Experiment 3 - Acid and Base Titration

А

n а l y t i c

MALAYAN COLLEGES LAGUNA

а

l C h e m i s t r y

C H E 1 3 0 L

EXPERIMENT NO. 3

ACID- BASE TITRATION I.

OBJECTIVES Upon completion of the experiment, the student should be able to: 

prepare standard solutions of NaOH and HCl;



standardize the acid and base solution;



determine the equivalent weight of an unknown acid; and



apply the various techniques involved in the preparation of solutions, standardization of solutions, comparison of solutions, and analysis of unknown solutions for acid-base titrations.

II.

A. LABORATORY EQUIPMENT / INSTRUMENTS Equipment/ Accessories 50 mL beaker 250 mL Erlenmeyer flask 500 mL volumetric flask Glass funnel Analytical balance Buret holder 50 mL buret 125 mL Erlenmeyer flask 25 mL volumetric pipet Stirring rod Iron stand

Quantity 2 2 1 2 1 2 2 2 1 1 1

B. CHEMICALS AND REAGANTS Chemical/ Reagent NaOH pellets 1.0 M HCl Phenolphtalein Distilled water in wash bottle Potassium acid phthalate (KHP, FM=204.2 g/mol) Unknown acid

Experiment 3: Acid and Base Titration

Page | 1

А

n а l y t i c


а

l C h e m i s t r y

C H E 1 3 0 L

EXPERIMENT NO. 3

ACID- BASE TITRATION III.

DISCUSSION OF FUNDAMENTALS Introduction There are many types of volumetric analyses in the study of analytical chemistry. One of the most common and easiest methods is titration/ titrimetry. This process involves an analyte of unknown concentration, and a titrant, that usually has a known concentration. One of the best working formulas for this is the dilution equation, which is;

  ,(Formula 1) wherein C

1

is

the molarity of the titrant, V1 the volume of the titrant that is used to complete the reaction, V 2 the volume of the analyte which has an unknown concentration, and C 2, which is what is usually the unknown. In titrimetry there are also lots o types, and in this experiment, we are focusing on acidbase reactions. The reaction involves mainly a base and an acid, from which either one has a known concentration. Certain procedures must be followed in the process of acid-base titrations, such as standardization of the acid/base solution, most laboratory techniques, etc. etc.

General Description Titration or titrimetric analysis is a method of slowly adding a titrant, usually of known concentration to an analyte to determine the amounts for a complete reaction (Figure 1, Hage and Carr 2011). Titration reactions are used for a number of purposes such as standardization, determination of amounts of analyte present in samples, and determining identities of unknown samples. An acid-base titration involves reactants which are acids and bases. The amounts of reactants needed to have a complete reaction are referred to as the stoichiometric amounts. In titration, this is referred to as the equivalence point, meaning that at this point, equivalent amount of the titrant has reacted to an equivalent amount of the base. Acid-base reactions do not show visible physical changes when products are formed (i.e. do not form solids nor vapors, nor show changes in temperature) and thus the equivalence point is difficult to determine. In this regard, aside from the acid and base, another important component of an acidbase titration is an acid-base indicator. The acid-base indicator changes its color upon adding a very slight excess of the titrant. This amount is a visual and experimental estimate of the equivalence point and is termed as the end point (Hage and Carr 2011). Another method that is used to determine the end point is to use a pH meter to monitor pH of the solution as the volume of added titrant changes and plotting these values to obtain a titration c urve (Figure 1, Hage and Carr 2011).


Page | 2

А

n а l y t i c


а

l C h e m i s t r y

C H E 1 3 0 L

EXPERIMENT NO. 3

ACID- BASE TITRATION

Figure 1. Typical Titration set-up and a sample of titration curve

Since the most usual application of acid-base reaction is determining the amount or concentrations of analytes, one of the reacting solutions should be of known concentration. The method of determining the exact concentration of standard solutions is referred to as standardization. The method makes use of a solid pure substance, a primary standard to determine the exact concentration of a secondary standard solution. Concentrations of solutions can be expressed in units of molarity (moles solute/liter of solution) or normality (equivalents solute/liter of solution).

Theory Titrations, from French origins ‘titre’ means rank, and from the Latin word ‘titulus’, it means

title. It generally means that titration is the determination of the rank of concentration of the


Page | 3

А

n а l y t i c


а

l C h e m i s t r y

C H E 1 3 0 L

EXPERIMENT NO. 3

ACID- BASE TITRATION solution, with water as a reference with a rank of pH 7. The analyte is the one being known of its rank, and the titrant is the one who will determine this rank. Titrimetry is based on how certain reactions will proceed, so therefore not all reactions can be considered to be used in a titration process. A type of reaction can be used in a titration process if it is: has a fast rate of reaction; it proceeds in a stoichiometric way (meaning, the moles of the unknown can be compared with that on the known substance); there is a known or a visual way of detecting the completion of the said reaction, and; if the free energy (Gibbs energy, represented by

) is large enough to make the reaction spontaneous.

For this specific instance, we are talking about acid-base titrations. The chemical reaction that is involved between acids and bases are called as neutralization reactions, from which the product is a salt, and a by-product of water. The water is produced by combining the hydrogen (hydronium) ions present in the acid, to the hydroxide ions present in the base. By knowing the concentration of an acid using a standardized basic solution is called as alkalimetry , and the other way around we call it acidimetry .

In an acid-base reaction, a stage which shows its completion is known as end point. Equivalence point is a stage from which the amount of the titrant (since we are talking about titrations) added is exactly and by stoichiometric basis, it is equivalent to the amount of the analyte in the titrated solution. The end point is detected by some visual changes, but if there are no visual changes present, an addition of an auxiliary reagent known as an 'indicator' will be very much needed. An indicator is a chemical reagent that is added to the analyte before the reaction is initialized, to recognize the completion of the titration/reaction. When the reaction is complete, it significantly gives off a recognizable color change. For a neutralization reaction, a pH indicator is needed, i.e. phenolphthalein. The end point and the equivalence point may not be identical. End point is usually detected only after adding a slight excess of the titrant, but the equivalence point is the true completion of the reaction.

Application Some applications of acid-base titrations would be: 

Creating household items



Chemical Industry



An important technique, titrating and back-titrating knowledge is very useful for a chemist



Balancing the pH of food, fertilizers, and the like.


Page | 4

А

n а l y t i c


а

l C h e m i s t r y

C H E 1 3 0 L

EXPERIMENT NO. 3

ACID- BASE TITRATION IV.

METHODOLOGY The first step was the cleaning of the buret, which was checked of leaks

and was then rinsed with small

The cooled solution was quantitavely

amount of the titrant. After, it was

transferred into a 500-mL volumetric

filled with the titrant up to the zero

flask.

mark.

NaOH pellets needed to prepare 500mL of 0.05M NaOH were fetched

from the prep room and the mass

The solution was diluted with distilled water up to the mark and was mixed by shaking.

was calculated.

The softdrink bottle was rinsed with In a 50-mL beaker, the required amount of NaOH pellets was weighed in a balance.

small portions of the NaOH solution. The NaOH remaining was transferred to this softdrink bottle.

The NaOH pellets were dissolved in

The bottle was labelled with the

water under the fumehood.

concentration, date of preparation and the name of the group who prepared the solution.

The beaker was cooled by placing it

Figure 2. Buret cleaning and preparation of

under running water, careful of not

standard NaOH solution.

getting any inside the beaker.


Page | 5

А

n а l y t i c


а

l C h e m i s t r y

C H E 1 3 0 L

EXPERIMENT NO. 3


The amount of 1.0M HCl needed to prepare 500-mL of 0.05M HCL solution was calculated.

This was diluted with distilled water

until the mark and was shake very well. A 500-mL volumetric flask was filled with 250-300 mL distilled water. The softdrink bottle was rinsed with small portions of the HCl solution. The HCl remaining was transferred to Using a volumetric pipet, calculated

this softdrink bottle.

volume was was transferred to the 500-mL volumetric flask. This was done under the fumehood. The bottle was labelled with the Figure 3. Preparation of standard HCl solution

concentration, date of preparation and the name of the group who prepared the solution.


Page | 6

А

n а l y t i c


а

l C h e m i s t r y

C H E 1 3 0 L

EXPERIMENT NO. 3


Accurately 0.25g KHP (previously dried at 105C for one hour) into a labeled Erlenmeyer flask was weighed.

The final volume reading in your buret was recorded.

It was dissolved in about 50 mL distilled water.

The normality of the standard NaOH solution was calculated into the proper number of significant figures for each trial.

2 to 3 drops of phenolphthalein

indicator was added when the solution was at room temperature. The average normality and the average deviation were computed The buret was filled with the t itrant up to the zero-mark. Initial volume

using the average normality for succeeding calculations.

reading in the buret was recorded.

The standard NaOH solution was kept The KHP solution was titrated with the 0.05 M NaOH solution to the faint

for future use. Figure 4. Standarization of NaOH solution

pink endpoint.


Page | 7

А

n а l y t i c


а

l C h e m i s t r y

C H E 1 3 0 L

EXPERIMENT NO. 3


One buret was filled with the standard HCl solution and another with the standard NaOH solution.

The volume used was calculated up to the proper number of significant

figures. The initial volume reading in each buret were both recorded.

The ratio of strengths of the acid and base was calculated.

From the buret, 25 mL of HCl was delivered into an Erlenmeyer flask.

The average ratio for each of the

This was done in two trials.

solutions was calculated.

About 20 mL distilled water was

The volume data from this part and

added into the Erlenmeyer flask along

the average normality for the NaOH

with 2 -3 drops of phenolphthalein.

solution (from Part IV) was used to calculate the concentration of standard HCl solution for the two trials.

The HCl solution was titrated with the NaOH solution up to the first faint pink endpoint. Near the endpoint, titrate dropwise and at the endpoint, split drops.

The average normality and the

average deviation were calculated and the average normality was used for succeeding calculations

The final volume reading in your buret was recorded.

The standard HCL solution was kept for future use. Figure 4. Comparisons of standard solutions


Page | 8

А

n а l y t i c


а

l C h e m i s t r y

C H E 1 3 0 L

EXPERIMENT NO. 3


0.10 g of the unknown acid was accurately weighed into a labeled Erlenmeyer flask. This was done in two trials.

About 50 mL distilled water was diluted and added 2-3 drops of phenolphthalein indicator.

The unknown acid solution was titrated with the standard NaOH solution up to the faint pink endpoint.

Calculate the equivalent weight of the unknown acid for each trial and the average EW. Figure 5. Determination of equivalent weight of unknown acid


Page | 9

А

n а l y t i c


а

l C h e m i s t r y

C H E 1 3 0 L

EXPERIMENT NO. 3

ACID- BASE TITRATION V.

DESCRIPTION OF THE APPARATUS / SET – UP

Figure 2. Burette

Figure 2.1 Burette with labels

A burette (commonly spelled as buret), is a laboratory apparatus that is mainly used for quantitative chemical analyses of liquids. It consists of a long, graduated glass tube with a stopcock (in a liquid burette’s case, on the bottom) that is handled by a burette clamp, which is connected to

an iron stand. The volume that the burette dispenses is controlled by the stopcock, and is accurately measured by the graduations of the glass tube.


Page | 10

А

n а l y t i c


а

l C h e m i s t r y

C H E 1 3 0 L

EXPERIMENT NO. 3

ACID- BASE TITRATION VI.

DATA SHEET

A.

PREPARATION OF STANDARD SOLUTIONS

Table 1. NaOH Standard solution

Condition

Data

Mass of NaOH pellets

1.04 g

Total volume of solution

500 mL

Table 2. HCl Standard Solution

Condition

Data

Volume of concentrated HCl solution

25 mL

Total volume of solution

B.

500 mL

STANDARDIZATION OF NaOH SOLUTION

Table 3. Standardization of NaOH Solution

Conditions

Trial 2

Trial 1

Primary standard Weight of flask, g

TARED Weight of flask + KHP, g Weight of KHP, g

0.25 g

0.25 g

24.5 mL

24.4 mL

NaOH solution Final volume reading, mL


Page | 11

А

n а l y t i c


а

l C h e m i s t r y

C H E 1 3 0 L

EXPERIMENT NO. 3

ACID- BASE TITRATION Initial volume reading, mL

0.00 mL

0.00 mL

Volume used, mL

24.5 mL

24.4 mL

0.04997 N

0.05017 N

Normality of NaOH solution

0.05007 N

Average N  average d

C.

COMPARISON OF STANDARD SOLUTIONS

Table 4. Volume Ratios of Standard Solutions

Trial 1

Trial 2

Final volume reading, mL

25.0 mL

25.0 mL

Initial volume reading, mL

0.00 mL

0.00 mL

Volume used, mL

25.0 mL

25.0 mL


17.8 mL

36.0 mL


0.00 mL

17.8 mL

Volume used, mL

17.8 mL

18.2 mL

Conditions HCl solution

NaOH solution

Ratios 1.000 mL HCl

___0.712__ mL NaOH

0.720 mL

Average  average d 1.000 mL NaOH Average  average d


___0.728_ mL NaOH

__1.404___ mL HCl

__1.374_ mL HCl 1.389 mL

Page | 12

А

n а l y t i c


а

l C h e m i s t r y

C H E 1 3 0 L

EXPERIMENT NO. 3

ACID- BASE TITRATION Normality of HCl solution

0.0356 N 0.03605 N

Average N  average d

D.

0.0365 N

DETERMINATION OF EQUIVALENT WEIGHT OF UNKNOWN ACID

Table 5. Determination of equivalent weight of unknown acid

Trial 1

Trial 2

Unknown Number

G

H

Mass of unknown

0.100 g

0.102 g

Conditions

Weight of flask, g

TARED

Weight of flask + unk acid, g Weight of unk acid , g

0.100 g

0.102 g


45.0 mL

34.1 mL


11.0 mL

11.0 mL

Volume used, mL

34.0 mL

34.1 mL

Equivalent weight

58.74 g/mol

59.74 g/mol

NaOH solution

Average EW  d Probable identity of acid


59.24 g/mol Glacial acetic acid (60.05 g/mol)

Page | 13

А

n а l y t i c


а

l C h e m i s t r y

C H E 1 3 0 L

VII.

EXPERIMENT NO. 3

ACID- BASE TITRATION SAMPLE COMPUTATIONS

                                          )(  )        (                                                                                  11.1

VIII.

RESULTS AND DISCUSSIONS The common unit used for expressing a solution’s concentration is m olarity (M) or molar

concentration. It is defined as the number of moles of solute per liter of solution. The concentration of a basic solution is determined, one way, by titrating it with a known volume of a standard acid solution of known concentration required to neutralize it. The main purpose of titration is the detection of the equivalence point of the solution with unknown concentration. The equivalence point is the point at which chemically equivalent amounts of the reactants have been mixed. The amount of reactants mixed at the equivalence point depends on the stoichiometry of the reaction. A known method for determining either the concentration of a solution of unknown molarity or the number of moles of a substance in a given sample is titration. A chemical reaction is used for this purpose, and the reaction must be fast, complete, and have a determinable end point. Reactions of strong acids and bases meet these criteria generally, and acid-base titrations are among the most important examples of this method.


Page | 14

А

n а l y t i c


а

l C h e m i s t r y

C H E 1 3 0 L

EXPERIMENT NO. 3

ACID- BASE TITRATION In the neutralization reaction of HCl and NaOH, the equivalence point occurs when one mole of HCl reacts with one mole of NaOH. Using an indicator such as phenolphthalein helps to determine when neutralization occurs. Phenolphthalein (C20H14O4) is an example of a substance which undergoes a distinct color change at or near the equivalence point. It changes its color to light pink at pH range 8.2-10.0. The point at which the indicator changes color and the titration is stopped is called the endpoint. Ideally, the endpoint should match with t he equivalence point. Phenolphthalein is colorless in acidic solution and reddish violet in basic solution. The experiment was started off with cleaning of the glasswares and rinsing of the buret with the titrant. This is done so that the buret will be clean of the titrant last used by students before us. After rinsing, there came the preparation of the standard NaOH solution, which proved challenging because filling the volumetric flask to the mark without very precise control took time. Too much distilled water won’t do so the whole method will have to be done again. Using stoichiometry, the

amount of 1.0M HCl to prepare 500-mL HCl solution was determined. Same method as the preparation of the NaOH solution, the HCl was diluted with distilled water and was kept of future use. The standardization of the NaOH solution part was where the titration was done. Acid-base titration involves a chemical reaction known as a neutralization reaction. Titrations allow the accurate determination of concentrations of unknown acids/bases. Analyzing unknown acids/bases, a "standard" solution is needed to react with the unknowns. Making a standard solution requires dissolving an accurately weighed amount of the substance and diluting it to an accurately measured volume. Having these given, the molar concentration can be calculated exactly. But, getting NaOH of sufficient purity to use as a primary standard is usually difficult. Obtaining a standard solution of NaOH using an indirect method is more practical. A solution of an approximate molarity was prepared and standardized against a primary standard of known purity. A KHP (potassium hydrogen pthalate) solution of known molarity was titrated to its endpoint using the standard NaOH solution, to know whether the pH of the prepared standard was accurate. KHP, a large compound (KHC8H4O4) having a molecular mass of 204.2 g/mol, served as the primary standard. It is a monoprotic acid and reacts with NaOH in a simple 1 to 1 relationship according to the following equation: NaOH(aq) + KHC8H4O4(aq) ---> KNaC8H4O4(aq) + H2O(l)


Page | 15

А

n а l y t i c


а

l C h e m i s t r y

C H E 1 3 0 L

EXPERIMENT NO. 3

ACID- BASE TITRATION After standardizing the base with potassium hydrogen pthalate, the standard base was used to standardize the acid to know the concentrations of both solutions for the titration and analysis of the unknowns.

IX.

SUMMARY AND CONCLUSIONS This experiment thoroughly explained us the acid-base reactions in actual, which is in the process of titration. By theory, when an acid is combined with a base that is compatible with each other, it will react to reach as what we call an ‘equivalence point’, from where the moles of the

analyte is equal to the moles of the titrant. This is commonly known as neutralization. But to do this, one must need to have a standardized solution of either an acid or the base, whichever it is that is known. Standardization is done to produce accurate results. Once the solution was standardized, we can now begin the titration process. But in the process, the salt that forms during neutralization is the only visual presence of reaching equivalence point, which is usually at a disadvantage since these salts are always colorless, and is dissolved to water, its by-product. To prevent this, one must add an indicator to the analyte, a chemical that changes in color from when the desired pH level is reached, or when the reaction has reached it equivalence point. Regardless to say, by using the dilution equation we can solve for the concentration of the analyte, with a known volume. All in all, acid-base titrations can be a powerful tool for knowing certain acid-base reactions, for knowing the concentrations of a base or an acid, as long as the titrant is standardized, to ensure accuracy and precision of results.

X.

POST LAB QUESTIONS Post lab Q1 . Will the volume used for titration be affected if the volume of water used to dissolve the KHP is increased to 75mL distilled water? Why or why not?

Answer : No, the concentration of the ac id would be the same since we will only convert the moles of KHP, whereas it the same as converting the grams of KHP to molarity KHP, then turning it to moles KHP by multiplying it to the volume, cancelling both figures.

Post lab Q2 . What is the ideal volume used for one trial of titration? Explain your choice in terms of precision of measurements. Using this volume, what mass of K HP must be used to standardize a solution of approximately 0.035 M NaOH?

Answer: By computing the average of our ex periments, the ideal volume for titration is supposed to be 24.45 mL. Using this volume, we c ompute: Experiment 3: Acid and Base Titration

Page | 16

А

n а l y t i c


а

l C h e m i s t r y

C H E 1 3 0 L

EXPERIMENT NO. 3


XI.

 (    )(  )                   

REFERENCES

Christian, Gary D. 2004. Analytical chemistry (6th ed.). John Wiley and Sons Inc. Hage, David S. and James D. Carr. 2011 . Analytical chemistry and quantitative analysis. New Jersey: Pearson Prentice Hall. th

Skoog, Douglas et. al. 2004. Fundamentals of Analytical Chemistry (8 ed.). Singapore: Thomson Learning.


Page | 17

Experiment 3 - Acid and Base Titration

Recommend Documents