neutralization titrations

Experiment No. ___________________

Date ___________________

NEUTRALIZATION TITRATIONS

INTRODUCTION
The neutralization of hydronium or hydroxide ion to form water is widely used as the basis for volumetric determinations of acids, bases and salts of weak acids. The reaction is characterized by a rapid change in pH near the equivalence point, a change that is readily detected by an acid-base indicator or that can be followed electrically by use of a pH meter. Neutralization titrations are performed with standard solutions of strong acids or strong bases. A standard solution (or standard titrant) is a reagent of exactly known concentration.
Standard solutions play a central role in all volumetric methods of analysis. The ideal standard solution for a volumetric method should
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be sufficiently stable so that it will not be necessary to determine its concentration (to standardize it) so frequently. react rapidly with the analyte so that the time required between additions of titrant is minimized. react more or less completely with the analyte so that satisfactory end points are realized. undergo a selective reaction with the analyte that can be described by a simple balanced equation. The accuracy of a volumetric method is closely related with the accuracy of the concentration of the standard solution used in the titration. Two basic methods are used to establish the concentration of standard solutions: In the direct method, a carefully weighed quantity of a primary standard is dissolved and diluted to an exactly known volume in a volumetric flask. In the second method, the solution is standardized by titrating
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a weighed quantity of a primary standard a weighed quantity of a secondary standard or a measured volume of another standard solution.

As a result, standardization is a process of determining the concentration of a substance in solution by adding to it a standard reagent of known concentration in carefully measured amounts until a reaction of definite and known proportion is completed, as shown by a color change, and then calculating the unknown concentration.
A titrant that is standardized against a secondary standard or against another standard solution is sometimes referred to as a secondary solution. A secondary standard is less desirable than a primary standard solution, because the concentration of the former is subject to greater uncertainty. A primary standard is a highly purified compound that serves as a reference material in all titrimetric methods. Important requirements for a primary standard are;
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High purity
Stability in air
Absence of hydrate water so that the composition doesn't change with variations in relative humidity
Reasonable solubility in the titration medium
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Reasonably large formula weight so that the relative error associated with weighing is minimized The Effect of Carbon Dioxide on Standard Base Solutions
In the preparation of acid and base solutions atmospheric CO2 may cause some problems.
Atmospheric CO2 is in equilibrium with aqueous carbonic acid whose concentration is about
1.5x10-5 M at ordinary temperature. So this small concentration of carbonic acid leads to no significant error. On the other hand, the distilled water is sometimes supersaturated with the gas and thus contains sufficient carbonic acid to cause detectable errors. In order to test the distilled water to be used in neutralization titrations, take about 500 mL from the source of production.
Add 5 drops of phenolphthalein and titrate with ~0.1 M NaOH. Less than 0.2 mL of base should be used for the first faint pink color. If a larger volume is needed, standard solutions should be prepared from the water that has been boiled for 2 to 3 minutes to remove CO 2 and then cooled to room temperature.
CO2 (from the atmosphere) reacts with the hydroxides of sodium, potassium, and barium
(solution or solid), producing the carbonate:
CO2 (g) + 2OH − → CO2− + H2 O
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Absorption of carbon dioxide by a standardized solution of sodium or potassium hydroxide leads to a negative systematic error (called carbonate error), in analyses in which a basic range indicator is used. Here, each carbonate ion (produced from two hydroxide ions) has reacted with only one hydronium ion when the indicator changes its color:
CO2− + H3 O+ → HCO− + H2 O
3
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However, no systematic error occurs when an acidic range indicator is used. Each carbonate ion produced will have reacted with two hydronium ions of the acid:
CO2− + 2H3 O+ → H2 CO3 + 2H2 O
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The amount of hydronium ion consumed by this reaction equals the amount of hydroxide lost during formation of the carbonate ion, thus, no error occurs.
Solid reagents used to prepare standard base solutions are always contaminated by carbonate ion. However, no carbonate error occurs if the same indicator is used for both standardization and analysis. But it leads to less sharp end points. It is therefore important to remove carbonate ion before a solution of a base is standardized.
Water used to prepare carbonate-free solutions must also be free of carbon dioxide. Distilled water, which is sometimes supersaturated with carbon dioxide, should be boiled to eliminate the gas. The water is then allowed to cool to room temperature before introducing the base, because hot alkali solutions absorb carbon dioxide rapidly.
Acid-Base Indicators
In all neutralization titrations, results are dependent upon the selection of the proper indicator.
The indicators most commonly used in neutralization titrations are highly colored organic compounds which have the property of changing color when the hydrogen ion concentration of the solution is changed over a certain range. The most widely used are phenolphthalein, methyl orange, methyl red and bromocresol green.

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The hydronium ion concentrations, H3O+, at the point of color change (end point) are widely different for various indicators. Phenolphthalein is colorless in solutions in which [H3O+]≥ 10-8 M and pink in solutions in which [H3O+]≤ 10-10 M. If the hydronium ion concentration changes from
10-10 to 10-8 M, the phenolphthalein indicator will change from pink to colorless. If the solution assumes a hydronium ion concentration intermediate between these two values, the color is also intermediate between deep pink and colorless. In other words, phenolphthalein shows a color change in the basic side of neutrality. Methyl orange, methyl red and bromocresol green give color changes on the acid side of neutrality, in the region of hydrogen ion concentration
10-3 to 10-6 M.
A list of common indicators with pH range is given in the following table.
Table 1. Some Important Acid-Base Indicators
Transition
Range, pH
1.2-2.8
3.1-4.4
3.8-5.4
4.2-6.3
6.0-7.6
8.3-10.0
9.3-10.5
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Common Name
Thymol blue
Methyl orange
Bromocresol green
Methyl red
Bromothymol blue
Phenolphthalein
Tymolphthalein
Alizarin yellow GG
*

Color
Change*
R-Y
R-O
Y-B
R-Y
Y-B
C-P
C-B
C-Y

Type**
1
2
2
2
1
1
1
2

B=blue, C=colorless, O=orange, P=purple, R=red, Y= yellow
HIn + H2 O

→
+
← H3 O

+ In−

(2) base type: In− + H2 O

**

→
+
← InH

+ OH −

(1) acid type:

The proper indicator for a titration is the one that will exhibit a color change at the hydronium ion concentration found at the equivalence point of the titration.
Standard Solutions for Neutralizations Titrations
The standard solutions used for neutralization titrations are usually solutions of strong acids or bases. The most frequently used acids are hydrochloric acid and sulfuric acid, either of which is suitable for use as a permanent reference standard.
Sodium, potassium and barium hydroxides are the most frequently used bases. Bases, however, are not as good as acids for permanent standards, because they absorb carbon dioxide whenever they come in contact with air. Solutions of bases also react with glass on long storage. Base solutions can, however, be stored in pyrex or plastic bottles (such as polyethylene) if the bottle is equipped with a siphon and the inlet tube is equipped with a tube of soda lime (mixture of Ca(OH)2 and Na2CO3) to remove CO2 from the entering air.
REAGENTS AND APPARATUS
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HCl, CH3COOH (unknown solutions, already prepared)
1.0 L of 0.1 M NaOH (for 2 students)
1.0 L of 0.1 M HCl (for 2 students)*
Potassium hydrogen phthalate, KHC8H4O4 (KHP) (primary standard)
Na2CO3 (primary standard)
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Bromocresolgreen, phenolphthalein, methyl orange indicators, bromothymol blue indicator (already prepared)
Vinegar (students` responsibility is to bring a bottle of vinegar to the lab.) buret 250mL conical flasks
100 mL graduated cylinder
1 L brown bottle for 2 students
0.25 L, 0.5 L and 1 L volumetric flasks for 2 students

* Do not pour standardized HCl at the end of the experiment which will be used in the analysis of carbonate mixture experiment. Store HCl in brown colored 1 L bottle. And do not forget to note the concentration of HCl which will be used in the analysis of carbonate mixtures experiment as a standard solution.
PROCEDURE
A. Preparation and Standardization of 0.1 M HCl
1)

Prepare 1.0 L of 0.1 M HCl solution from the concentrated HCl solution for two students.

2)

Transfer 0.20 to 0.25 g (0.1 mg) of primary standard Na2CO3 to a 250 mL erlenmeyer flask and add 50 mL of distilled water and 2-3 drops of bromocresol green indicator. Make sure that Na2CO3 dissolves in water.

3)

Titrate the solution with 0.1 M HCl until the solution just changes its color from blue to green
(V1). The net reaction during titration can be written as follow:
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CO 3  2H  CO 2  H2 0 (H2CO 3 )

4)

Place the flask on a wire gauze and heat just to gentle boiling for two or three minutes to remove CO2 produced. Cool to the room temperature and complete the titration.
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

If there is CO2 in the solution, the indicator should change from green to blue as CO2 is removed during heating. When you see blue color, titrate this solution with 0.1 M HCl till green color is observed and this additional amount of HCl has to be added V1 to the to determine end point.
As shown below, two end points are observed in the titration of sodium carbonate. The first at about pH 8.3, corresponding to the conversion of carbonate to hydrogen carbonate, the second, involving the formation of carbon dioxide, is observed at about pH 3.8.

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The second end point is always used for standardization because the change in pH is greater than that of the first. An even sharper end point can be achieved by boiling the solution to eliminate carbonic acid and carbon dioxide. The sample is titrated to the first appearance of the acid color of the indicator (such as bromocresol green or methyl orange). At this point, the solution contains a large amount of dissolved carbon dioxide and small amounts of carbonic acid and unreacted hydrogen carbonate. Boiling destroys this buffer by eliminating the carbonic acid:
H2 CO3 (g) → CO2 (g) + H2 O(l)
The solution then becomes alkaline again due to the residual hydrogen carbonate ion.
5)

Repeat the titration with one additional sample.

6)

Your partner will do the standardization experiment twice. Record the results of all titrations
(totally 4 titrations) and calculate the molarity of the HCl solution for four replicates and at the end calculate the average of molarity of HCl solution using the following equation. g Na 2CO 3
V
xM
2
HCl
HCl
gfw Na 2CO 3

VHCl is the volume of HCl (in L) used.
MHCl is the molarity of HCl solution. g Na2CO3 is the weight of Na2CO3 taken. gfw Na2CO3 is the molar mass of Na2CO3.
B. Preparation and Standardization of 0.1 M NaOH
1)

Boil 500 mL of distilled water and then cool it. Use this distilled water to prepare 500 mL of
0.1 M NaOH solution for two students.

2)

Transfer 0.10 to 0.20 g (1 mg) of potassium hydrogen phthalate (KHC8H4O4) into a 250 mL conical flask. Dissolve the salt in 50 mL of distilled water and add 3 to 5 drops of phenolphthalein indicator.

3)

Titrate the solution with 0.1 M NaOH prepared until a faint pink color appears and persists at least 20 seconds. The net reaction during titration can be written as follow:
HP   OH  H2O  P2

Where, HP- is hydrogen phthalate ion coming from KHP, OH- is hydroxide ion coming from
NaOH and P2- is phthalate ion.
4)

Repeat the titration with one additional sample.

5)

Your partner will do the standardization experiment twice. Record the results of all titrations
(totally 4 titrations) and calculate the molarity of the NaOH solution for four replicates and at the end calculate the average molarity of NaOH solution using the following equation: g KHP
V
xM

NaOH
NaOH gfw KHP

VNaOH is the volume of NaOH (in L) used.
MNaOH is the molarity of NaOH solution.
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g KHP is the weight of KHP taken. gfw KHP is the molar mass of KHP.
C. Analysis of Weak Acids
1)

Take two replicate unknown sample from your assistant

2)

Treat each unknown solution individually.

3)

Add 50 mL of distilled water and 2 drops of phenolphthalein and then titrate with standardized NaOH solution to the first permanent (~30 s) pink color. The net reaction during the titration is
CH3COOH  OH  CH3COO   H2O

4)

Calculate the total acidity as mg CH3COOH for each unknown sample separately and then calculate the average of the results.

VNaOH x MNaOH 

g CH3COOH gfw CH3COOH

D. Analysis of HCl and CH3COOH Mixture
1)

Take two replicate unknown sample from your assistant.

2)

Treat each unknown solution individually.

3)

Add 50 mL of distilled water and 2 drops of methyl orange indicator and then, titrate with standard NaOH solution until color changes.

4)

Now add 2 drops of phenolphthalein indicator and titrate until the appearance of pink color.

5)

Determine the amount of HCl from the volume and molarity of NaOH used to obtain first end point (methyl orange end point).

6)

Determine the amount of CH3COOH from the volume and molarity of NaOH to get phenolphthalein end point.

E. Determination of Acetic Acid in Commercial Vinegar
1) Pipette exactly 2.00 mL of the commercial vinegar sample into a 250 mL Erlenmeyer flask and add 50 mL of distilled water.
2) Add 3 drops of phenolphthalein indicator and titrate the acetic acid with the standard NaOH solution to a pale pink equivalence point. Record the burette readings.
3) Repeat the titration twice more using a fresh aliquot of vinegar. Results should agree within ±0.2 mL or additional titrations are required.
4) Use the newly calculated molarity of the NaOH solution and calculate the moles of acetic acid titrated and the % (w/v) of acetic acid in the vinegar.
PRE-LAB STUDIES
1)

Read pages 338-349, 368-383, 428-435 from the textbook (8th Ed)
Read pages 303-314, 322-337, 382-387 from the textbook (9th Ed)

2)

Define the followings:
Standardization, titrant, primary standard, acid/base indicators, indicator range, end point, equivalance point, titration curve, buffer solution.

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3)
4)

What are the requirements for a primary standard?
Write the transition range (pH) and color change for the following indicators:
Phenolphthalein, bromocresol green, methyl orange.
5) What is the aim of standardization?
6) What is the difference between primary standard and standard solution?
7) Define and compare end-point and equivalence point.
8) Why the weak acids cannot be used as a titrant in the titrations?
9) Why do the indicators change their colors for different pH’s. (Hint: Consider the chemical behavior of the indicators.)
10) Describe the preparation of 1.0 L of 0.1 M HCl from concentrated HCl.*
11) Describe the preparation of 1.0 L of 0.1 M NaOH from solid NaOH.
* Write the density and weight percent values from the concentrated HCl solution bottle in the lab. POST-LAB STUDIES
1) Explain the effect, if any, of each of the following sources of error upon the molarity of the base as determined in the experiment; i.e., would the experimental value for molarity be too high or too low, and why?
a) If the beaker in which the titration was performed contained several mL of distilled water from the rinsing at the time the KHP was massed into it.
b) If the tip of the buret was not filled with solution before the initial reading was taken.
c) If a bubble appeared in the tip during the titration.
d) If liquid splashed from the titration beaker before the end point had been reached.
e) If the buret was not rinsed with the NaOH solution following the rinsing with distilled water. 2) Suppose that, instead of using NaOH, a base such as Ba(OH)2 had been used. What changes in the calculations would then have to be made to determine the molar concentrations of the base? Answer this question in words, and illustrate you answer by calculating Molarity from the following data: mol KHP used 0.040 mol
Ba(OH)2 Initial buret level = 0.020 mL
Final buret level = 36.70 mL
3) Write the name of primary standards used for the standardization of HCl and NaOH? Could you use them interchangeably? Why should we pay attention to measure the mass of the primary standard precisely to standardize standardization of NaOH and HCl?
4) Why do we boil and then cool distilled water before the preparation of NaOH? Why this process is not necessary for the preparation of HCl?
5) In the titration of weak acid with a strong base, why did you use phenolphthalein as indicator? 6) In the titration of weak and strong acid mixture with a strong base, which acid is titrated first?
Write the related titration reactions by drawing the corresponding titration curve.
7) For the titration of a strong acid, a neutral range indicator is used and for the titration of weak acids a basic range indicator is used. However, for the titration of weak and strong acid mixture you used an indicator of acidic range for the first end-point. Why?
8) Write the related reactions and draw the titration curves for the titration of CH3COOH and
CH3COOH&HCl mixture with NaOH.

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Name surname:

Section:

Date:

REPORT SHEET FOR NEUTRALIZATION TITRATIONS
A. Standardization of HCl using Na2CO3
Na2CO3 Replicate
Number
1
2
3
4

Na2CO3 (g)

Volume of 0.1 M
HCl (mL)

M HCl

̅?), M
Average (?
B. Standardization of NaOH using KHP
KHP Replicate
Number
1
2
3
4

KHP (g)

Volume of 0.1 M
NaOH (mL)

M NaOH

̅?), M
Average (?
C. Analysis of Weak Acids
Replicate Number

Volume of 0.1 M NaOH (mL) phenolphthalein end point

1
2
̅?) mg=
Average mg CH3COOH, (?
True value, mg CH3COOH=
D. Analysis of HCl and CH3COOH Mixture
Replicate
Number
1
2

Volume of 0.1 M HCl (mL) methyl orange end point

Volume of 0.1 M HCl (mL) phenolphthalein end point

̅?) mg=
Average mg HCl, (?
̅?) mg=
Average mg CH3COOH, (?
True value, mg HCl =
True value, mg CH3COOH=
E. Determination of Acetic Acid in Commercial Vinegar
Replicate

Volume of 0.1 M NaOH (mL) phenolphthalein end point

1
2
3
TA`s Name and Signature:

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