ACID BASE TITRATION EXPERIMENT 1 (PRACTICAL ANALYTICAL CHEMISTRY) PDF

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Summary

PART A: PREPARATION OF STANDARD SOLUTION
PART B: STANDARDIZATION OF THE STANDARD SOLUTION
PART C: ANALYSIS OF VINEGAR SOLUTION
POST LAB QUESTION...

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UNIVERSITY MALAYSIA SARAWAK FAKULTY OF RESOURCE SCIENCE AND TECHNOLOGY

STK1211- PRACTICAL ANALYTICAL CHEMISTRY INSTRUMENTATION AND MEASUREMENTS

Experiment 1: Acid-Base Titration Report Due Date : 29 October 2021 Lecturer’s Name : Dr Showkat Ahmad Bhawani Group Members: 1. Gibson Aznil Ak Kinchun (79478) 2. Aida Farhana Binti Ahmad Salimi (78676) 3. Olivia Yeo (80926) 4. Izzfeesya Irdina Binti Idris (78073)

1

Table of Content

Contents

Page

Introduction

3-4

Objectives

4

Procedure & Apparatus,Reagents

5-7

Result

8-9

Calculation

10-13

Discussion

14-20

Questions and answers

21-22

Conclusion

23

Attachments

24-25

References

26-27

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EXPERIMENT 1: ACID-BASE TITRATIONS

INTRODUCTION Titration is a laboratory technique used to determine the concentration of a solution using another solution with a known concentration. One of the solutions involved in a titration is used as a standard solution. A standard solution of titrant in a burette is gradually applied to react with an analyte with an unknown concentration in an Erlenmeyer flask . The standard solution can be classified as either primary or secondary. A primary standard solution is prepared by dissolving an accurately weighed pure solid of a known molar mass in a known volume of distilled water. A primary standard is used to determine the molarity of the other standard solution, known as a secondary solution. A standard solution can be prepare by accurate weighing of the solute, followed by precise dilution to an exactly known volume in a volumetric flask .The most common standard solutions used in acid-base titration . An acid-base titration is an experimental procedure used to determined the unknown concentration of an acid or base by precisely neutralizing it with an acid or base of known concentration. Solutions of sodium hydroxide are commonly used in titration analysis of samples containing an acidic solute. The NaOH solution used in titration need to be standardized because they contain impurities . Solid NaOH is hygroscopic (it absorbs moisture). Thus it is difficult to obtain its accurate mass. The standardized NaOH becomes the secondary solution and can then be used to determine the concentration of other acids such as H2SO4 acid. NaOH solutions are generally prepared to be approximately a given concentration. . By measuring how many volume of the approximately prepared NaOH are necessary to react completely with a weighed

sample of a known primary standard acidic substance, the concentration of the NaOH solution can be calculated. Once prepared, the concentration of a NaOH will change with time. As a consequence, NaOH solutions must be used relatively quickly. The purpose of the titration is the detection of the equivalence point, the point at which chemically equivalent amounts of the reactants have been mixed called end point. In titration analysis, a pH indicator is usually added in the analyte solution to indicate the end point of titration.. In an acid-base titration analysis, there should be a sudden change in pH when the reaction is complete. For example, if the sample being titrated is an acid, then the titrant to be used will be basic. When one excess drop of titrant is added, the solution being titrated will suddenly become basic.The equivalence point is found in a titration by adding trace amounts of a substance, turns color when the equivalence point is reached..When a strong acid is titrated with a strong base, or vice versa, the pH of the solution will be about 7.0 at the equivalence point. Phenolphthalein is the indicator used in this experiment. Phenolphthalein is colorless in acidic solutions and turns pink in alkaline solutions.

OBJECTIVES i.

Standardize 2g NaOH pellets

ii. Prepare a standard solution of NaOH iii. Measure the amount of acetic acid in a solution of vinegar iv.

Acquire the correct techniques of titration

v.

Understand the use of indicators in titrations

APPARATUS Burette with stand Pipette 1L plastic bottle with stopper 250 mL Erlenmeyer flasks Retort stand with clamp 12

REAGENTS Sodium hydroxide pellets Potassium hydrogen phthalate (KHP) Phenolphthalein Unknown vinegar

PROCEDURE Part A: Preparation of the Sodium Hydroxide Solution 1. 500 mL volumetric flask and stopper were cleaned and rinsed. The flask was labelled “Approx. 0.1M NaOH”. 250mL of distilled water was put into the flask.

2. 2.0g of sodium hydroxide pellets were weighed out and transferred to 500mL flask. The flask was stopper and shook to dissolve the sodium hydroxide. 3. When all the sodium hydroxide pellets have dissolved, distilled water was added to the bottle until the mark on the neck of the flask. Stopper and shake thoroughly to mix.

Part B: Standardization of the Sodium Hydroxide Solution 1. The burette was set up in the burette clamp. The burette was rinsed and filled with the sodium hydroxide solution that was just prepared. 2. Three 250mL Erlenmeyer flasks were cleaned with water, and then were rinsed with distilled water. They were labelled as 1, 2, and 3. 3. The bottle of dried KHP was removed from the oven. When the KHP was completely cool, three samples of KHP between 0.6 and 0.8g were weighed, one for each of the Erlenmeyer flasks. The exact weight of each KHP sample was recorded to the nearest mg (±0.001g). 4. 100mL of distilled water was added to KHP sample. 2-3 drops of Phenolphthalein indicator solution were added. The KHP sample was swirled to dissolve completely. 5. The initial reading of the NaOH solution in the burette was recorded to the nearest 0.02mL. 6. NaOH solution from the burette was added to the sample in the Erlenmeyer flask, the flask was constantly swirled during the addition. 7. When the titration approached the endpoint, one drop of NaOH was added at a time, with constant swirling until one single drop of NaOH caused a permanent pale pink color that does not to fade on .

8. Step 4-7 was repeated with the other two KHP samples. 9. The molecular mass of KHP given was 204.2, calculate the number of moles of KHP in samples 1,2 and 3. 10. From the number of moles of KHP present in each sample, and from the volume of NaOH solution used to titrate the sample, calculate the molar concentration (M) of NaOH in the titrant solution. The reaction between NaOH and KHP is 1:1 stoichiometry.

Part C: Analysis of Vinegar Solution Vinegar is a dilute solution of acetic acid and can be effectively titrated with NaOH using the Phenolphthalein end point. 1. Three Erlenmeyer flasks were cleaned and labelled as samples 1,2, and 3. 2. The 5 mL pipette was rinsed with small portions of the vinegar solution and the rinsing was discarded. 3. By using the pipette, 5 mL of the vinegar solution was pipetted into each of the Erlenmeyer flasks. About 100 mL of distilled water and 2-3 drops of Phenolphthalein indicator solution were added into each flask. 4. The burette was refilled with the NaOH solution and the initial reading of the burette was recorded to the nearest 0.02 mL. Sample 1 of vinegar was titrated in the same manner as in the standardization until one drop of NaOH causes the appearance of the pale pink color. 5. The final reading of the burette was recorded to the nearest 0.02mL. 6. The titration was repeated for the other two vinegar samples. 7. The molar concentration of the vinegar solution was calculated based on the volume of vinegar sample taken and on the volume and average concentration of NaOH titrant used.

8. The percent by mass of acetic acid in the vinegar solution was calculated given that the formula mass of acetic acid was 60.0 and the density of the vinegar solution was 1.01 g/mL.

RESULTS Part A: Standardization of the Sodium Hydroxide Solution

Particulars

Trial 1

Trial 2

Trial 3

Mass of KHP taken (g)

0.72

0.72

0.75

Final burette reading (mL)

37.00

37.00

38.00

Initial burette reading (mL)

0.00

0.00

0.00

Volume of NaOH used (mL)

37.00

37.00

38.00

Molarity of NaOH solution (M)

0.0958

0.0958

0.0958

Average molarity of NaOH solution 0.0958 (M)

PartB:Anal ys i sofaVi negarSol ut i on

Particulars

Trail 1

Trail 2

Trail 3

Volume of vinegar solution used (mL)

5.00

5.00

5.00

Final burette reading (mL)

44.20

45.20

44.80

Initial burette reading (mL)

0.00

0.00

0.00

Volume of NaOH used (mL)

44.20

45.20

44.80

Molarity of NaOH solution (M)

0.0958

0.0958

0.0958

Molarity of vinegar solution (M)

0.847

0.866

0.858

% mass of acetic acid in vinegar (%)

5.03

5.15

5.09

Average molarity of vinegar solution (M)

0.857

Average % mass of acetic acid in vinegar (%)

5.09

CALCULATION

DISCUSSION

Theoretical background: Titration describes a process where the concentration of an unknown substance is determined by comparing it with a solution of known concentration. The concept that makes titrations possible is finding the equivalence point, identifying when the quantity of the unknown substance is equal to the quantity of the known substance. A titration is a process in which a solution containing a known amount of a substance is allowed to react with a second solution containing an unknown concentration of another substance that will react with the first substance in a known and reproducible manner. The substances are allowed to react until there is some indication that equivalent amounts of the substances have reacted. The solutions are measured from a buret, a long, graduated glass tube with a stopcock at the bottom. In the present case a solution of potassium hydrogen phthalate is prepared in an erlenmeyer flask by dissolving an exactly known weight of pure potassium hydrogen phthalate in water. The sodium hydroxide solution (whose exact concentration is unknown), is delivered from a buret until an amount equivalent to the amount of potassium hydrogen phthalate has been added. This point in the process is called the equivalence point of a titration.

Interpretation of result: This experiment will be done in two parts: (1) preparation and standardization of the sodium hydroxide solution and (2) analysis of a vinegar solution. Standardization is the process of determining concentration in an unknown solution by titrating it with a solution of known concentration.In this experiment,KHP ( Pottasium Hydrogen Phthalate) that used was a brittle, white, crystalline material used in the experiment. Before the crystals could dissolve in water, they had to be vigorously stirred. The acidic solution that resulted was translucent, with a little amount of undissolved KHP granules. The transparent NaOH solution turned pink when it came into touch with the transparent phenolphthalein in the KHP solution and then became transparent after shaking . Three samples of known potassium hydrogen phthalate were titrated with the prepared 0.1 M sodium hydroxide solution during the standardisation process. Each sample contained 0.6 g until 0.8 g potassium hydrogen phthalate (measured to the nearest 0.1 mg) in 100 mL water, as well as two drops of phenolphthalein indicator to lighten the solution and signal the conclusion of the titration . Based on the data that has been given, we obtained data of mass of KHP , final and initial burette reading for trial 1,2 and 3 for part A. We also obtained data for part B which are volume of vinegar solution used, final and initial burette reading for trial 1,2 and 3. From the data part A given, we can calculate the molarity of NaOH in titrant solution. The molarity of NaOH in titrant solution is 0.0958 M for trial 1 , 2 and 3 since the mass of KHP taken and volume of NaOH are the same. So, the average molarity of NaOH in titrant solution is 0.0958M . To calculate it, we used mass of KHP taken to find mole of itself by using the following formula :

 Mol of KHP = mass of KHP / MW ( g / mol)  Molarity of NaOH = mol of KHP / volume of NaOH (L)

For part B, the volume NaOH used to get the end point of titration with 5mL vinegar in trial 1 is 44.20 mL, in trial 2 is 45.20 mL , while in trial 3 is 44.80 mL. We can calculate the Molarity of vinegar based on the concentration of volume and average concentration of NaOH titrant used , by using the following formula :

M1 V1 = M2 V2  M1 : Molarity of vinegar

; V1 : Volume of vinegar

 M2: Molarity of average NaOH

; V2 : Volume of NaOH used (Part B)

* volume must in liter (L) Next, we can calculate % mass of acetic acid in vinegar by using the following formula:  Moles of Acetic acid (n) = ( M NaOH)( V NaOH)  Mass of Acetic Acid (g) = (n acetic)(MW)  Mass of Vinegar = (V vinegar in mL )( density of vinegar g/mL) = (5 mL)( 1.01 g/mL) = 5.05 g  % mass of acetic acid = ( mass of acetic acid, g / mass of vinegar, g )( 100% )

We can calculate the number of moles of NaOH that we have added to the flask by taking the given molarity of the NaOH , which we will say is 0.0958 M (M means moles of substance/L). The number of moles of NaOH will be the same as the moles of acetic acid, which means that the number of moles of acetic acid in the 5.00 mL of vinegar is also 0.0958 moles.

Since we have mole of acetic acid from ratio 1:1 with NaOH from the stoichiometry equation, will say 0.0958 moles of acetic acid , we can calculate molarity of vinegar solution by dividing mole of solute (mole of acetic acid / mole of NaOH x Volume of NaOH ) with volume of vinegar solution . For example, in trial 1 we got 0.847 M , trial 2 we got 0.866M , and trial 3 we got 0.858M.

Safety precaution & suggestion to reduce error

Moreover,there are some weakness of this experiment in which it was by allowing the KHP turns to dark pink colour.The impact implies when there were accuracy concerns in the computation of the concentration due to the excessive use of NaOH.But,this errors can avoid by reducing the rate of flow of NaOH from the burette as soon as the solution begins to turn dark pink, and shut off the supply when the pink can no longer be removed. We can avoid adding too much NaOH this way. End point determination is perhaps the most common source of systematic error in titrations. Because the end point is determined visually, there can be a significant difference in when a titration is halted from one individual to the next. Because each person's ability to see a slight change in colour differs, a "faint pink" will appear differently to each person, resulting in a different amount of titrant being delivered depending on who is conducting the

experiment, resulting in a difference in the molarity of the NaOH solution calculated after the experiment. As in addition, by allowing the KHP to flow along the conical flask's interior walls,inaccuracies would have resulted as a result of this. It's possible that the NaOH didn't react with the exact amount of KHP that was intended. To acquire a more precise volume measurement, make sure the conical flask is directly under the pipette, with no contact with the interior walls.

Moreover, vinegar is a popular home item that contains both acetic acid and other compounds. By titrating a sample of vinegar with a standard solution of NaOH, the molar concentration of acetic acid in the sample can be determined. NaOH(aq) + CH3COOH(aq) = CH3COONa(aq) + H2O (l) A neutralising reaction happens when basic sodium hydroxide is added to acidic acetic acid. The vinegar also contains an indication known as phenolphthalein. When too much NaOH is added to make the solution more basic, this indicator colours it dark pink. As a result, when the solution develops a pale pink colour, it has been neutralised satisfactorily. Because they are both at a 1:1 ratio in the preceding equation, the amount of NaOH needed to standardise the vinegar can be used to determine the amount of acetic acid in the vinegar. As a result, the number of moles of NaOH needed to neutralise the acid must be the same as the number of moles of the acid.

The reactions that occurred during the experiment were neutralisation reactions, which meant that at the end of the experiment, the moles of acid equaled the moles of base. This factor was applied to the formula ‘moles = concentrations x volume to get

the molar concentration of acetic acid. There were numerous errors that influenced the outcomes of this experiment, contributing to the total 50 percent inaccuracy. The strength of the sodium and sodium hydroxide was a crucial component that influenced the outcome of this experiment. If one of these compounds is exposed to the air, their strength begins to deteriorate. Sodium and sodium hydroxide were both kept open to interact with the environment for a period of time during the experiment. Because the strength was weakened, the final response did not match the theoretical value precisely, implying that the values used to calculate the molar concentration were not as reliable. To avoid this error from impacting the experiment's results, try to maintain the sodium and sodium hydroxide in an enclosed area at all times, limiting the amount of time they are exposed to the atmosphere. Furthermore, because all pieces of apparatus contain an element of uncertainty, the equipment utilised may have contributed to the inaccuracy. These uncertainties are then employed in calculations to maintain the same level of uncertainty as the amount of material used. This value takes into account both the uncertainties that existed at the time the solutions were made (for example, the uncertainty of the mass balance used to calculate the amount of sodium required to make sodium hydroxide) and the uncertainties that existed during the transfer of the solution from one instrument to another. More precise equipment, such as a more accurate mass balance, can help to reduce these uncertainties. Limiting the amount of solution transferred from one container to another will also help to reduce mistake. Because the individual performing the experiment was forced to read off multiple readings from the pipette and burette, some of the inaccuracy in this experiment can

be attributed to human judgement. This error can be minimised by making sure that readings are taken at eye level and that the same person is doing the readings every time, as judgement differs from person to person. In conclusion, this experiment found the molar concentration of acetic acid in vinegar to be 0.857 mol/L. However, this value was 50% inaccurate due to a errors that occurred while conducting the investigations.

QUESTIONS AND ANSWERS 1.

Give the definition of indicators.

Definition of indicators are dye which shows different colour in acidic and basic condition. This chemical indicates a change in pH by a clean colour change and the equivalence point is exactly when the titration mixture changes colour permanently. There are many different indicators for different pH ranges and combinations of weak or strong acids reacting with weak or strong bases. It is very sensitive to pH change. The point in th...