What is Acid-Base Indicators and Titration

Introducing :Acid - Base Indicators

The most common method to get an idea about the pH of solution is to use an acid base indicator. An indicator is a large organic molecule that works somewhat like a " color dye". Whereas most dyes do not change color with the amount of acid or base present, there are many molecules, known as acid - base indicators , which do respond to a change in the hydrogen ion concentration. Most of the indicators are themselves weak acids.

The most common indicator is found on "litmus" paper. It is red below pH 4.5 and blue above pH 8.2. 

Color	Blue Litmus	Red Litmus
Acid	turns red	stays same
Base	stays same	turns blue

Other commercial pH papers are able to give colors for every main pH unit. Universal Indicator, which is a solution of a mixture of indicators is able to also provide a full range of colors for the pH scale.

A variety of indicators change color at various pH levels. A properly selected acid-base indicator can be used to visually "indicate" the approximate pH of a sample. An indicator is usually some weak organic acid or base dye that changes colors at definite pH values. The weak acid form (HIn) will have one color and the weak acid negative ion (In^-) will have a different color. The weak acid equilibrium is:

HIn --> H+ + In^-
For phenolphthalein: pH 8.2 = colorless; pH 10 = red

For bromophenol blue: pH 3 = yellow; pH 4.6 = blue

See the graphic for more indicators, colors, and pH ranges.

Explain the color indicator change:

Use equilibrium principles to explain the color change for phenolphthalein at the end of the demonstration.

Solution:

The simplified reaction is: H⁺ + OH^- --> HOH

As OH^- ions are added, they are consumed by the excess of acid already in the beaker as expressed in the above equation. The hydroxide ions keep decreasing and the hydrogen ions increase, pH decreases.

See lower equation: The indicator equilibrium shifts left, In^- ions decrease. Below pH 8.2 the indicator is colorless. As H⁺ ions are further increased and pH decreases to pH 4-5, the indicator equilibrium is effected and changes to the colorless HIn form.

Equilibrium: HIn --> H⁺ + In^-
colorless red

Molecular Basis for the Indicator Color Change:

Color changes in molecules can be caused by changes in electron confinement. More confinement makes the light absorbed more blue, and less makes it more red.

How are electrons confined in phenolphthalein? There are three benzene rings in the molecule. Every atom involved in a double bond has a p orbital which can overlap side-to-side with similar atoms next to it. The overlap creates a 'pi bond' which allows the electrons in the p orbital to be found on either bonded atom. These electrons can spread like a cloud over any region of the molecule that is flat and has alternating double and single bonds. Each of the benzene rings is such a system.

See the far left graphic - The carbon atom at the center (adjacent to the yellow circled red oxygen atom) doesn't have a p-orbital available for pi-bonding, and it confines the pi electrons to the rings. The molecule absorbs in the ultraviolet, and this form of phenolphthalein is colorless.

In basic solution, the molecule loses one hydrogen ion. Almost instantly, the five-sided ring in the center opens and the electronic structure around the center carbon changes (yellow circled atoms) to a double bond which now does contain pi electrons. The pi electrons are no longer confined separately to the three benzene rings, but because of the change in geometry around the yellow circled atoms, the whole molecule is now flat and electrons are free to move within the entire molecule. The result of all of these changes is the change in color to pink.

Naturally occurring pH indicators[edit]

Many plants or plant parts contain chemicals from the naturally colored anthocyanin family of compounds. They are red in acidic solutions and blue in basic. Anthocyanins can be extracted with water or other solvents from a multitude of colored plants or plant parts, including from leaves (red cabbage); flowers (geranium, poppy, or rose petals); berries (blueberries, blackcurrant); and stems (rhubarb). Extracting anthocyanins from household plants, especially red cabbage, to form a crude pH indicator is a popular introductory chemistry demonstration.

Litmus, used by alchemists in the Middle Ages and still readily available, is a naturally occurring pH indicator made from a mixture of lichen species, particularly Roccella tinctoria. The word litmus is literally from 'colored moss' in Old Norse (see Litr). The color changes between red in acid solutions and blue in alkalis. The term 'litmus test' has become a widely used metaphor for any test that purports to distinguish authoritatively between alternatives.

Hydrangea macrophylla flowers can change color depending on soil acidity. In acid soils, chemical reactions occur in the soil that make aluminium available to these plants, turning the flowers blue. In alkaline soils, these reactions cannot occur and therefore aluminium is not taken up by the plant. As a result, the flowers remain pink.

Indicator	Low pH color	High pH color
Hydrangea flowers	Blue	pink to purple
Anthocyanins	Red	blue
Litmus	Red	blue

Indicators

Indicators are substances whose solutions change color due to changes in pH. These are called acid-base indicators. They are usually weak acids or bases, but their conjugate base or acid forms have different colors due to differences in their absorption spectra.

Do you know that the color of hydrangea flower depends on the pH of the soil in which it is grown? This picture shows various colors of hydrangea flowers.

Indicators are complicate organic weak acids or bases with complicated structures. For simplicity, we represent a general indicator by the formula HIn, and its ionization in a solution by the equilibrium,

HIn = H⁺ + In^-,

and define the equilibrium constant as K_ai,

        [H⁺][In^-]

  K_ai = ----------.

         [HIn]

Which can be rearranged to give

    [In^-]     K_ai

   ------- = -----

    [HIn]     [H⁺]

When [H⁺] is greater than 10 K_ai, In^- color dominates, whereas color due to HIn dominates if [H⁺] < K_ai / 10. The above equation indicates that the color change is the most sensitive when [H⁺] = K_ai in numerical value.

We define pK_ai = - log(K_ai), and the pK_ai value is also the pH value at which the color of the indicator is most sensitive to pH changes.

Taking the negative log of K_ai gives,

                           [In^-]

  -log K_ai = -log[H⁺] - log------

                           [HIn]

or

                 [In^-]  

  pH = pK_ai + log-----

                 [HIn]

This is a very important formula, and its derivation is very simple. Start from the definition of the equilibrium constant K, you can easily derive it. Note that pH = pK_ai when [In^-] = [HIn]. In other words, when the pH is the same as pK_ai, there are equal amounts of acid and base forms. When the two forms have equal concentration, the color change is most noticeable.

Colors of substances make the world a wonderful place. Because of the colors and structures, flowers, plants, animals, and minerals show their unique characters.

Many indicators are extracted from plants. For example, red cabbage juice and tea pigments show different colors when the pH is different. The color of tea darkens in a basic solution, but the color becomes lighter when lemon juice is put into a tea. Red cabbage juice turns blue in a basic solution, but it shows a distinct red color in an acidic solution.

Some Common Indicators

Name	Acid Color	pH Range of Color Change	Base Color
Methyl violet	Yellow	0.0 - 1.6	Blue
Thymol blue	Red	1.2 - 2.8	Yellow
Methyl orange	Red	3.2 - 4.4	Yellow
Bromocresol green	Yellow	3.8 - 5.4	Blue
Methyl red	Red	4.8 - 6.0	Yellow
Litmus	Red	5.0 - 8.0	Blue
Bromothymol blue	Yellow	6.0 - 7.6	Blue
Thymol blue	Yellow	8.0 - 9.6	Blue
Phenolphthalein	Colorless	8.2 - 10.0	Pink
Thymolphthalein	Colorless	9.4 - 10.6	Blue
Alizarin yellow R	Yellow	10.1 - 12.0	Red

Some common indicators and their pK_ai (also referred to as pK_a) values are given in a table form. Since the table is an HTML file, we can not include the table in the DOS version, but the HTML version allows you to see this table below:

There is a separate file for this, and it can also be accessed from the Chemical Handbook menu.

What is Titration?

The word titration comes from the Latin word "titulus", which means inscription or title. The French word title means rank. Therefore, Titration means the determination of concentration or rank of a solution with respect to water with a pH of 7. 

The standard solution is usually added from a graduated vessel called a burette. The process of adding standard solution until the reaction is just complete is termed as titration and the substance to be determined is said to be titrated.

All chemical reactions cannot be considered as titrations. A reaction can serve as a basis of a titration procedure only if the following conditions are satisfied:

1.    The reaction must be a fast one.

2.    It must proceed stoichiometrically.

3.    The change in free energy (ΔG) during the reaction must be sufficiently large for spontaneity of the reaction.

4.    There should be a way to detect the completion of the reaction.

End point and Equivalent point:

For a reaction, a stage which shows the completion of a particular reaction is known as end point. Equivalence point is a stage in which the amount of reagent added is exactly and stoichiometrically equivalent to the amount of the reacting substance in the titrated solution. The end point is detected by some physical change produced by the solution, by itself or more usually by the addition of an auxiliary reagent known as an 'indicator'. 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. In many cases, the difference between these two will fall within the experimental error.
 

Indicator:

It is a chemical reagent used to recognize the attainment of end point in a titration. After the reaction between the substance and the standard solution is complete, the indicator should give a clear colour change.

When a titration is carried out, the free energy change for the reaction is always negative.
  That is, during the initial stages of the reaction between A & B, when the titrant A is added to B the following reaction takes place.

Equilibrium constant,

a = activity co-efficient.

Large values of the equilibrium constant K implies that the equilibrium concentration of A & B are very small at the equivalence point. It also indicates that the reverse reaction is negligible and the product C & D are very much more stable than the reactants A & B. Greater the value of K the larger the magnitude of the negative free energy change for the reaction between A & B. Since,

Where,

R = Universal gas Constant = 8.314 JK^-1mol^-1,
T = Absolute Temperature.

The reaction of the concentration of A & B leads to the reduction of the total free energy change. If the concentrations of A & B are too low the magnitude of the total free energy change becomes so small and the use of the reaction for titration will not be feasible.

Expressions of Concentration of Solutions:

The concentration or strength of solution means the amount of solute present in a given amount of the solution. The concentration may be expressed in physical or chemical units.

1.    Normality (N): It is defined as number of gram equivalents of the solute present in 1 litre (1000mL.) of the solution. If W g of solute of equivalent weight E is present in V mL of the solution, the normality of the solution is given by:

 

2.    Molarity (M): It is defined as the number of moles of the solute present in 1 litre (or 1000 mL) of the solution. A one molar solution contains 1 mole of the solute dissolved in 1 litre of the solution.

3.    Molality (m): It is defined as the number of moles of solute dissolved in 1000 g of the solvent. One molal solution contains one mole of the solute dissolved in 1000 g of the solvent.

           

Normal solution:

A solution containing one gram equivalent weight of the solute dissolved per litre is called a normal solution; e.g. when 40 g of NaOH are present in one litre of NaOH solution, the solution is known as normal (N) solution of NaOH. Similarly, a solution containing a fraction of gram equivalent weight of the solute dissolved per litre is known as subnormal solution. For example, a solution of NaOH containing 20 g (1/2 of g eq. wt.) of NaOH dissolved per litre is a sub-normal solution. It is written as N/2 or 0.5 N solution.

Formulae used in solving numerical problems on volumetric analysis;

1.    Strength of solution = Amount of substance in g litre^-1.

2.    Strength of solution = Amount of substance in g moles litre^-1.

3.    Strength of solution = Normality × Eq. wt. of the solute = molarity × Mol. wt. of solute.

4.    Molarity = Moles of solute/Volume in litre.

5.    Number of moles = Wt.in g/Mol. wt = M × V (initial) = Volume in litres/22.4 at NTP (only for gases).

6.    Number of milli moles = Wt. in g × 1000/mol. wt. = Molarity × Volume in mL.

7.    Number of equivalents= Wt. in g/Eq. wt = x × No. of moles × Normality × Volume in litre (Where x = Mol. wt/Eq. wt).

8.    Number of mill equivalents (meq.) = Wt. in g × 1000 / Eq. wt = normality × volume in mL.

9.    Normality = x × No. of mill moles (Where x = valency or change in oxi. number).

10.  Normality formula, N₁V₁ = N₂V₂, (Where N₁, N₂ → Normality of titrant and titrate respectively, V₁, V₂ → Volume of titrant and titrate respectively).

11.  % by weight = Wt. of solvent/Wt. of solution × 100 .

  

A solution is a homogeneous mixture of two or more components, the composition of which may be changed. The substan3333333ce which is present in smaller proportion is called the solute, while the substance present in large proportion is called the solvent.

Volumetric Analysis:

It involves the estimation of a substance in solution by neutralization, precipitation, oxidation or reduction by means of another solution of accurately known strength. This solution is known as standard solution.

Volumetric analysis depends on measurements of the volumes of solutions of the interacting substances. A measured volume of the solution of a substance A is allowed to react completely with the solution of definite strength of another substance B. The volume of B is noted. Thus we know the volume of the solutions A and B used in the reaction and the strength of solution B; so the strength of the other solution A is obtained. The amount (or concentration) of the dissolved substance in volumetric analysis is usually expressed in terms of normality. The weight in grams of the substance per litre of the solution is related to normality of the solution as,

Weight of the substance (g per litre) = Normality × gram equivalent weight of the substance.

Conditions of Volumetric Analysis:

i) The reaction between the titrant and titrate must be expressed.
ii) The reaction should be practically instantaneous.
iii) There must be a marked change in some physical or chemical property of the solution at the end point.
iv) An indicator should be available which should sharply define the end point.

Different methods to determine the endpoint include:

·         pH indicator:

A pH indicator is a substance that it changes its colour in response to a chemical change. An acid-base indicator changes its colour depending on the pH (e.g., phenolphthalein). Redox indicators are also frequently used. A drop of indicator solution is added to the titration at the start; at the endpoint has been reached the colour changes.

·         A potentiometer

It is an instrument that measures the electrode potential of the solution. These are used for titrations based on a redox reaction; the potential of the working electrode will suddenly change as the endpoint is reached.

·         pH meter:

It is a potentiometer that uses an electrode whose potential depends on the amount of H+ ion present in the solution. (It is an example of an ion-selective electrode.) This allows the pH of the solution to be measured throughout the titration. At the endpoint, there will be a sudden change in the measured pH. This method is more accurate than the indicator method and is very easily automated.

·         Conductance:

The conductivity of a solution depends on the ions present in it. During many titrations, the conductivity changes significantly. (i.e., during an acid-base titration, the H⁺ and OH^- ions react to form neutral H₂O, this changes the conductivity of the solution.) The total conductance of the solution also depends on the other ions present in the solution, such as counter ions. This also depends on the mobility of each ion and on the total concentration of ions that is the ionic strength.

·         Colour change:

In some reactions, the solution changes colour without any added indicator. This is often seen in redox titrations, for instance, when the different oxidation states of the product and reactant produce different colours.

·         Precipitation:

In this type of titration the strength of a solution is determined by its complete precipitation with a standard solution of another substance.

eg: 

Acid base titration:

The chemical reaction involved in acid-base titration is known as neutralisation reaction. It involves the combination of H₃O⁺ ions with OH^- ions to form water. In acid-base titrations, solutions of alkali are titrated against standard acid solutions. The estimation of an alkali solution using a standard acid solution is called acidimetry. Similarly, the estimation of an acid solution using a standard alkali solution is called alkalimetry.

 The Theory of Acid–Base Indicators: 

Ostwald, developed a theory of acid base indicators which gives an explanation for the colour change with change in pH. According to this theory, a hydrogen ion indicator is a weak organic acid or base. The undissociated molecule will have one colour and the ion formed by its dissociation will have a different colour.

Let the indicator be a weak organic acid of formulae HIn. It has dissociated into H⁺ and In^- . The unionized molecule has one colour, say colour (1), while the ion, In^- has a different colour, say colour (2). Since HIn and In^- have different colours, the actual colour of the indicator will dependent upon the hydrogen ion concentration [H⁺]. When the solution is acidic, that is the H⁺ ions present in excess, the indicator will show predominantly colour (1). On other hand, when the solution is alkaline, that is, when OH^- ions present in excess, the H⁺ ions furnished by the indicator will be taken out to form undissociated water. Therefore there will be larger concentration of the ions, In^-. thus the indicator will show predominantly colour (2). 

Some indicators can be used to determine pH because of their colour changes somewhere along the change in pH range. Some common indicators and their respective colour changes are given below. 

Indicator	Colour on Acidic Side	Range of Colour Change	Colour on Basic Side
Methyl Violet	Yellow	0.0 - 1.6	Violet
Bromophenol Blue	Yellow	3.0 - 4.6	Blue
Methyl Orange	Red	3.1 - 4.4	Yellow
Methyl Red	Red	4.4 - 6.2	Yellow
Litmus	Red	5.0 - 8.0	Blue
Bromothymol Blue	Yellow	6.0 - 7.6	Blue
Phenolphthalein	Colourless	8.3 - 10.0	Pink
Alizarin Yellow	Yellow	10.1 - 12.0	Red

i.e., at pH value below 5, litmus is red; above 8 it is blue. Between these values, it is a mixture of two colours.

Indicators Used for Various Titrations:

1. Strong Acid against a Strong Base:

Let us consider the titration of HCl and NaOH. The pH values of different stages of titration shows that, at first the pH changes very slowly and rise to only about 4. Further addition of such a small amount as 0.01 mL of the alkali raises the pH value by about 3 units to pH 7. Now the acid is completely neutralized. Further of about 0.01 mL of 0.1 M NaOH will amount to adding hydrogen ions and the pH value will jump to about 9. Thus, near the end point, there is a rapid increase of pH from about 4 to 9.

An indicator is suitable only if it undergoes a change of colour at the pH near the end point. Thus the indicators like methyl orange, methyl red and phenolphthalein can show the colour change in the ph range of 4t0 10. Thus, in strong acid- strong base titrations, any one of the above indicators can be used. 

2. Weak Acid against Strong Base:

Let us consider the titration of acetic acid against NaOH. The titration shows the end point lies between pH 8 and 10. This is due to the hydrolysis of sodium acetate formed. Hence phenolphthalein is a suitable indicator as its pH range is 8-9.8. However, methyl orange is not suitable as its pH range is 3.1 to 4.5. 

3. Strong Acid against Weak Base:

Let us consider the titration ammonium hydroxide against HCl. Due to the hydrolysis of the salt, NH₄Cl, formed during the reaction, the pH lies in the acid range. Thus, the pH at end point lies in the range of 6 to 4. Thus methyl orange is a suitable indicator while phenolphthalein is not suitable.

StrongAcids	StrongBases	WeakAcids	WeakBases
HCl	NaOH	Acetic acid	Ammonia
HNO3	KOH	Hydrocyanic  acid	Magnesium  hydroxide
HBr	etc	HF	Pyridine
H2SO4		Oxalic acid	Sodium carbonate
HI		Ethanoic acid	Potassium carbonate
HClO4		etc	etc

 

Precipitation Titration:

A titrimetric method based on the formation of a slightly soluble precipitate is called a precipitation titration. The most important precipitation process in titrimetric analysis utilizes silver nitrate as the reagent (Argentimetric process).

Many methods are utilized in determining end points of these reactions, but the most important method, the formation of a coloured precipitate will be considered here.

1.    In the titration of a neutral solution of chloride ions with silver nitrate, a small quantity of potassium chromate solution is added to serve as the indicator. At the end point the chromate ions combine with silver ions to form the sparingly soluble brick-red silver chromate. This is a case of fractional precipitation, the two sparingly soluble salts being AgCl (Ksp = 1.2 x 10^-10) and Ag₂CrO₄ (Ksp = 1.7x10^-12).

AgCl is the less soluble salt and initially chloride concentration is high, hence AgCl will be precipitated. Once the chloride ions are over and with the addition of small excess of silver nitrate solution brick red colour silver chromate becomes visible. The titration should be carried out in neutral solution or in very faintly alkaline solution. i.e. within the pH range 6.5-9.

In acid solutions following reaction occurs.

Consequently the chromate ions concentration is reduced and the solubility product of silver chromate may not be exceeded. In markedly alkaline solution, silver hydroxide (Ksp = 2.3 x 10⁸⁾ might be precipitated.

2.    The titration can be carried out with dichlorofluorescein as the indicator. Dichlorofluorescein is an e6 lo8xample of an adsorption indicator. Adsorption indicators have the interesting property of changing colour when they stick (adsorb) to the surface of a precipitate. During the titration the dichlorofluorescein molecules exist as negatively charged ions (anions) in solution. As the AgCl precipitate is forming, the excess Cl^- ions in the solution form a layer of negative charge on the precipitate surface. As the equivalence point is reached and passed, the excess Cl^- ions on the precipitate surface are replaced by excess Ag+ ions, giving the surface a positive charge. The negatively charged indicator will be attracted to the positively charged precipitate surface where it absorbs and changes colour. The suspended precipitate will have a pink tinge because of some premature displacement of chloride ion by the dichlorofluorescein ion. When the pink colour starts to persist for slightly longer periods of time, the drip rate is lowered. The end point is reached when the entire solution turns pink. It is important that the AgCl precipitate be prevented from coagulation during the titration. For this reason a small amount of dextrin is added to the solution.

Complexometric Titration:

This type of titration depends upon the combination of ions (other than H⁺ and OH^-) to form a soluble ion or compound as in the titration of a solution of a cyanide with AgNO₃.

Principle of Complexometric Titration:

Complexometric titrations are particularly useful for determination of a mixture of different metal ions in solution. Ethylene diamine tetra acetic acid (EDTA), is a very important reagent for complex formation titrations. EDTA has been assigned the formula II in preference to I since it has been obtained from measurements of the dissociation constants that two hydrogen atoms are probably held in the form of zwitter ions.

 

 

EDTA behaves as a dicarboxylic acid with two strongly acidic groups. For simplicity EDTA may be given the formula H₄Y, the disodium salt is therefore Na₂H₂Y and it has the complex forming ion H₂Y^2- in aqueous solution. The reactions with cationsmay be represented as;

M²⁺ + H₂Y ₂^-→ MY^2- + 2H⁺
M³⁺ + H₂Y ₂^-→ MY^- + 2H⁺
M⁴⁺+ H₂Y ₂^-→ MY + 2H⁺

One gram ion of the complex-forming ion H₂Y^2- reacts in all cases with one gram ion of the metal. EDTA forms complexes with metal ions in basic solutions. In acid-base titrations the end point is detected by a pH sensitive indicator. In the EDTA titration metal ion indicator is used to detect changes of pM. It is the negative logarithm of the free metal ion concentration, i.e., pM = - log [M²⁺]. Metal ion complexes form complexes with specific metal ions. These differ in colour from the free indicator and a sudden colour change occurs at the end point. End point can be detected usually with an indicator or instrumentally by potentiometric or conductometric (electrometric) method.