Acid-base titration - Pharmaceutical Analysis 1 B. Pharma 1st semester

Acid-base titration

Content

• Different types of acid base titration

• Theories of indicators

   Ostwald's theory

   Quinonoid theory

•The different types of neutralization curve

   Strong acid/strong base

   Strong base/weak acid

   Strong acid/weak base

   Weak acid/weak base

Learning Objectives

At the end of this lecture, the student will be able to

• Explain the different types of acid base titration

• Explain the theories of indicators

•Explain the different types of neutralization curves

•Discuss about the colour change and pH range of various acid base indicators

Acids

• Acids are sour tasting

• Blue litmus paper turns blue

• Arrhenius acid: Any substance that, when dissolved in water, increases the concentration of hydronium ion (H3O+)

Bronsted Lowry acid: A proton donor

• Lewis acid: An electron acceptor

Base

• Bases are bitter tasting and slippery

• Red litmus turns to blue

• Arrhenius base: Any substance that, when dissolved in water, increases the concentration of hydroxide ion (OH-)

Bronsted Lowery base: A proton acceptor

Lewis acid: An electron donor

What Happens During Titration of Acid Base Titration

Different types of acid base titration:

• Acidimetry

• Alkalimetry

• Displacement titration

• Nonaqueous titration

Acidimetry: Titration of a base with a standard acid

Example: Sodium hydroxide with Hydrochloric acid

NaOH + HCl → NaCl+H2O

Alkalimetry: Titration of an acid with a standard based

Example: Hydrochloric acid with Sodium hydroxide

HCl+ NaOH → NaCl+H2O

Displacement titration: Titration which involves the displacement of a weaker anion and a cation by a stronger anion and cation

• They are of two types:

Titration of weaker anion by strong acid

Na2 B4O7 + H2SO4→ Na2SO4+ H3BO3

Titration of weaker cation with a strong base

NH4Cl+ NaOH→NH4OH+ NaCl

Non aqueous titration: Titration carried out in non-aqueous media for determining the strength of weak acid or weak base

• They are of two types: Acidimetry, Alkalimetry

Choosing a Titrant

In theory, any strong acid or strong base can be used as titrant

The reason for this is that most reaction involving a strong acid or a strong base is QUANTITATIVE

Strong base titrant

Weak base titrant

Strong acid titrant

Sodium hydroxide, NaOH

Potassium hydroxide, KOH

Magnesium hydroxide,Mg(OH)2

Barium hydroxide, Ba(OH)2

Ammonium hydroxide, NH4OH Amine acetate Carbonate, CO32-

Flouride ion, F-

Sodium carbonate, Na2CO3

Hydrochloric acid, HCl

Sulphuric acid, H2SO4

Perchloric acid, HClO4

 

pH – A Measure of Acidity

pH = -log [H+]

Solution

 

At 250C

 

neutral

[H+] = [OH-]

[H+] = 1 x 10-7

pH = 7

acidic

[H+] > [OH-]

[H+] > 1 x 10-7

pH < 7

basic

[H+] < [OH-]

[H+] < 1 x 10-7

pH > 7

Indicator and Choice of Indicator

Acid –Base Indicators

• The acid-base indicator function by changing colour just after the equivalence point of a titration; this colour change is called the end point

• The end point is most often detected VISUALLY.  Most acid-base indicators are organic dye molecules which are either acids or bases.

• Indicators can be monoprotic (HIn) or diprotic (H2In) acids.

• The acid form of an indicator is usually coloured; when it loses a proton resulting in anion (In-), or base form of the indicator, exhibiting different colour

The two forms exist in equilibrium with one another as follows:

Hin ---------------------à H+   +  In-

(colour A)                                               (colour B)

A diprotic acid indicator, H2In, ionizes in TWO steps as follows:

H2In ---------------------à HIn-   + H+ ---------------------à H+   +  In-

(colour A)                                      (colour B)                                                                (colour C)

Selecting and Using Indicators

• Choose an indicator that undergoes a distinct colour change at the equivalence point (true end point) of an acid-base titration

• Choose an indicator whose middle half of the pH transition range (greatest colour change) encompasses/overlap the pH at the equivalence point or the pH at the steepest part of the titration curve.

• For indicator color change to be detectable by normal eyesight, a rough tenfold excess of one or the other form of the   indicator must be present. This corresponds to a pH range of + 1 pH unit about the pK of the indicator

Acid-Base Indicators

Common Name

 

Transition range

 

Colour Change

ACID

Colour Change

BASE

Crystal violet

0.1 – 1.5

Yellow

Blue

Thymol blue

1.2 – 2.8

Red

Yellow

Mrthyl yellow

2.4 – 4.0

Red

Yellow

Methyl orange

3.1 – 4.4

Red

Yellow

Bromocresol green

3.8 – 5.4

Yellow

Blue

Methyl red

4.2 – 6.3

Red

Yellow

Chlorophenyl red

4.5 – 6.4

Yellow

Red

Bromothymol blue

6.0 – 7.6

Yellow

Blue

Phenol red

6.4 – 8.0

Yellow

Red

Thymol blue

8.0 – 9.6

Yellow

Blue

Phenolpthalein

8.0 – 9.7

Colourless

Red

Alizarin yellow

10.0 – 12.0

colourless

Violet

 

Properties of an indicator:

Must have an easily observed colour change

Must change immediately in the required pH range over the addition of ‘half’ a drop of reagent

To be useful, an indicator must change over the “vertical” section of the curve where there is a large change in pH for the addition of a very small volume of alkali.

The indicator used depends on the pH changes around the end point - the indicator must change during the ‘vertical’ portion of the curve.

 In the example, the only suitable indicator is PHENOLPHTHALEIN

Theory of indicator

• An acid/base indicator is a weak organic acid or a weak organic base whose undissociated form differs in colour from its conjugate base or conjugate acid form

Examples: Methyl orange, methyl red, ---etc

Theory of indicator is explained based on:

 Ostwald's theory

 Quinonoid theory

Based on Ostwald's theory:

HIn + H2O ó  In- + H3O+

or

In + H2O   ó   InH+ + OH-

• The behaviour of a base type indicator is described by the equilibrium

• The equilibrium constant takes the form:

[H3O+][In-] / [HIn]  = Ka

Rearranging:

[H3O+] = Ka  [HIn-]/[In-]

– HIn pure acid color: [HIn]/[In-] ≥ 10

– HIn pure base color: [HIn]/[In-] ≤ 0.1

~The ratios change from indicator to indicator~

– Substitute the ratios into the rearranged Ka: [H3O+] = 10Ka  (acid color), [H3O+] = 0.1Ka (base color)

– pH range for indicator = pKa ± 1

• acid color pH = -log(10Ka) = pKa + 1

• base color pH = -log(0.1Ka) = pKa – 1

Quinonoid theory:

According to this theory

(a) The acid-base indicators exist in two tautomeric forms having different structures. Two forms are in equilibrium

One form is termed benzenoid form and the other quinonoid form.

(b) The two forms have different colors. The color change in due to the interconversation of one tautomeric form into other

(c) One form mainly exists in acidic medium and the other in alkaline medium

• Thus, during titration the medium changes from acidic to alkaline   or   vice-versa.   The   change   in   pH   converts   one tautomeric form into other and thus, the colour change occurs.

Example for Phenolphthalein:

• Has benziod form in acidic medium and thus, it is colourless

• While it has quinonoid form in alkaline medium which has pink colour.

Methyl orange solution: has  quinonoid  form  in  acidic  solution and benzenoid form in alkaline solution The color of benzenoid form is yellow while that of quinoniod form is red

Mixed Indicators

• With ordinary acid base indicators, the color change is not sharp and abrupt, but it extends over 2 units of pH

• However, sometimes, it is desirable to have sharp color change over a narrow and selected range of pH

• This can be achieved by the use of mixture of indicators whose pKind values are close together and the overlapping   colors are complementary at an intermediate pH value

For example:

A mixture of 3 parts of phenolphthalein and 1 part of naphtholphthalein shows a change in color from pale rose to violet at pH = 8.9. --- Useful for the titration of phosphoric acid to diproti stage

Similarly, a mixture of 3 parts of thymol blue and 1 part of cresol red shows change in color from yellow to violet at pH = 8.3 --- useful for the titration of CO32- to HCO3-  stage

Neutralization Curve pH Curves

Types: There are four types of acid-base titration; each has a characteristic curve.

strong acid (HCl)  v.  strong base (NaOH)

weak acid (CH3COOH)  v.  strong alkali (NaOH)

strong acid (HCl)  v.  weak base (NH3)

weak acid (CH3COOH)  v.  weak base (NH3)

In the following examples, alkali (0.1M) is added to 25cm3 of acid (0.1M)

End points need not be “neutral‘ due to the phenomenon of salt hydrolysis

Strong Acid (HCl) v. Strong Base (NaOH)

Any of the indicators listed will be suitable - they all change in the ‘vertical’ portion

Strong Acid (HCl) v. Weak Base (NH3)

Only methyl orange is suitable - it is the only one to change in the ‘vertical’ portion

Weak Acid (CH3COOH) v. Strong Base (NaOH)

Only phenolphthalein is suitable - it is the only one to change in the ‘vertical’ portion

Weak Acid (CH3COOH) v. Weak Base (NH3)

There is no suitable indicator- none change in the ‘vertical’ portion.

The end point can be detected by plotting a curve using a pH meter.

Other pH curves - Acid v. Carbonate

Sodium carbonate reacts with hydrochloric acid in two steps...

Step 1        Na2CO3    + HCl  ——> NaHCO3    + NaCl

Step 2        NaHCO3   + HCl  ——> NaCl  + H2O  + CO2

Overall      Na2CO3    + 2HCl ——> 2NaCl  + H2O  + CO2

There are two sharp pH changes

The second addition of HCl is exactly the same as the first because the number of moles of HCl which react with the NaHCO3 is the same as that reacting with the Na2CO3.

First rapid pH change around pH = 8.5 due to the formation of NaHCO3 Can be detected using phenolphthalein

Second rapid pH change around pH = 4 due to the formation of acidic CO2 Can   be   detected   using   methyl orange

Other pH curves - polyprotic acids (H3PO4)

Phosphoric acid is triprotic; it reacts with sodium hydroxide in three steps...

Step 1       H3PO4      +  NaOH     ——>   NaH2PO4      +  H2O

Step 2       NaH2PO4   + NaOH   ——>   Na2HPO4   +  H2O

Step 3       Na2HPO4   + NaOH   ——>   Na3PO4    +  H2O

There are three sharp pH changes

Each successive addition of NaOH is the same as equal number of moles are involved

Sodium Bi Carbonate

Assay Principle

Acid base, Acidic metric titration

NaHCO3 + H2SO4 ==> Na2SO4 + H2O + CO2

Indicator: Methyl orange

Colour change: yellow to Red

Boric acid

Assay Principle

Neutralization titration, indirect titration alkali metric

Indicator: phenolphthalein

Reagent: Glycerine

Colour change: colourless to pale pink

Summary

• Classification of acid base titration:

Acidimetry

Alkalimetry

Displacement titration

Non aqueous titration

Theory of indicators: Ostwalds theory, Quinonoid theory

   Neutralization Curves:

Strong acid vs strong base

Strong acid vs weak base

Strong base vs weak base

Weak acid vs weak base

• Indicators used for different types of acid base titrations

Sodium bi carbonate: Acid base Acidic metric titration

Boric acid: Neutralization titration, indirect titration alkali metric

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