COLOR CHANGES AND pH RANGE OF SOME INDICATORS

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Construction of a titration curve (Neutralization curve)

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The mechanism of neutralization process can be understood by studying the changes in the hydrogen ion concentration during the course of the approximate titration. The change in pH in the vicinity of the equivalence point is of greatest importance as it enables an indicator to be selected with smallest titration error.

The curve (pH vs the moles of base or acid added) may be evaluated experimentally by
determination of the pH at various stages during the titration by a potentiometric method or by usage of universal indicator.

Universal indicator is a mixture of suitable indicators which gives a color change over a considerable portion of pH range. They are not suitable for quantitative titrations, but may be employed for the determination of approximate pH of a solution.


Procedure


Pipette out 25.00 mL of the solution A into a titration flask and add few drops of the universal indicator. Compare the color with the color chart and record the corresponding pH value of the solution. Fill the burette with the solution B. Add following amounts (X) of solution B from the burette to the flask and record the corresponding pH value after each addition by comparing the color charts provided.

X mL = 12.50, 18.50, 22.50, 24.50, 24.75, 24.80, 24.85, 24.90, 24.95, 25.00, 25.05 25.10,25.15, 25.20, 25.25, 25.30, 25.35, 25.40, 25.45, 25.50, 27.50, 31.50 37.50, 42.50, 50.00

(i) Construct the titration curve using your results
(ii) Thus obtain the ideal pH range for an indicator
(iii) Suggest possible indicators which could be used in this titration



NOTE:-
You can use a pH meter instead of the universal indicator. Follow the same procedure usingthe pH meter and construct the titration curve. You can find the equivalence point in this curve which is the exact pH (a point) the neutralization takes place. When you use the indicators it is a pH range what you get, not the exact point.

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Titration of a strong base with a strong acid(Determination of the concentration of sodium hydroxide using 0.10 M HCl)

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Theory

All acid base titrations basically involve the following reaction

As the titration proceeds, initially the pH drops slowly, but nearer the end point, the pH of the solution drops rapidly. i.e. in the vicinity of the equivalence point the rate of change of pH of the solution is very rapid.

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Procedure

1. Prepare a 250.00 mL solution of 0.10M HCl using the concentrated HCl acid provided.
2. Pipette out 25.00 mL of the given NaOH solution in to the titration flask, add few drops of phenolphthalein indicator. Titrate this solution with the HCl solution you prepared. Repeat the titration using methyl orange as the indicator. 
3. Obtain two readings in each.
4. Calculate the concentration of NaOH solution using your results.

The concentration of NaOH calculated when methyl orange is used as the indicator
will be slightly higher than the concentration when phenolphthalein is used.
Phenolphthalein is a weak base which has a pH range 8.0-10.0. Methyl orange is a
weak acid which have the pH range 3.0-5.0. More HCl is consumed when methyl orange
is used as the indicator because it shows the acidic color when the pH of the
medium fall to pH 3.

Total free energy for above neutralization reaction (Room temperature 30o C)

Δ G = - RT ln K
= 8.314Jmol-1 K-1 × 303K × 2.303 × log 10 ^14
= -8.1222 × 1014 J
Where Δ G = Free energy change
R= Universal gas constant
T= Absolute temperature
K= Equilibrium constant

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Titration of a weak acid with a strong base key words: dissociation of acetic acid

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Theory

Dissociation of acetic acid can be given as below

For dilute solutions

The pH of the solution at the equivalence point is
pH = 1/2 pKw + 1/2 pKa - 1/2 pC

Where
= the ionic product of water = 1× 10^14
= the dissociation constant of acetic acid = 1.82 ×10^-5
= the concentration of the salt in moldm-3
Accordingly, for the neutralization of 0.1M acetic acid with 0.1M NaOH at the equivalence point,


pH = 7 + 2.37 - 1/2 (1.3) = 8.37

The initial pH of the 0.1M acetic acid can be calculated using its Ka value to be 2.87.
For other concentrations corresponding pH value can be calculated by,

 

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Procedure

Pipette out 25.00mL of the given acetic acid solution into a titration flask and add few drops of thymolphthalein indicator. Titrate this solution with 0.10M NaOH solution.

1. Calculate the concentration of the acetic acid solution.
2. Sketch the titration curve.
3. Comment on choosing indicators for this titration.
4. Derive an equation for the effective or conditional formation constant of water on the employment of a weak acid, HA, with an acid dissociation constant (Ka)
5. Calculate the conditional formation constant of water for the above titration.
6. Calculate the standard free energy change of this titration (Room temperature = 30 oC
7. A sample of vinegar, weighing 12.0g, was titrated with a 0.500M NaOH solution. 17.50 mL being required to obtain a phenolphthalein end point and 5.5 mL being required to obtain a methyl orange end point. Explain the use of correct indicator system in this titration and calculate the percentage of acetic acid in vinegar.

Discussion

Acetic acid is a weak acid and NaOH is a strong base. When acetic acid is titrated with NaOH, with the final drop of NaOH at the end point, the medium turns alkaline. So an alkaline indicator is suitable for this titration. Thymolphthalein is an alkaline indicator with the pH range of 9.3-10.5, the pH range where this reaction has its end point.
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Titration of a weak base with strong acid (Key words: Dissociation constant of ammonia)

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Determination of the concentration of ammonia using 0.10M HCl

Theory

The pH of the solution at the equivalence point is,

pH = 1/2 pKw -1/2 pKb + 1/2pC


Where,
Kb = the dissociation constant of ammonia = 1.85 x 10^-5

At the equivalence point of the neutralization of 0.1M aqueous ammonia with 0.1M HCl,

pH 7 – 2.87 + ½(1.3) = 5.28

For other concentrations, the pH may be calculated from,

 

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Procedure

Pipette out 25.00 mL of the given aqueous ammonia solution into a titration flask, add few drops of methyl orange indicator. Titrate this solution with the 0.100M HCl acid you have prepared.

1. Calculate the concentration of the aqueous ammonia solution.
2. Draw the titration curve
3. Comment on the selection of a suitable indicator for this reaction.
4. Calculate the pH of the solution at the equivalence point if 100.00 mL of 0.100M methylamine (pKa = 10.64) solution is titrated with 0.100M HCl acid.
5. Calculate the standard free energy change for the above titration at room temperature.

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Neutralization of a weak acid with a weak base

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Theory

For the following reaction, pH of the solution at the equivalence point is,

For the titration of 0.1M acetic acid with 0.1m aqueous ammonia the pH at the equivalence point is 7.1

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 There is no sudden pH change and hence no sharp end point can be formed with any simple indicator.

• Calculate the standard free energy change for the neutralization of 0.1M acetic acid with 0.1M aqueous ammonia.
• Compare this value with the values you obtained for the standard free energy change in previous experiments and comment on the feasibility of this titration.
• Discuss the feasibility of titration of a solution of formic acid pKa = 3.75) with a solution of piperidine (pKb = 2.80)

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Neutralization of a polyprotic acid with a strong base (Key words: Phosphoric acid

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Determination of the concentration of phosphoric acid using NaOH

Theory

The shape of the titration curve will depend on the relative magnitudes of the various dissociation constants. If a dibasic acid H2A has two equivalent independent ionization processes with two dissociation constants K1 & K2, the ratio K1/K2 should be greater than 104 (K1/K2 > 104) in order to titrate the two hydrogen ions separately.
i.e. the acid behaves as a mixture of two acids.
The common polyfunctional acid used for titration is phosphoric (V) acid (orthophosphoric acid). The stepwise dissociation of the acid is represented by the following set of equilibria

Hence the acid will behave as a mixture of three monoprotic acids. The neutralization proceeds almost completely to the end of the primary stage before the secondary stage is affected, and the secondary stage proceeds almost to completion before the tertiary stage apparent.

The pH of the equivalence point is,
pH = 1/2pK1 + 1/2 pK2 = 4.6
At the second equivalence point,

pH = 1/2pK2 + 1/2pK3

and at the third equivalence point,

pH = 1/2pKw + 1/2 pK3 - 1/3 pC

for 0.1M phosphoric acid, pH at the third equivalence point = 12.35

In the third ionization stage, HPO42- behaves as an exceedingly weak acid and the titration curve is very flat (the magnitude of the free energy change is almost zero) and no indicator is available for direct titration. Therefore phosphoric acid behaves identical to a dibasic acid.

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Procedure

Pipette out 25.00 mL of the phosphoric acid into a titration flask and add few drops of bromocresol green indicator and titrate with the 0.20M NaOH solution.
Repeat the titration using thymolphthalein indicator.

I. Calculate the concentration of phosphoric acid.
II. Draw the titration curve and indicate the use of correct indicators in the titration.
III. Suggest an experimental procedure that you would carry out to determine the dissociation constant of the acid. Carry out the suggested procedure in consultation with your demonstrator. Thus determine the concentration of the acid.
IV. Discuss the possibility of using citric acid as a triprotic acid. The dissociation constants of the acid are:

                  K1 = 9.2 x 10^-4, K2 = 2.7 x 10^-5, K3 = 1.3 x 10^-6

V. You are provided with a mixture of two acids H2A and HB the corresponding dissociation constants are given below.
For H2A K1 = 2 x 10^-2, K2 = 1 x 10^-13
For HB K1 = 3.2 x 10^-7

• Is it possible to determine the concentration of the two acids in the mixture by titrating with a standard solution of NaOH? Explain your answer giving reasons and calculations.
• Draw the appropriate titration curve indicating all the necessary features.
• Thus determine the concentration of individual acids in the mixture if the concentration of NaOH is 0.01M.

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