Exp 6 acid and base titration

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CHM 420

CHM 420

GENERAL CHEMISTRY 

GENERAL CHEMISTRY 

NAME

NAME

:

:

SITI NORMALIA

SITI

NORMALIA BINTI

BINTI SULAIMAN

SULAIMAN

:

2012237674

:

2012237674

GROUP

GROUP

:

:

ASB1Ac

ASB1Ac

EXPERIMENT

EXPERIMENT 6

6

:

:

ACIDS

ACIDS AND

AND BASES

BASES

LECTURE

LECTURE NAME

NAME

:

:

ANISAH

ANISAH RAFIDAH

RAFIDAH BT

BT AHMAD

AHMAD

EXPERIMENT

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EXPERIMENT 6

 Acids and Bases

OBJECTIVE

The main purpose of this experiment is to study the properties of acidic or basic substances using indicators and a pH meter.

INTRODUCTION

Acid and bases undergo complete or incomplete when dissolved in water and are called strong acids, strong bases, weak acids or weak bases. The experimental determination of the pH of a solution commonly performed either by the use of indicators or pH meter. The acid or base dissociation (ionization equilibrium) constant, Ka or Kb can be determined experimentally. A sample of weak acid (HA) is dissolved in water and then divided into two equal-volume portions. When one portion is titrated with a sodium hydroxide solution, all HA molecules present converted into A-ions.

OH- + HA H2O + A

-Ka = [H+][A]/[HA] = [H+] 10-pH

The number of A ion produced is equal to the number of moles of HA in the original half-portion, and is also equal to the number of moles of HA in the unused portion of weak acid. The value of Ka can be determined by measuring the pH of a half-neutralised sample of the acid.

When the acid-base titration is carried out, the end point of the titration can be determined when the indicator changes colour. The change of pH can be measured using a pH meter.

CHEMICAL AND APPARATUS

Sample solution A Test tubes pH meter Sample solution B Conical flask 

Unknown acid solution Beaker Indicators (thymol blue. bromophenol blue)

methyl orange and phenolphthalein)

Pipette (25mL) Burette

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PROCEDURE

For each of the following reactions, the observation was recorded.  A. pH using indicator

1. 1 to 2 mL of sample solution A was placed in small test tube 1, 2 and 3. 2. 1 to 2 drops of thymol blue was added to test tube 1.

3. The colour of the solution was then recorded.

4. Step 2 was repeated by replacing the indicator. Bromophenol blue into test tube 2 and methyl orange into test tube 3.

5. The colour of the solution was then recorded.

B. Determination of Ka for a weak acid.

1. 0.2 M of NaOH solution was filled in the burette.

2. 25mL of an unknown weak acid solution was pipette into a conical flask. 3 to 4 drops of phenolphthalein indicator was added into the flask.

3. The acid solution was titrated until a faint persistent pink colour appeared.

4. 25mL of unknown acid solution was added into the neutralised solution. The solution was stirred well. The pH of the solution was recorded.

C. Strong acid-base titration

1. The burette was filled with 0.2 M HCl.

2. 25mL of 0.2 M NaOH was transferred into the 100mL beaker using a pipette. 3 to 5 drops of phenolphthalein was added into the solution.

3. A pH meter was calibrated at pH 7 and pH 4. 4. The pH of NaOH was measured and recorded.

5. 10mL of HCl was carefully added into the NaOH solution. The solution was mixed gently and the pH was measured.

6. The acid was continued to be added and the pH was measured a indicated.

PRECAUTIOUS

1. Safety goggles must be worn during the entire experiment.

2. In part C of the procedure ensures to submerge the tip of the pH meter into the solution for entire experiment.

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DATASHEET EXPERIMENT 6  ACIDS AND BASES  A. pH using indicator

Indicator Colour change pH of colour change

Colour change in sample 1

Colour change in sample 2 Phenolphthalein Red-yellow 8.2-10.0 Colourless Colourless to

light pink  Bromophenol blue Yellow-blue 3 - 4.7 Colourless to pale yellow Colourless to pale blue Methyl orange Red-yellow 3.2 - 4.4 Colourless to

red

colourless to orange

B. Determination of the Ka for a weak acid

pH of half neutralisation solution of unknown weak acid : 4.01

C. Strong acid-base titration

a. HCl added (mL) 0 10 15 20 23 25 26 27 30 b. Measured pH 12.8 12.2 11.8 10.6 6.0 5.9 5.8 5.8 5.7

QUESTIONS

1. Estimate the pH of Sample 1 and Sample 2.

Bromophenol blue is a dye used as a pH indicator, changing from yellow to blue over the pH range 3.0 to 4.7. Methyl orange is another form of dye used as an indicator, changing from red to orange-yellow over the pH range 3.1—4.4.

Sample 1

As phenolphthalein was added into the solution, the colourless solution remains unchanged showing us that the solution could never be in a basic condition. When bromophenol blue was added, the colourless solution then turns from colourless to pale yellow coloured solution. Next, when it was added a few drops of methyl orange, the colourless solution become reddish solution. Both bromophenol blue and methyl orange proves that sample 1 has low pH. Therefore, from this observation we can conclude that the range of pH for sample 1 is higher than 2.8 though less than 4.4.

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Sample 2

When phenolphthalein was added into sample 2, the colourless solution turns to pink. This indicates that the pH is in obviously basic solution. When bromophenol blue was added, the colourless solution then turns from colourless to pale blue coloured solution, another prove that the solution has a high pH value. Next, when it was added a few drops of methyl orange, the colourless solution become orange solution, also indicating the sample to be somehow in higher range of pH value. From this observation we can conclude that the range of pH for sample 2 is higher than 7 though less than 9.

The most possible pH is 8.

2. Which indicators bracketed the pH colour change of sample 1?

In my opinion, the indicator should be methyl orange as when it was added into sample 1, the colourless solution quickly change to red, indicating directly that the solution is acidic. Even when bromophenol blue do change the colour of the solution, the acidity of the solution is uncertain as it only change to pale yellow instead of bright yellow.

3. Which indicators bracketed the pH colour change of sample 2?

In my opinion, the indicator should be phenolphthalein as just when the indicator was added into the sample, it quickly turns from colourless to pink. Phenolphthalein only shows colour change whenever it was in a basic solution thus proving us directly that the solution is in basic condition.

4. From the observed pH of the unknown weak acid, calculate a) [H+] in the solution and Ka

pH =

(log [H+]) 4.01 =

(log [H+]) [H+] = 10(

4.01)

= 9.77 ×10−5 

In water, NaOH dissociated to form

 → 

+ +



From the equation,

pH = 4.01

[



+] = 9.77 ×10−5 

[A¯] = 9.77×10−5 

NaOH = 0.2

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b) Percent ionization

Percent ionization = [

+]



×100%

= 9.77 ×10−5 / 0.2 × 100%

= 4.885×10−4×100%

= 0.04885%

5. Construct a titration curve by plotting measured pH versus volume HCl (mL) added. a) What is the pH range for the colour change of phenolphthalein as shown in the

plotted graph? 6 <



<10.8

b) What is the pH of the equivalence point in this titration? =10.8+6

 2

= 16.8 2

= 8.4

c) Explain why phenolphthalein was used in this experiment?

Phenolphthalein is another commonly used indicator for titrations, and is another weak acid. In this case, the weak acid is colourless and its ion is bright pink. Adding extra hydrogen ions shifts the position of equilibrium to the left, and turns the indicator colourless. Adding hydroxide ions removes the hydrogen ions from the equilibrium which tips to the right to replace them - turning the indicator pink. The half-way stage happens at pH 8.2. Since a mixture of pink and colourless is simply a paler pink, and therefore made it easy for us to detect whenever the reaction had happen.

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0 2 4 6 8 10 12 14 0 5 10 15 20 25 30 35 pH

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DISCUSSION

The main purpose of this experiment is to study the properties of acidic or basic substances using indicators and a pH meter.

Part A of the experiment is when we use various indicators to determine the pH value of two unknown sample, sample 1 and sample 2. Thymol blue, bromophenol blue and methyl orange was used here. These 3 different indicators help us to determine the pH of the solution.

Thymol blue (thymolsulphonephthalein) is a brownish-green or reddish-brown crystalline powder that is used as a pH indicator. It is insoluble in water but soluble in alcohol and dilutes alkali solutions. It transitions from red to yellow at pH 1.2 to 2.8 and from yellow to blue at pH 8.0 to 9.6. Bromophenol blue is a dye used as a pH indicator, changing from yellow to blue over the pH range 3.0 to 4.7. Methyl orange is another form of dye used as an indicator, changing from red to orange-yellow over the pH range 3.1—4.4.

Basically, the acid-base indicator serves as an easy-to-spot signal during the experiment. Once the pH strays into the specified range, the solution starts to change color. Whenever thymol blue indicator was added into sample 1, the colourless solution remains unchanged. This indicates that the pH could not be any less than 2.8. When bromophenol blue was added, the colourless solution then turns from colourless to pale yellow coloured solution. Next, when it was added a few drops of methyl orange, the colourless solution become reddish solution. Both bromophenol blue and methyl orange proves that sample 1 has low pH. Therefore, from this observation we can conclude that the range of pH for sample 1 is higher than 2.8 though less than 4.4. The most possible pH is for sample 1 is 3. As for sample 2, whenever thymol blue indicator was added, the colourless solution turns to light pink  solution. This indicates that the pH is in much basic solution. When bromophenol blue was added, the colourless solution then turns from colourless to pale blue coloured solution, another prove that the solution has a high pH value. Next, when it was added a few drops of  methyl orange, the colourless solution become orange solution, also indicating the sample to be somehow in higher range of pH value. From this observation we can conclude that the range of pH for sample 2 is higher than 7 though less than 9. The most possible pH is 8.

Next, the experiment proceed to part B where as it displays the simulation of strong base and weak acid reaction. NaOH and an unknown weak acid were used in this experiment. pH meter was used to determine the pH of the mixed solution. From the pH value that we have obtained, we can determine the [H+] value by using the formula:

pH =

(log [H+]) 4.01 =

(log [H+]) [H+] = 10(

4.01)

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And once the concentration is obtained, the Ka value is also obtainable. In water, NaOH dissociated to form

 → 

+ +



From the equation,

pH = 4.01 [



+] = 9.77×10−5  [A¯] = 9.77 ×10−5  NaOH = 0.2 Ka = ([H+] [A¯]) / [HA] Ka = ([



+] [



−]) / [



] Ka = (9.77×10−5 )(9.77×10−5 )/ 0.2 = 9.55×10−9/0.2 Ka = 4.77×10−8 Percent ionization = [

+]



×100% = 9.77×10−5 / 0.2× 100% = 4.885 ×10−4×100% = 0.04885%

Proceeding to the part C of the experiment, is the simulation of the reaction happen when strong acid and strong base were in contact with each other. When an acid is titrated with a base, there is typically a sudden change in the pH of the solution at the equivalence point. If  a few drops of indicator solution have been added, this sharp decrease in pH causes an abrupt change in color, which is called the endpoint of the indicator. The actual magnitude of  the jump in pH, and the pH range which it covers depend on the strength of both the acid and the base involved, and so the choice of indicator can vary from one titration to another. In our case, we use sodium hydroxide and hydrochloric acid. The pH changes quite slowly at  the start of the titration, and almost all the decrease in pH takes place in the immediate vicinity of the endpoint. For the first part of the graph, we have an excess of sodium hydroxide. The curve will be exactly the same as when we add hydrochloric acid to sodium hydroxide. Once the acid is in excess, there will be a difference.

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CONCLUSION

The experiment is a success. For part A, the pH value for sample 1 is 3, pH value for sample 2 is 8. Part B resulting [H+] is 9.77×10−5 , Ka is 4.77×10−8 and percent ionisation is 0.4885%. And as for

part C, the endpoint of the titration is at pH 8.4. REFERENCE  http://www.answers.com/topic/ph-indicator  http://www.thebigger.com/chemistry/ionic-equilibria/calculate-the-percent-ionization-of-0-20-m-solution-of-hydrocyanic-acid-hcn-ka-for-hcn-4-9-x-10-10/  http://www.chemguide.co.uk/physical/acidbaseeqia/phcurves.html  http://malaysia.answers.yahoo.com/question/index?qid=20100115103012AAkuaWB  Shakhashiri, B. Z., Chemical Demonstrations: A Handbook for Teachers of Chemistry, Vol. 1 ;

The University of Wisconsin

 Press: Madison 1983; pp 280–285.

 Basic Chemistry, 9th Edition International Student Version Leo J. Malone (Saint Louis

Univ.), Theodore Dolter (Southwestern Illinois College) March 2012, ©2013 0 2 4 6 8 10 12 14 0 10 20 30 40 pH volume of HCl added (mL) Acid-base titration curve

Figure

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References

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