide generates heat (exothermic). Take care in

handling the erlenmeyer flask with the mixture.

DISPOSAL INFORMATION:

All solutions in this experiment should be disposed in the proper waste containers in the fume hood as

provided by the instructor in the laboratory.

6 Sodium Hydroxide Safety Sheet. MSDS, n.d. Web. 4 May. 2012. URL: < http://www.sciencelab.com/msds.php?msdsId=9924999 > 7 Acetic Acid Safety Sheet. MSDS, n.d. Web. 6 May. 2012. URL: <https://www.sciencelab.com/msds.php?msdsId=9922769> 8 Phenolphthalein Safety Sheet. MSDS, n.d. Web. 7 May. 2012. URL: <http://www.sciencelab.com/msds.php?msdsId=9926477>


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PROCEDURE:

1. Put on your safety goggles.

2. Close all the doors and windows in order to help the room maintaining the room temperature

constant.

3. Measure the room temperature with a thermometer before each trial in order to make sure that the

temperature of the flasks and the beaker are kept constant.

4. Measure the room pressure with a barometer before each trial.

5. Construct a titration set-up with a burette, burette holder, ring stand and funnel (Figure 1).

6. Fill the burette with 50.00 mL (the volume capacity of the burette) distilled water from the wash

bottle.

7. Open the stopcock and let the distilled water run out into the sink.

8. Prepare a 600 mL beaker in the presence of constant room temperature and pressure, which are

19.5°C and 1067.0 hPa respectively.

9. Fill the burette with 1M, 30.00 mL NaOH(aq) by using a funnel and let it run out into the 600 mL

beaker so as to get rid of the impurities in the burette.

10. Pour the NaOH(aq) in the beaker to the waste container in the fume cupboard, as instructed.

11. Clean the beaker with tap water, rinse it with distilled water, and dry it out with paper towel.

12. Fill the burette with 50.00 mL, 1 M, NaOH(aq).While doing this, put the funnel on the mouth of the

burette, which will prevent the spilling of NaOH(aq).

13. Read the initial volume of NaOH(aq) on the burette and record the value. While reading the value

on the burette, use the burette paper. The black part on the paper should be placed just below the

meniscus of NaOH(aq) on the other side of the burette. Read the lowest point of the meniscus at eye

level to decrease the errors caused by the refraction of light.

14. Prepare a clean, dry 100. mL erlenmeyer flask and place 30.0 mL of unknown concentration of

CH3COOH(aq) using the 25.0 mL graduated cylinder in the flask. Use the same unknown

concentration of the acid solution in all trials by using the same stock in each trial.

15. Get 3 drops of phenolphthalein indicator with a dropper and add them in the erlenmeyer flask. Use

the same phenolphthalein solution from the same stock in all trials.

16. Construct a water bath set-up with a bunsen burner, beaker, ring stand and a test tube clamp

(Figure 2) by adding 300 mL tap water into the clean, dry 600 mL beaker.

17. Place the erlenmeyer flask with CH3COOH(aq) and phenolphthalein in the 600 mL beaker which

was filled with tap water. Measure the temperature of water with a thermometer. If required, heat

up the beaker to 20.0 °C before titration. Record this value of temperature as the temperature of

CH3COOH(aq). The erlenmeyer flask will remain in the water bath during titration.


8

18. Open the stopcock and allow the NaOH(aq) to drop into the flask and start a reaction with

CH3COOH(aq). Swirl the erlenmeyer flask during the addition of NaOH(aq) into the flask.

19. Check the temperature of the reaction mixture in between the time intervals 0.0-20.0 seconds in

the flask after adding NaOH(aq) and heat the solution if needed.

20. Close and open the stopcock until the CH3COOH(aq) solution becomes temporarily pink. At the

moment when the light pink colour is observed in the erlenmeyer flask, slow down the addition of

NaOH(aq), start the chronometer and wait for 20.0 seconds while swirling the solution. When the

time elapses and the color is still visible, close the stopcock.

21. Read the lowest point of the meniscus at eye level on the burette in order to measure the volume of

NaOH(aq) by using a burette paper to decrease the parallax errors and then record the volume.

22. Pour the mixture in the erlenmeyer flask into the waste container.

23. Clean the flask with tap water, rinse it with distilled water, and dry it out with paper towel.

24. Add some more 1 M NaOH(aq) into the burette until 30.00 mL level is read on the burette. Read the

volume on the burette by using the burette card and read the lowest point of the meniscus at eye

level in order to decrease the errors caused by refraction.

25. Calculate the molarity of CH3COOH(aq) solution using the data collected by titration.

26. Repeat steps 1-4 and 14-25 four more times in order to complete 5 trials and acquire a data group

of 15 for titration.

27. Prepare a clean, dry 100. mL erlenmeyer flask and place 30.0 mL of CH3COOH(aq) of known

molarity using the 25.0 mL graduated cylinder in the flask and determine the pH of CH3COOH(aq)

in the erlenmeyer flask at 20.0°C by using the pH meter. First, submerse the electrode of the pH

meter in the acetic acid solution, then press the measure button and record the value. After

recording the pH, calculate the hydrogen and acetate ion concentration by using the pH value.

28. Determine the acid dissociation constant of CH3COOH(aq) at 20.0°C using the acetate and

hydrogen ion concentrations found on step 27 and record the value.

29. Clean the pH meter with distilled water and dry it with paper towel.

30. Fill a clean, dry erlenmeyer flask with 30.0 mL acetic acid solution of the same known molarity.

31. Obtain a thermometer and place the flask in the water bath which is on top of the bunsen burner.

Start heating the water bath and stop heating when the thermometer reads 35.0°C.

32. Determine the pH of CH3COOH(aq) in the erlenmeyer flask at 35.0°C by using the pH meter. Then

calculate the hydrogen ion concentration by using the pH value.

33. Determine the acid dissociation constant of CH3COOH(aq) at 35.0°C using the acetate and

hydrogen ion concentrations found on step 33 and record the value.

34. Repeat steps 1-4 and 29-33 but heat the water bath to 50.0°C in step 31and determine the acid

dissociation constant of the acetic acid solution.

35. Repeat steps 1-4 and 27-34 four more times in order to have 5 trials and have a data group of 15.


9

Figure 1: The titration set-up Figure 2: The water bath set-up

in the investigation.9 in the experiment.

10

D:

A.1:C

A.2:C

A.3:C

9 Titration. n.p., n.d. Web. 24 April. 2012 URL: <http://water.me.vccs.edu/courses/env211/changes/titration.gif> 10 Molecular Weight of A Volatile Liquid. n.p., n.d. Web. 19 April. 2012 URL: <http://intro.chem.okstate.edu/HTML/P8.HTM>


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DATA COLLECTION

Table 1: The initial and final readings on the burette containing the aqueous sodium hydroxide solution, the volume of the aqueous sodium hydroxide solution

obtained from the difference of the initial and final reading on the burette, the molarity of aqueous sodium hydroxide, volume of aqueous acetic acid solution

used, the temperature of the water bath, the room temperature and room pressure in each trial during titration.

TITRATION OF CH3COOH(aq) WITH NaOH(aq)

TRIAL

INITIAL READING

ON BURETTE

CONTAINING

NaOH(aq) SOLUTION

( 0.02 mL)

FINAL READING

ON BURETTE

CONTAINING

NaOH(aq) SOLUTION

( 0.02mL)

VOLUME

OF NaOH(aq)

SOLUTION

USED

( 0.04 mL)

VOLUME OF

CH3COOH(aq)

SOLUTION

USED

( 0.2 mL)

MOLARITY

OF NaOH(aq)

SOLUTION

(M)

TEMPERATURE

OF THE WATER

BATH/TITRATION

MIXTURE

( 0.20C)

NUMBER OF

PHENOLPHTHALEIN

SOLUTION DROPS

ROOM

TEMPERATURE

( 0.20C)

ROOM

PRESSURE

( 0.2 hPa)

1 30.00 1.10 28.90 30.0 1 20.2 3 19.5 1067.0

2 30.00 0.39 29.61 30.0 1 20.3 3 19.5 1067.0

3 30.00 0.11 29.89 30.0 1 20.1 3 19.5 1067.0

4 30.00 0.25 29.75 30.0 1 20.3 3 19.5 1067.0

5 30.00 0.34 29.66 30.0 1 20.0 3 19.5 1067.0


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DETERMINATION OF THE ACID DISSOCIATION CONSTANT OF CH3COOH(aq)

FOR 20.0°C ACETIC ACID SOLUTION

TRIAL

INITIAL TEMPERATURE

OF CH3COOH(aq)

SOLUTION

( 0.2 °C)

FINAL TEMPERATURE

OF CH3COOH(aq)

SOLUTION (JUST

BEFORE pH READING)

( 0.2 °C)

pH OF

CH3COOH(aq)

SOLUTION

(±0.01)

VOLUME OF

CH3COOH(aq)

SOLUTION USED

( 0.2 mL)

TEMPERATURE OF

THE WATER BATH

( 0.20C)

1 19.5 20.0 2.52 30.0 20.2

2 19.4 20.0 2.52 30.0 20.3

3 19.7 20.0 2.51 30.0 20.0

4 19.5 20.0 2.51 30.0 20.1

5 19.6 20.0 2.51 30.0 20.0


TRIAL

INITIAL TEMPERATURE

OF CH3COOH(aq)

SOLUTION

( 0.2 °C)

FINAL TEMPERATURE

OF CH3COOH(aq)

SOLUTION(JUST

BEFORE pH READING)

( 0.2 °C)

pH OF

CH3COOH(aq)

SOLUTION

(±0.01)

VOLUME OF

CH3COOH(aq)

SOLUTION USED

( 0.2 mL)

TEMPERATURE OF

THE WATER BATH

( 0.20C)

1 19.2 35.0 2.47 30.0 35.0

2 19.3 35.0 2.39 30.0 35.1

3 19.1 35.0 2.46 30.0 35.0

4 19.3 35.0 2.45 30.0 35.3

5 19.0 35.0 2.47 30.0 35.0


TRIAL

INITIAL TEMPERATURE

OF CH3COOH(aq)

SOLUTION

( 0.2 °C)

FINAL TEMPERATURE

OF CH3COOH(aq)

SOLUTION (JUST

BEFORE pH READING)

( 0.2 °C)

pH OF

CH3COOH(aq)

SOLUTION

(±0.01)

VOLUME OF

CH3COOH(aq)

SOLUTION USED

( 0.2 mL)

TEMPERATURE OF

THE WATER BATH

( 0.20C)

1 33.1 50.0 2.57 30.0 50.2

2 33.3 50.0 2.56 30.0 50.1

3 33.0 50.0 2.58 30.0 50.0

4 33.2 50.0 2.57 30.0 50.2

5 33.0 50.0 2.59 30.0 50.0

Table 2: The initial and final temperatures of the acetic acid solution, pH of the acetic acid, volume of

acetic acid solution used and the temperature of the water bath for 20.0, 35.0 and 50.0°C.


12

Note: The solutions were prepared by the lab technician and therefore, the uncertainties for the

concentrations of the solutions could not be provided.

Qualitative Data:

CH3COOH(aq), phenolphthalein and NaOH(aq) were colourless before titration.

The indicator phenolphthalein did not change the colour of the mixture immediately when it was

added to the flask that contained acetic acid solution.

When NaOH(aq) started to drop into the flask with CH3COOH(aq) and phenolphthalein, the colour of

the solution started to turn into pink temporarily.

While swirling the flask, the temporary light pink colour observed faded away and the mixture of

CH3COOH(aq) and phenolphthalein became colourless again.

The light pink colour observed before did not fade away at the endpoint of titration for about 20.0

seconds.

The temperature of the water bath increased when NaOH(aq) was added to the flask of

CH3COOH(aq) and one could feel the heat liberated from the erlenmeyer flask by touching it, which

indicated that the reaction was exothermic.

Acetic acid had a strong vinegar odor that could be smelled from a distance.

Apparently, sodium hydroxide solution was odorless. A smell could not be detected from afar and

the solution could not be smelled closely due to safety reasons.

DATA PROCESSING

1. FINDING THE MOLARITY OF ACETIC ACID (ETHANOIC ACID) SOLUTION

USED

During titration, the end point, which is marked by the indicator by the change of color, is the

point where the stoichiometric coefficients of the titrant (sodium hydroxide) and the analyte (acetic

acid) are equal; “the amount of titrant is sufficient to fully neutralize or react with the analyte.”11

So

the number of moles of the acid and base are equal.

where n is the number of moles.

11 Whitney, W.D., Smith, B.E. (1911). "Titrimetry". The Century Dictionary and Cyclopedia. The Century co. pg. 6504. Print.


13

In order to find the number of moles of acid and base at that point, the molarities should be multiplied

by the volume of the acid or base used.

where n is the number of moles, M is the molarity in molars and V is the volume in liters.

Balanced equations aid in observing stoichiometric coefficients and must be used to find the molarity

of the acetic acid solution. One should assume that the indicator changes color exactly at the

equivalence point, due to the fact that “some indicators change near the point.”12

The rightmost significant figures are shown in bold in every calculation.

Unless otherwise stated, all values given in tables are rounded according to their uncertainties.

The unrounded, rounded and average values are given in 2 Tables: At the end of the titration

process and at the end of the determination of acid dissociation constants in the results table.

Example for the First Trial of Titration:

THE TITRATION OF ACETIC ACID WITH SODIUM HYDROXIDE

X mol Y mol

Since there is 1 mol-1 mol stoichiometric ratio between acetic acid and sodium hydroxide in the

reaction equation above, the number of moles of sodium hydroxide and acetic acid are equal at the end

point.

12 Acetic Acid Titrations and Indicators. n.p., n.d. Web. 14 April. 2012. URL: <www. scifun.chem.wisc.edu/.../pdf/aceticacid.pdf>


14

(The molarity of NaOH(aq) has one significant figure, so the final answer also has only one.)

Uncertainty of Molarity of Acetic Acid in the First Trial of Titration:

Note that the percentage uncertainty of the molarity of cannot be included in the

uncertainty calculations, as the 1 M was prepared by the lab instructor prior to the

experiment.

(Percentage uncertainty of ) + (Percentage uncertainty of

)

Consequently,

By conducting these steps for the other trials, one can find the molarity of acetic acid and the

percentage uncertainty as:

Trial 2: 0.9787 M 1 M 0.4351% +/- 0.43%

Comment [a1]: Should be V

Comment [a2]: This is not clear

Comment [a3]: % uncertainties smaller than 2 should be written in 2 sig figs.

Comment [a4]: Unroudned value should be writtent first.


15

Trial 3: 0.9878 M 1 M 0.4338% 0.43%

Trial 4: 0.9856 M 1 M 0.4344% 0.43%

Trial 5: 0.9826 M 1 M 0.4349% 0.43%

TITRATION

TRIAL

UNROUNDED MOLARITY OF ACETIC ACID

SOLUTION USED

(M)

ROUNDED MOLARITY OF ACETIC

ACID SOLUTION USED

(M)

1 0.9633 ± 0.4027 % 1 ± 0.40 %

2 0.9787 ± 0.4351 % 1 ± 0.43 %

3 0.9878 ± 0.4338 % 1 ± 0.43 %

4 0.9856 ± 0.4344 % 1 ± 0.43 %

5 0.9826 ± 0.4349 % 1 ± 0.43 %

AVERAGE 0.9796 ± 0.42818 % 1 ± 0.43 %

Table 3: The unrounded and rounded molarities of acetic acid solutions found by the titration of acetic

acid with sodium hydroxide solution in all five trials with their percentage uncertainties and average.

2. FINDING THE CONCENTRATIONS OF ACETATE AND HYDROGEN IONS AT

20.0°C, 35.0°C and 50.0°C

The pH of a solution is an excellent way to find the hydrogen ion concentration in a solution.

By using a pH meter, one can find the pH of the acetic acid solution at 20.0°C.13

Since acetic acid is a

weak acid and partially ionizes into its ions, the equilibrium must be examined in order to find the

molarities of both of the ion concentrations. As the molarity of acetic acid is known, the reaction can

be established and the concentration of the acetate ions can be found at 20.0°C. The same can be

applied for 35.0°C acetic acid solution and 50.0°C acetic acid solution.

The pH of the medium and the hydrogen ion concentration can be expressed in an equation as:

or where is the hydrogen ion concentration in the solution.

Example for the First Trial of Determining the Acetate and Hydrogen Ion Concentrations at

20.0°C:

THE PARTIAL IONIZATION OF ACETIC ACID

13 Beginner Chemistry: pH and Appliances. n.p., n.d. Web. 27 April. 2012. URL: <www.elmhurst.edu/~chm/.../184ph.html>


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In: 1 M

Rxn: - Y M + Y M + Y M Ionized part : Y M

Eq: (1-Y) M Y M M

(The ionized part can be neglected due to the fact that it is so small that it would not have a major

effect on the calculations in determining the Ka of acetic acid.)

By conducting this method for the other trials, one can find the molarity of the acetate ions as:

Trial 2: M

Trial 3: M

Trial 4: M

Trial 5: M

ACETIC ACID SOLUTION AT 20.0°C

TRIAL pH OF ACETIC ACID SOLUTION

(± 0.01)

HYDROGEN ION CONCENTRATION IN

ACETIC ACID (M)

1 2.52

2 2.52

3 2.51

4 2.51

5 2.51

Table 4: The pH values of the acetic acid solutions found by a pH meter and the hydrogen ion

concentrations that are calculated from the pH values at 20.0°C in all five trials respectively.


35.0°C:


17


In: 1 M


Eq: (1-Y) M Y M M




Trial 2: M

Trial 3: M

Trial 4: M

Trial 5: M



(± 0.01)


ACETIC ACID (M)

1 2.47

2 2.39

3 2.46

4 2.45

5 2.47

Table 5: The pH values of the acetic acid solutions found by a vernier and the hydrogen ion

concentrations that are calculated from the pH values at 35.0°C in all five trials respectively along

with their uncertainties.


50.0°C:


18


In: 1 M


Eq: (1-Y) M Y M M




Trial 2: M

Trial 3: M

Trial 4: M

Trial 5: M



(± 0.01)


ACETIC ACID (M)

1 2.57

2 2.56

3 2.58

4 2.57

5 2.59

Table 6: The pH values of the acetic acid solutions found by a vernier and the hydrogen ion

concentrations that are calculated from the pH values at 50.0°C in all five trials respectively along

with their uncertainties.


19


TRIAL CONCENTRATIONS OF ACETATE AND HYDROGEN IONS IN ACETIC ACID SOLUTION

(M)

1

2

3

4

5


TRIAL

CONCENTRATIONS OF ACETATE AND HYDROGEN IONS IN ACETIC ACID SOLUTION

(M)

1

2

3

4

5


TRIAL

CONCENTRATIONS OF ACETATE AND HYDROGEN IONS IN ACETIC ACID SOLUTION

(M)

1

2

3

4

5

Table 7: The concentrations of acetate and hydrogen ions obtained from the pH of the solutions at

20.0°C, 35.0°C and 50.0°C respectively for all five trials.

3. FINDING THE ACID DISSOCIATION CONSTANT OF ACETIC ACID AT 20.0°C,

35.0°C and 50.0°C

Since the concentrations of acetate and hydrogen ions are known, the acid dissociation

constant at 20.0°C, 35.0°C and 50.0°C can be found. The constant for acetic acid can be expressed as:

where:

is the acid dissociation constant


20

is the concentration of the acetate ions in the acetic acid solution

is the concentration of the hydrogen ions in the acetic acid solution

is the concentration of acetic acid.

Note that the average value and uncertainty of the molarity of will be used in the

calculations for determining the acid dissociation constants and their uncertainties.

Example for the First Trial of Determining the Acid Dissociation Constant at 20.0°C:

Average M

M

M

M

Uncertainty of the Acid Dissociation Constant of Acetic Acid for the First Trial:

(Percentage uncertainty of average ) + (Percentage uncertainty of pH of )

Consequently,

Comment [a5]: Unit?

Comment [a6]: The unrounded sum should first have been shown


21

By conducting these steps for the other trials, one can find the acid dissociation constants and its

percentage uncertainties as:

Trial 2:

Trial 3:

Trial 4:

Trial 5:


M

M


(Percentage uncertainty of ) + (Percentage uncertainty of pH of )


22

Consequently,



Trial 2: M

Trial 3: M

Trial 4: M

Trial 5: M


M

M


(Percentage uncertainty of ) + (Percentage uncertainty of pH of )


23

Consequently,



Trial 2: M

Trial 3: M

Trial 4: M

Trial 5: M


24


TRIAL ACID DISSOCIATION CONSTANT OF ACETIC ACID

(M)

1

2

3

4

5


TRIAL

ACID DISSOCIATION CONSTANT OF ACETIC ACID

(M)

1

2

3

4

5


TRIAL

ACID DISSOCIATION CONSTANT OF ACETIC ACID

(M)

1

2

3

4

5

Table 8: The acid dissociation constants found in each trial at 20.0°C, 35.0°C and 50.0°C with their

percentage uncertainties.

Determining the Average Value of the Acid Dissociation Constants:

Note that unrounded values were used in this step so as to prevent imprecise calculations.


25

Average Value of the Acid Dissociation Constant at 20.0°C:

M

Average Percentage Uncertainty:


M



26


M


TEMPERATURE

OF ACETIC ACID AVERAGE ACID DISSOCIATION CONSTANTS OF ACETIC ACID (M)

20.0°C

35.0°C

50.0°C

Table 9: The average acid dissociation constants of the acetic acid solution found in different

temperatures in M with percentage uncertainties.

4. FINDING THE THEORETICAL ACID DISSOCIATION CONSTANTS OF

ACETIC ACID AT 20.0°C, 35.0°C and 50.0°C

The theoretical value of the acid dissociation constant of acetic acid at 20.0°C, 35°C and 50°C can

be found using Van’t Hoff’s equation.5

Note that the acid dissociation constants found by the Van’t Hoff’s equation will be the literature

values discussed in the Conclusion and Evaluation section.


27

where:

KT2 is the equilibrium constants at T2, at 20.0oC (293K), 35.0°C (308 K) and 50.0°C (323 K).

KT1 is the equilibrium constant at T1, at 25.0oC (298K), which is 1.754x10

-5 M.

14

is the standard enthalpy change of the dissociation of acetic acid, which is - 385 J/mol. 15

R is the universal gas constant, which is 8.314 J mol-1

K-1

.16

Since the acid dissociation constant at 25.0°C is known, one can find the acid dissociation constant

at 20.0°C, 35.0°C and 50.0°C.

Determining the Theoretical Acid Dissociation Constant at 20.0°C:

M

Percentage Uncertainty of the Theoretical Acid Dissociation Constant at 20.0°C:


14 Housecroft, C. E.; Sharpe, A. G. (2008). “Inorganic Chemistry” (3rd ed.). Chapter 6: Acids, Bases and Ions. Print. 15 Green, J., & Damji. (2008) “S. Chemistry”. Melton: IBID. Print. 16 Jensen, William B. (2003). "The Universal Gas Constant R". J. Chem. Educ. 80 (7): 731. Print.


28

M



M


°C

Comment [a7]: It is not clear how this one has 1 sig fig


28

Table 10: Results of the calculations, concentration of acetic acid values obtained from titration with their percentage uncertainties and the acid dissociation

constants of acetic acid when the acetic acid solution is at 20.0°C, 35.0°C and 50.0°C, with their percentage uncertainties.

TRIAL

TITRATION 20.0°C CH3COOH(aq) 35.0°C CH3COOH(aq) 50.0°C CH3COOH(aq)

Unrounded

Molarity

of

Acetic Acid

(M)

Rounded

Molarity

of

Acetic Acid

(M)

Uncertainty

of the

Molarity of

Acetic Acid

(± %)

(M)

Unrounded

Ka of

Acetic Acid

(M)

Rounded

Ka of

Acetic Acid

(M)

Uncertainty

of the

Ka of

Acetic Acid

(± %)

(M)

Unrounded

Ka of

Acetic Acid

(M)

Rounded

Ka of

Acetic Acid

(M)

Uncertainty

of the

Ka of

Acetic Acid

(± %)

(M)

Unrounded

Ka of

Acetic Acid

(M)

Rounded

Ka of

Acetic Acid

(M)

Uncertainty

of the

Ka of

Acetic Acid

(± %)

(M)

1 0.9633 1 0.40

2 0.9787 1 0.43

3 0.9878 1 0.43

4 0.9856 1 0.43

5 0.9826 1 0.43

AVERAGE 0.9796 1 0.43


29

ERROR PROPAGATION

1. THE THEORETICAL ACETIC ACID CONCENTRATION:

The theoretical value of the acetic acid concentration was provided by the instructor at the end of

the experiment.

Theoretical Value of the Concentration of CH3COOH(aq): 1 M

2. THE THEORETICAL ACID DISSOCIATION CONSTANTS OF ACETIC ACID:

If one looks at the percentage error in this experiment, one will attain the accuracy of the

experiment. Since there are 3 separate dissociation constants, 3 error values will be found.

Theoretical Acid Dissociation Constant of CH3COOH(aq) at 20.0°C: M




30

For 20.0°C ACETIC ACID SOLUTION:



The following calculations involve the determination of the precision of the data group. The

values found here will be discussed explicitly in the Conclusion and Evaluation section.

In order to determine precision, one must add and subtract the absolute uncertainty of the average

datum for each temperature value from the individual data acquired from every average value in

Table 10. If all of the trials are in that range, then it is precise.17

The data acquired from Table 10 are rounded to 3 significant figures in order to closely examine

the precision. The calculations were done using the unrounded values, though.

17 Accuracy and Precision. n.p., n.d. Web. 17 April. 2012. URL: <http://astro.physics.uiowa.edu/ITU/glossary/percent-error-formula/>


31

Average value of Ka for 20.0°C: M

Rounded Ka acquired from Table 10: , , , ,

M



M



M

In order to thoroughly examine the results, all of the results from the tables will be given as

unrounded values in the Conclusion and Evaluation section.

For the discussion of accuracy and precision, the ranges bound by the uncertainties of the average

experimental and theoretical values will be given as unrounded values in the Conclusion and

Evaluation section, as the ranges are indistinguishable between different temperatures of acetic

acid when rounded.

DCP:

A.1:C

A.2:C

A.3:C


32

CONCLUSION AND EVALUATION

The purpose of this experiment was to determine the effect of increasing the temperature of

30.0 mL CH3COOH(aq) solution of an unknown concentration to 20.0, 35.0 and 50.0°C, on the acid

dissociation constant of CH3COOH(aq). The concentration of acetic acid was found by titrating 30.0

mL CH3COOH(aq) with 1 M NaOH(aq) and then the acid dissociation constant of acetic acid at different

temperatures were found by changing the temperature of the CH3COOH(aq) solution and measuring the

pH with a pH meter at that temperature. The room pressure, 1067.0 hPa, type of acid, acetic acid,

volume of acetic acid used in each trial, 30.0 mL and the number of phenolphthalein solution drops,

three drops, were kept constant for ensuring that only one agent was responsible for the change in the

dissociation constant.

It was found that the temperature of the acetic acid solution had an effect on the acid

dissociation constant of acetic acid, and the research question, “What is the effect of increasing the

temperature of 30.0 mL CH3COOH(aq) solution of an unknown concentration to 20.0, 35.0 and 50.0°C,

on the acid dissociation constant of CH3COOH(aq), whose concentration is found by titrating 30.0 mL

CH3COOH(aq) with 30.0 mL, 1 M NaOH(aq) and its acid dissociation constant found by changing the

temperature of the CH3COOH(aq) solution and measuring the pH with a pH meter at that temperature

when room pressure, 1067.0 hPa, type of acid, acetic acid, volume of acetic acid used in each trial,

30.0 mL, volume of sodium hydroxide used in each trial, 30.0 mL and the number of phenolphthalein

solution drops, three drops, are kept constant?” was answered by the data in Table 10. As it was

observed from all of the trials for 35.0°C CH3COOH(aq) and 50.0°C CH3COOH(aq) and from the

average values of the acid dissociation constants at 35.0°C and 50.0°C in Table 9, the value of the

dissociation constant decreased as the temperature of the acid was increased with the exception in the

trials conducted at 20.0°C CH3COOH(aq).

When the data in Table 10 is scrutinized, some fluctuating patterns in trials are detected. When

determining the concentration via titration, the data followed an increasing, then a decreasing pattern:

There was an increase from 0.9633 M to 0.9878 M in the first three trials and then a slight decrease

from 0.9878 M to 0.9826 M in the last two trials conducted. As for the acid dissociation constants, the

data for the Ka at 20.0°C decreased and then became constant on trials 3 and 4, and then decreased

once more. Two identical values were also seen at trials 1 and 2 of the data group for the Ka at 35.0°C,

but then an increase was recorded. Another two identical values were also observed on trials 1 and 4 in

the data group for the the Ka at 50.0°C, but a decrease in between and an increase were recorded.

When perceived individually, there was always an increasing or a decreasing pattern in between the













33

trials in titration and in the trials involving the determination of the acid dissociation constant at

different temperatures, and the constant trials indicated reliability.

The data in Table 10 suggests average values that are crucial in determining the errors in the

experiment. From the data, it was found that the average experimental value of the concentration of

CH3COOH(aq) was M. The value of the initial

concentration of acetic acid was compared with the theoretical value of the concentration, which was 1

M according to the instructor, and the percentage error was calculated to be 0.02%.

In addition, the values of the acid dissociation constants were in

average for 20.0°C CH3COOH(aq) , in average for 35.0°C CH3COOH(aq) and

in average for 50.0°C CH3COOH(aq) in the experiment. When they were

compared to their theoretical (literature) values found by Van’t Hoff’s equation, which were

± 1.0% M for 20.0°C CH3COOH(aq), ± 0.57% M for 35.0°C

CH3COOH(aq) and ± 0.40% M for 50.0°C CH3COOH(aq), it was calculated that the

percentage errors of the acid dissociation constants of acetic acid were (50%),

% (20%) and % (60%) for 20.0°C CH3COOH(aq), 35.0°C CH3COOH(aq) and

50.0°C CH3COOH(aq) respectively.

Although some of the trials did not depict the apprehended pattern, an increase in the acid

dissociation constant was expected since the reaction of the dissociation of acetic acid is exothermic,

which means heat is on the side of the products. So when heat from the surroundings entered the

equilibrium, the products of the equilibrium were in excess and the equilibrium had to balance the

stress made. According to Le Chatelier's principle, the reaction had to shift towards the reactants when

the acid solution was heated in order to establish the equilibrium again which explained the increase in

the concentration of acetic acid and thus, the decrease in the acid dissociation constant. However, it

was seen that there was a slight change from the expected pattern in between 20.0°C CH3COOH(aq) and

35.0°C CH3COOH(aq), at which the acid dissociation constant increased as temperature was increased

by 15.0°C. When the data obtained from the temperatures 35.0°C CH3COOH(aq) and 50.0°C

CH3COOH(aq) were compared, however, an increase in the acid dissociation constant was observed,

which was expected.

For example, in Millikan’s oil drop experiment, Millikan used 58 oil drops in order to

calculate the charge of the electron.15

Since there are five trials in this investigation and the conclusion

is based on only five trials, more trials can be conducted to accurately investigate the relationship

between the acid dissociation constant of acetic acid and the temperature of acetic acid.

When the percentage errors and the average percentage uncertainties are examined for both the

concentration and the acid dissociation values, the small gap between the percentage uncertainty of the

Comment [a8]: ?

Comment [a9]: As the temperature increases?

Comment [a10]: You did not determine the acid concentration at different temperature values though.


34

average values of the acetic acid concentration acquired, which is 0.43%, and the percentage error

obtained for the trials 1-5 in titration, which is 0.02%, the difference is slight and may be related to

possible systematic errors that occurred during the investigation. However, the large gap between the

percentage uncertainty of the average values of the acid dissociation constants acquired, which are

0.92% for 20.0°C CH3COOH(aq), 0.84% for 35.0°C CH3COOH(aq) and 0.82% for 50.0°C

CH3COOH(aq), and the percentage error values obtained for the trials 1-5 at each temperature, which

are 50% for 20.0°C CH3COOH(aq), 20% for 35.0°C CH3COOH(aq) and 60% for 50.0°C CH3COOH(aq),

indicate that the random errors cannot be the only reason for this deviation from the theoretical value

and there should also be systematic errors in the experiment.

The percentage errors infer accuracy and precision, two crucial terms in statistical analysis.

Accuracy is related to the difference between the theoretical values and the average experimental

values found for each temperature. The concentration values of CH3COOH(aq) acquired from the

titration part of the experiment in each trial are close to the theoretical value of 1 M, which indicates

that the titration is accurate. All of the trials are slightly smaller than the theoretical value obtained, the

smallest datum collected being 0.9633 M in trial 1 and the largest being 0.9878 M in trial 3. The

percentage error of 0.02% also proves the assertion that the data collected in the titration process is

accurate.

For the accuracy of the acid dissociation constants of CH3COOH(aq), however, the data

obtained in each trial for 20.0°C CH3COOH(aq) is not close to the theoretical value, denoting

inaccuracy due to systematic errors. All of the trials for 20.0°C CH3COOH(aq) are significantly smaller

than the theoretical value obtained; the smallest datum collected was M in trials 1

and 5 and the largest was M in trials 2, 3 and 4. The calculations show that for

20.0°C CH3COOH(aq), all of the trials are out of the range of the uncertainty, 1.0%, of the theoretical

value, which supports the claim made. The range bound by this uncertainty was found to be

M.

The data acquired for 50.0°C CH3COOH(aq) is not close to the theoretical value either,

denoting inaccuracy. All of the trials for 50.0°C CH3COOH(aq) are significantly smaller than the

theoretical value obtained; the smallest datum collected was M in trial 2 and the

largest was M in trial 5. All of the trials are also out of the range of the uncertainty,

0.40%, of the theoretical value, which also advocates the claim made. The range bound by this

uncertainty was found to be M. The percentage errors that are

50% and 60% for 20.0°C CH3COOH(aq) and 50.0°C CH3COOH(aq) respectively support these

assertions, as well.


35

The data collected for 35.0°C CH3COOH(aq), partly different than the other temperature values,

connotes a little more accuracy compared to the other temperature values. However, all of the trials for

35.0°C CH3COOH(aq) are smaller than the theoretical value obtained like the other temperatures

recorded; the smallest datum collected was in trials 1 and 2 and the largest was

in trial 5. As seen, the value in the largest trial is close to the theoretical value than

the value in other trials. The last trial is therefore the trial responsible for the smaller percentage error.

All of the trials are out of the range of the uncertainty, 0.50%, of the theoretical value, which supports

the fact that the data is highly inaccurate. The range bound by this uncertainty was found to be

M. An error of 20% also advocates the claim that the data is

inaccurate, but the values are observed to be closer to the theoretical value than the values at 20.0°C

CH3COOH(aq)and 50.0°C CH3COOH(aq) .

Precision is related to the difference between the average values and the data found in each

trial for each temperature. The concentration of CH3COOH(aq) acquired from titration in each trial are

close to each other, which suggests that the measurements are reliable. The concentration values of

CH3COOH(aq) acquired from titration in most trials are within the range of the uncertainty, 4.3%, of

the average value. The range bound by this average was calculated to be 0.9754-0.9838 M, so all the

trials lie in this interval except trial 1.

For the precision of the acid dissociation constants, the calculations from the end of the Data

Collection and Processing section show that for 20.0°C CH3COOH(aq), all of the trials are located in

the range of the uncertainty, 0.92%, of the average value. The range bound by this average was found

to be M.

For 35.0°C CH3COOH(aq), however, it was found that none of the trials except trial 4 are

located in the range of the uncertainty, 0.84%, of the average value. The range bound by this average

was found to be M. This may be due to the large value

obtained in trial 5, which increased the average value and can be explained with random errors, though

there are random errors in every trial. All the other data collected in trials 1-4 are close to each other.

For 50.0°C CH3COOH(aq), none of the trials are located in the range of the uncertainty, 0.82%,

of the average value, showing that the results are nor precise for this temperature. The range bound by

this average was found to be M. This may be due to the small

value obtained in trial 2 and the large value obtained in trial 5, which altered the average value and can

be explained with random errors, even though there are random errors in every trial. All the other data

collected in trials 2-4 are close to each other, though.


36

The low accuracy and precision in the determination of the Ka, along with the large difference

between the average percentage uncertainties and the percentage errors for the Ka of acetic acid found

at 20.0°C, 35.0°C and 50.0°C can be explained with several error sources in the experiment.

Sodium hydroxide used in the titration might have reacted with the carbon dioxide present in the

air. This can result in the formation of aqueous sodium carbonate, which can interfere with the pH,

as the carbonate ion is a base. It can also result in a decrease in the concentration of sodium

hydroxide and also a decrease in the determined concentration of the acetic acid solution. This

might be a reason for the slightly lower acetic acid concentration values during titration.

The uncertainty of the molarity of sodium hydroxide was not taken into consideration during

calculations, so it could have affected the precision of the experiment..

It did not enter into my attention that the pH meter had to be calibrated until after completing all

the trials in the experiment. This was obviously my mistake and I am aware of this mistake now.

Not calibrating the pH meter before the investigation might be the reason for why all the Ka values

are smaller than the theoretical values and may also be the reason for the large difference between

the average percentage uncertainties of the acid dissociation constants and the percentage errors.

The increase in the temperature resulted in the evaporation of the acetic acid solution, which may

have caused an increase in the molarity of the acetic acid solution during the experiment.

However, it was taken as 1 M in the calculations of the experimental values of the dissociation

constants.

The exact same temperature could not be obtained in each trial since it was not an easy task to

keep the water bath temperature constant during titration. Except for trial 5 in titration, trial 5 in

the determination of the acid dissociation constant at 20.0°C, trials 3 and 5 in the determination of

the acid dissociation constant at 35.0°C and for trials 1 and 4 in the determination of the acid

dissociation constant at 50.0°C, all the values of temperature were higher than 20.0°C in titration,

also higher than 20.0°C, 35.0°C and 50.0°C for the steps for finding the acid dissociation

constants. As the dissociation of acetic acid is exothermic, the reaction shifts towards the reactants

and cause a decrease in the value of the acid dissociation constant with an increase in temperature.

Thus, it might have caused a decrease in the experimental values of Ka of acetic acid in the

determination of the acid dissociation constants in each trial.

Although the temperature of the acetic acid solution was increased from 20.0°C to 35.0°C and

then from 35.0°C to 50.0°C, a significant change in the theoretical value of the acid dissociation

constant could not be observed. The temperature difference between 20.0°C and 35.0°C and

between 35.0°C and 50.0°C might not have that big effect on the acid dissociation constant since

the acid dissociation constants are calculated only for weak acids and are very small values.

Therefore, in order to have a significant difference in the value of the acid dissociation constant,

the difference between the values of temperature should be greater next time.

Comment [a11]: Should not it be a result of incerase in the conc of acid?


37

There are a variety of ways in which the investigation can be improved by reducing the error

sources.

Random errors could be reduced by conducting more trials. More trials are definitely needed for

this experiment to have a better picture of the effect of temperature on the dissociation constant.

Sodium hydroxide is “industrially produced as a 50% solution by variations of the electrolytic

chloralkali process, and sodium carbonate is insoluble in 50% NaOH(aq) solution.”18

Carbonate-

free solutions can be obtained by diluting 50% NaOH(aq) solution, so a 2 M sodium hydroxide

solution can be prepared and diluted to 1 M in order to get rid of the impurities. This can result in

more accurate measurements. An alternate solution can be to prepare the sodium hydroxide

solution by the investigator just before titration, when there is no time limitation on conducting

this experiment, rather than leaving it to the lab technician. This way can enable the investigator to

take the uncertainties of the mass of the sodium hydroxide solid and the volume of water used in

the preparation of sodium hydroxide solution into account and can provide the uncertainty for the

molarity of the sodium hydroxide solution. This can also ensure that the contact of the NaOH

solution with the air is minimized and the results of the experiment are more precise since the

uncertainty of the molarities of the solutions used will also be included in the calculations.

The pH meter should be calibrated before the experiment. This is essential since the glass

electrode does not give a reproducible electromotive force over longer periods of time. Calibration

should be performed with at least two standard buffer solutions that span the range of pH values to

be measured. Buffers at pH 4 and pH 7 should be used to calibrate the pH meter in this

investigation. The procedure for calibrating the pH meter should be followed.19

1. Select two pH buffers that bracket the expected sample pH. The first buffer should be pH 7.00 (zero

point adjustment) and the second buffer should be near the expected sample pH (pH 4).

2. Before starting calibration, be sure the sensor and the buffer solution are at the same temperature. If not,

allow time for temperature equilibration.

3. Pour the necessary amount of buffer solutions into separate glass beakers. Buffer solutions will remain

stable in a glass beaker for an hour.

4. Close the buffer containers promptly to avoid carbon dioxide absorption with parafilm. Discard the used

buffer.

5. Place the electrode into the first buffer. When the reading is stable, set the pH meter to the pH value of

the first buffer at the measured temperature.

6. Place the electrode into the second buffer. When the reading is stable, set the pH meter to the pH value

of the second buffer at the measured temperature.

18 Kurt, C, Bitter, J. (2005) "Sodium Hydroxide". Ullmann's Encyclopedia of Industrial Chemistry, Weinheim: Wiley-VCH. Print. 19 pH Meter Calibration. n.p., n.d. Web. 20 April. 2012. URL: <http://www.all-about-ph.com/ph-meter-calibration.html>

Comment [a12]: How so?The diluting will improve which error source?

http://en.wikipedia.org/wiki/Ullmann%27s_Encyclopedia_of_Industrial_Chemistry


38

The rate of evaporation is directly proportional to the temperature. Therefore, instead of increasing

the temperature of the solution to 35.0°C and 50.0°C, the temperature of the acetic acid should be

decreased by placing some ice into the water bath in order to reduce the rate of evaporation and to

see the effect of lower temperature instead of a higher temperature on the acid dissociation

constant.

A digital thermometer with a temperature range -10.0°C- 110.0°C and a digital water bath with a 1

L capacity and temperature range 15.0°C-65.0°C should be used to reduce the fluctuations in the

temperature of the reaction mixture in each trial. The desired temperature for the titration should

be obtained from the digital water bath by using the adjustment handle of the temperature on the

bath first. The acid solution and the phenolphthalein in the erlenmeyer flask should be brought to

the desired temperature by placing them into the bath before titration. The sodium hydroxide

solution should also be brought to the desired temperature by placing it into a beaker and then

placing the beaker into the bath. The sodium hydroxide solution could then be poured into the

burette just before the titration. Although its temperature will change, this way can reduce the

error caused by the temperature fluctuations during the titration process. The temperature can be

increased to the desired temperatures and the steps to determine the acid dissociation constants can

be conducted with minimal loss of heat to the surroundings. This way can ensure that the error due

to temperature fluctuations is reduced during the investigation.

CE:

A.1:C

A.2:C

A.3:C

Comment [a13]: Her temperature icin titration yapmak ve de, titrationdan hemen once pH olcmek

deney prosedurundeki faultlari da duzeltebilir.sicaklik arttigi icin eval=poration

olmustu ve cionc degismisti demistik ya.


39

BIBLIOGRAPHY

The number before the source indicates the footnote number of that source in the report.

Books

(2) Atkins, P.W. (1978). “Physical Chemistry”. Oxford University Press. Print.

(15) Green, J., & Damji, S. (2008). “Chemistry”. Melton: IBID. Print.

(14) Housecroft, C. E., Sharpe, A. G. (2008). “Inorganic Chemistry” (3rd ed.) Chapter 6: Acids, Bases

and Ions. Print.

(16) Jensen, W.B. (2003). “The Universal Gas Constant R”. J. Chem. Educ. 80 (7): 731. Print

(19) Kurt, C, Bitter, J. (2005). “Sodium Hydroxide”. Ullmann's Encyclopedia of Industrial Chemistry,

Weinheim: Wiley-VCH. Print.

(5) Prant, J.A. (1989). “Chemical Components”. Cambridge University Press. Print.

(3) Rossotti, F.J.C., Rossotti, H. (1961). “The Determination of Stability Constants”. McGraw–

Hill. Chapter 2. Print.

(11) Whitney, W.D., Smith, B.E. (1911). “Titrimetry”. The Century Dictionary and Cyclopedia. The

Century co. pg. 6504. Print.

Websites

(18) Accuracy and Precision. n.p., n.d. Web. 17 April. 2012. URL:

<http://astro.physics.uiowa.edu/ITU/glossary/percent-error-formula/>

(7) Acetic Acid Safety Sheet. MSDS, n.d. Web. 6 May. 2012. URL:

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<www. scifun.chem.wisc.edu/.../pdf/aceticacid.pdf>

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<www.elmhurst.edu/~chm/.../184ph.html>

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< http://www.britannica.com////>

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