Unit 9: Acid-Base Chemistry — Long Questions

Ans. General properties of acids, bases and indicators

• Scientists have recognized that acids and bases have distinct characteristic properties, since the earliest days of experimental chemistry.
• Acids have a sour taste while bases are bitter.
• Acids and bases change the colour of certain dyes called indicators, such as litmus and phenolphthalein.
• Acids change litmus from blue to red and phenolphthalein from pink to colourless.
• Bases change litmus from red to blue and phenolphthalein from colourless to pink. From these colour changes, acids and bases neutralize each other.
• During neutralization, acids and bases react with each other to produce ionic substances called salts.

Ans. Bronsted-Lowry Concept

Introduction In 1923, the Danish chemist J.N. Bronsted and the English chemist T.M. Lowry independently expanded the Arrhenius theory of acid-base.

Need of Bronsted-Lowry concept
The limitations of Arrhenius theory are that it describes the reaction of an acid and base only in the aqueous medium. There are many reactions that occur in solvents other than water or in the absence of any solvent.

Bronsted-Lowry pointed out that acid-base reactions can be seen as proton-transfer reactions and that acids and bases can be defined in terms of this proton (H+) transfer.

Definition of Acid An acid is the specie donating a proton in a proton-transfer reaction.

Definition of Base A base is a specie that accept proton from the other species in a proton transfer reaction.

When HCl dissolves in water, an H+ ion (a proton) is transferred from HCl to water, where it becomes attached to a lone pair of electrons on the O atom and forms H3O+. In effect, HCl (the acid) has donated the H+ and H2O (the base) has accepted it.

HCl(g) + H2O(l) ⇌ H3O+(aq) + Cl−(aq)
Acid Base Conjugate acid Conjugate base

Conjugate Acid A specie formed after a Bronsted base accepts a proton from the acid is called the conjugate acid.

Example The hydronium ion (H3O+) is the conjugate acid of water.

Conjugate Base A specie formed when an acid donates a proton to a base is called the conjugate base. Cl− is conjugate base of HCl. HCl-Cl− and H2O-H3O+ are conjugate acid-base pairs.

Example When ammonia dissolves in water, proton transfer also occurs. An H+ from H2O attaches to the N atom's lone pair and NH4+ and OH− are formed. With fewer H+, the water molecule becomes OH− ion:

NH3(g) + H2O(l) ⇌ OH−(aq) + NH4+(aq)
Base Acid Conjugate base Conjugate acid

H2O (the acid) has donated the proton and NH3 (the base) has accepted it. NH3-NH4+ and H2O-OH− are conjugate acid-base pair.

Definition of Amphoteric An amphoteric substance is a species that can act as either an acid and a base (it can lose or gain a proton), depending on the other reactant.

H2O is amphoteric in nature because it acts as a base in one case and as an acid in the other.

Application of Bronsted-Lowry theory
It can be applied to acids in solvents other than water or even solventless reactions. The reaction between gaseous ammonia and HCl gives solid NH4Cl.

HCl(g) + NH3(g) ⟶ NH4Cl(s)

HCl is an acid because in the reaction it donates a proton to the NH3 molecule. The NH3 acts as a base, even though hydroxide ion OH− is not present, and accepts a proton from the HCl molecule.

Ans. Lewis concept of acids and bases

Introduction In 1923, G.N. Lewis proposed a generalized definition of acid-base behavior in which acids and bases are identified by their ability to accept or to donate a pair of electrons and form a coordinate covalent bond.

Lewis Acid A Lewis acid is any species (molecule or ion) that can accept a pair of electrons.

Lewis Base A Lewis base is any species (molecule or ion) that can donate a pair of electrons.

Reaction of NH3 and Ag+
A Lewis acid-base reaction occurs when a base donates a pair of electrons to an acid. The two ammonia molecules act as Lewis bases, donates a pair of electrons to a positively charged silver ion (the Lewis acid).

Adduct The sum of charges on the left side is +1, so the acid-base adduct on the right hand side must carry a charge of +1.

Reaction of BF3 and F−
The boron atom in boron trifluoride, BF3, has only six electrons in its valence shell. The boron atom has an incomplete octet, it can behave as an electron pair acceptor. As a result, BF3 is a very good Lewis acid and reacts with many Lewis base; a fluoride ion in the Lewis base in this reaction, donating one of its lone pairs:

Adduct The negative charge on the adduct is the sum of charges on the left hand side of the equation is -1, the sum of charges on the right hand side must also be -1.

Ans. Definition The product of concentration of H+ and OH− ions in water is called ionic product of water or dissociation constant of water (Kw).

Explanation

• Pure water is a poor conductor of electricity but its conductance is measurable.
• Water undergoes self-ionization reversibly as follows,

H2O(c) + H2O(l) ⇌H3O+(aq) + OH−(aq)
Net reaction H2O(l) ⇌ H+(aq) + OH−(aq)

The equilibrium constant for this reaction can be writeth as follows:

Kc = [H+ ][OH−] = 1.8×10−16mol dm−3
[H2O]

• The concentration of H2O, [H2O] in pure water may be calculated to be 1000g dm−3 divided by 18g mol−1 giving 55.5 mol dm−3
• Water is present in very large excess and very few of its molecules undergo ionization, so its concentration remain effectively constant.
• Constant concentration of water is taken on left hand side and multiplied with Kc to get another constant called Kw.

Kc [H2O] = [H+ ][OH−]
1.8 × 10−16 × 55.5 = 1.01 × 10−14 = [H+][OH−]
This value, 1.01 × 10−14 is called Kw of water at 25°C.
Kc [H2O] = [H+][OH−]
Kw = [H+][OH−] = 10−14 at 25°C

Effect of temperature on Kw: The effect of temperature on Kw is shown in table. The value of Kw increases almost 75 times when temperature is increased from 0°C to 100°C.

For neutral water
[H+] = [OH−]
[H+][H+] = 10−14
[H+]2 = 10−14 (at 25°C)
[H+] = 10−7 moldm−3
The concentration of OH− is also the same in neutral water.
[OH−] = 10−7 moldm−3

Effect of addition acid and base on Kw
Whenever some quantity of acid or base is added to water, then Kw remains the same, but [H+] and [OH−] are no more equal.

Addition of acid to water: In case of addition of small amount of acid.
[H+] > [OH−]

Addition of base to water: In the case of addition of few drops of a base.
[OH−] > [H+]

During both of these additions, the values of Kw will remain the same, i.e. 10−14 at 25°C.

Ans. pH

Definition The negative logarithm of molar concentration of H+ ions in a mixture is called pH.

pOH

Definition The negative logarithm of molar concentration of OH− ions in a mixture is called pOH.

Explanation

• In all the aqueous solutions, the concentration of H+ and OH− are too low to be conveniently expressed and used in calculations.
• In 1909, Sorenson, a Danish biochemist, introduced the term pH and pOH. So, the scales of pH and pOH were developed. These two quantities can be calculated as:

pH = -log[H+]
pOH = -log[OH−]

For neutral water,
pH = -log[10−7] = 7
pOH = -log[10−7] = 7
pH + pOH = 14

• The value of pH normally varies between 0 → 14 at 25°C. The pH can have a negative value or greater than 14 pH can be possible.
• The pH values of some familiar aqueous solutions are shown in figure. This table helps to understand the acidic or basic nature of commonly used solutions.

Ans. Ionization constant of acids

Definition The value of Ka called the dissociation constant of an acid, is the quantitative measure of the strength of the acid.

Explanation
Acids and bases, when dissolved in water, may or may not be completely dissociated. Many acids are weak electrolytes and ionize to an extent which is much less than 100%.

Derivation Suppose we have an acid HA dissolved in water,

HA + H2O ⇌ H3O+ + A−

Kc for the reversible reaction will be written as follows:

Kc = [H3O+][A−]/([HA][[H2O])

Kc[H2O]= [H3O+][A−]/[HA]

Let Kc[H2O]= Ka

Hence Ka = [H3O+][A−]/[HA]

• This equation can be used to calculate Ka for any acidic solution if we know the pH or [H+] of that solution and the initial concentration of the dissolved acid [HA].
• This can also be used to calculate the equilibrium concentration of H3O+ and A− produced if we know the initial concentration of acid HA and its Ka value.

Effect of Ka on strength of acid

(i) Ka < 10−3 acid is weak.
(ii) Ka = 1 to 10−3 acids are moderately strong.
(iii) Ka> 1 acid is strong.

Ans. Calculating H3O+ ions from Ka

• The weak acids do not completely dissociate in water.
• The concentration of hydronium ions is not equal to the initial concentration of the acid.
• The equilibrium expression is used to calculate the H3O+ concentration.

For a weak acid HA, the dissociation in water can be represented as:

HA(aq) ⇌ H+(aq) + A−(aq)

In water, H+ combines with water to form H3O+:

H+(aq) + H2O ⇌ H3O+(aq)

The equilibrium equation in terms of H3O+(aq) is:

HA(aq) ⇌ H3O+(aq) + A−(aq)

Let the initial concentration of the weak acid HA be C mol dm−3.

At equilibrium [H3O+] = x; [A−] = x and [HA] = C-x

The expression for the acid dissociation constant (Ka) is:

Ka = [H3O+][A−]/[HA]

Substitute the equilibrium concentrations into this expression:

Ka = (x × x)/(C − x) = x²/(C − x)

For weak acids, x is usually very small compared to C, so C-x ≈ C. Therefore, the expression simplifies to:

Ka ≈ x²/C

x² ≈ Ka × C

Taking the square root of both sides:

x ≈ √(Ka × C)

Thus, the concentration of H3O+ is given by:

x ≈ √(Ka × C)

Ans. Common Ion Effect

Definition The suppression of ionization of a weak electrolyte by adding a common ion to it is called common ion effect.

Explanation

Purification of sodium chloride NaCl is purified by passing hydrogen chloride gas through brine (saturated solution of NaCl). Sodium chloride is fully ionized in the solution. Equilibrium constant expression for this process can be written as follows:

NaCl(s) ⇌Na+(aq) + Cl−(aq)

Kc = [Na+][Cl−(aq)]/[NaCl]

HCl also ionizes in solution:

HCl(g) ⇌H+(aq) + Cl−(aq)

• On passing HCl gas, concentration of Cl− ion is increased.
• NaCl crystallizes out of the solution to maintain the constant value of the equilibrium constant. This type of effect is called the common ion effect.
• The addition of a common ion to the solution of a less soluble electrolyte suppresses its ionization and the concentration of unionized species increases, which may come out as a precipitate.

Na+(aq) + Cl−(aq) ⇌NaCl(s)

More examples of common ion effect

(i) Solubility of KCℓO3

The solubility of less soluble salts KClO3 in water is suppressed by the addition of a more soluble salt KCl by common ion effect. K+ is a common ion. The ionization of KClO3 is suppressed and it settles down as precipitate.

KClO3(s) ⇌ K+(aq) + ClO3−(aq)

KCl(s) ⇌ K+(aq) + Cl−(aq)

(ii) Solubility of H2S

The dissociation of a weak acid H2S in water can be suppressed by the addition of stronger acid HCl. H+ is a common ion. H2S becomes less dissociated in acidic solution. In this way, low concentration of S2− ion is produced.

H2S(g) ⇌ 2H+(aq) + S2−(aq)

This low concentration of S2− ions helps to do the precipitation of radicals of second group basic radicals during salt analysis.

HCl(aq) ⇌ H+(aq) + Cl−(aq)

(iii) Solubility of NH4Cℓ

An addition of NH4Cl in NH3 solution suppresses the concentration of OH− due to the presence of a large excess of NH4+ from NH4Cl. Actually, NH4Cl is a strong electrolyte. The combination of these two substances is used as a group reagent in third group basic radicals for salt analysis.

NH3Cℓ(aq) ⇌NH3(aq) + Cℓ−(3)

NH3 + H2O(aq) ⇌ NH4(aq) + OH−(aq)

Common ion effect finds extensive applications in the qualitative analysis and the preparation of buffers.

Ans. Applications of common ion effect

(i) Purification of NaCℓ

NaCl is purified by passing hydrogen chloride gas through the saturated brine. Sodium chloride is fully ionized in the solution.

NaCl(s) ⇌Na+(aq) + Cl−(aq)

HCl(g) ⇌H+(aq) + Cl−(aq)

On passing HCℓ gas, concentration of Cℓ− ions is increased. Therefore, NaCℓ crystallizes out of the solution to maintain the constant value of equilibrium constant.

(ii) Precipitation of KCℓO3

The solubility of a less soluble salt KCℓO3 in water is suppressed by the addition of more soluble salt KCℓ by common ion effect. K+ is a common ion. The ionization of KCℓO3 is suppressed and it settles down as precipitate.

KClO3(s) ⇌ K+(aq) + ClO3−(aq)

KCl(s) ⇌ K+(aq) + Cl−(aq)

(iii) Precipitation of 2nd group basic radicals

The dissociation of a weak acid H2S in water can be suppressed by the addition of stronger acid HCℓ. H+ is a common ion. H2S becomes less dissociated in acidic solution. In this way, low concentration of S2− is developed.

This low concentration of S2− ions helps to do the precipitation of radicals of second group basic radicals during salt analysis.

(iv) Preparation of buffers

Common ion effect has its extensive applications in the qualitative analysis and preparation of buffer solutions.

Ans. Buffer Solutions
Definition: The solutions which resist the change in their pH when a small amount of an acid or a base is added to them, are called buffer solutions.
They have a specific constant value of pH and their pH values do not change on dilution and on keeping for a long time.

Types of Buffer
Buffer solutions are mostly prepared by mixing two substances.
(i) Acidic Buffers
• By mixing a weak acid and its salt with a strong base. Such solutions give acidic buffers with pH less than 7.
• Mixture of acetic acid and sodium acetate is one of the best examples of acidic buffers.
(ii) Basic Buffers
• By mixing a weak base and its salt with a strong acid. Such solutions will give basic buffers with pH more than 7.
• Mixture of NH4OH and NH4Cl is one of the best examples of basic buffers.

Mechanism of action of Buffers
• The example of an acidic buffer consisting of CH3COOH and CH3COONa. CH3COOH being a weak electrolyte undergoes very little dissociation.
• When CH3COONa, which is a strong electrolyte, is added to CH3COOH solution, then the dissociation of CH3COOH is suppressed, due to common ion effect of CH3COO-.
CH3COOH(aq) + H2O(l) ⇌ CH3COO-(aq) + H3O+(aq)
CH3COONa(aq) ⇌ CH3COO-(aq) + Na+(aq)
• If one goes on adding CH3COONa in CH3COOH solution, then the added concentrations of CH3COO- decrease the dissociation of CH3COOH and the pH of solution increases.
The table 9.4 tells us how the pH value of a mixture of two compounds is maintained. Greater the concentration of acetic acid as compared to CH3COONa, lesser is the pH of solution.

Table 9.4 Effect of addition of acetate ions on the pH of acetic acid solution
[CH3COOH] (mol dm-3) | [CH3COO-] (mol dm-3) | % Dissociation | pH
0.10 | 0.00 | 1.3 | 2.89
0.10 | 0.05 | 0.036 | 4.44
0.10 | 0.10 | 0.018 | 4.74
0.10 | 0.15 | 0.012 | 4.92

• A buffer mentioned above is a large reservoir of CH3COOH and CH3COO- components.
• When an acid or H3O+ ions are added to this buffer, they will react with CH3COO- to give back acetic acid and hence the pH of the solution will almost remain unchanged. The reason is that CH3COOH being a weak acid will prefer to remain un-dissociated.
• The buffer solution consisting of NH4Cl and NH4OH, can resist the change of pH and pOH, when acid or base is added from outside.
• When a base or OH- ions are added in it, they will react with H3O+ to give back H2O and the pH of the solution again will remain almost unchanged.

Ans. Calculating the pH of a Buffer

Consider a weak acid HA and its salt NaA with a strong base say NaOH. The reversible reactions for dissociation of HA and NaA are as follows:

HA ⇌ H+ + A−
NaA ⇌ Na+ + A−

pH of a buffer solution can be calculated by using Henderson equation as given below:

pH = pKa − log [Acid]/[Salt]

Interchanging the numerator and denominator, the sign of log changes.

pH = pKa + log [Salt]/[Acid]

The relationship is called Henderson's equation.

• The equation shows that two factors evidently govern the pH of a buffer solution.
• First the pKa of the acid used and the second is the ratio of the concentrations of the salt and the acid.
• The best buffer is prepared by taking equal concentration of salt and acid.
• pH is controlled by pKa of the acid.

Example For the buffer solution of acetic acid and sodium acetate, the pH can be calculated as follows.

[CH3COOH] = [CH3COONa]

Then pH = pKa + log [CH3COONa]/[CH3COOH]

pH = pKa + log(1)

So, pH = pKa + 0 = pKa

pH = 4.74

It means that the pH of this buffer is just equal to the pKa of the acid.

Ans. Solubility Product
Definition: The solubility product is the product of the concentrations of ions raised to an exponent equal to the co-efficient of the balanced equation.
• The value of Ksp is a measure of the dissociation of sparingly soluble salt.
• The dissociation is completed at equilibrium stage for NaCl when it is mixed in water. Because it is a completely soluble salt.
• The dissociation is not complete at equilibrium stage in slightly soluble salts like PbCl2, PbSO4 etc.
• Ksp is only for slightly or sparingly soluble salts.

Slightly soluble ionic compounds
(i) Lead sulfate (PbSO4): When it is shaken with water the solution contains Pb2+, SO42- and undissociated PbSO4.it means that equilibrium exists between solid solute, PbSO4 and the dissolved ions, SO42- and Pb2+.

PbSO4(s) ⇌[H2O] Pb2+(aq) + SO42-(aq)

Kc = [Pb2+ ][SO42- ]/[PbSO4]

Being a sparingly soluble salt, the concentration of lead sulphate (PbSO4) almost remains constant. Bring [PbSO4] on left hand side (L.H.S) with Kc.

Kc[PbSO4] = [Pb2+ ][SO42- ]

If Kc[PbSO4] = Ksp

Ksp = [Pb2+ ][SO42- ] = 1.6×10-8 at 25°C

Ksp is called the solubility product of PbSO4. It is the product of molar solubilities of two ions at equilibrium stage.

(ii) Lead sulphate (PbCl2): It is a well-known sparingly soluble compound and it dissociates to a very small extent like PbCl2.
Ksp = [Pb2+][Cl-]2

Effect of temperature
Ksp is usually a very small quantity at room temperature. The value of Ksp is temperature dependent. It increases by increase in temperature.
The following table 9.5 shows the Ksp values of some slightly soluble ionic compounds.

Table 9.5 Ksp values of some ionic compounds
Salt | Ion Product | Ksp | Salt | Ion Product | Ksp
AgBr | [Ag+][Br-] | 5.0×10-13 | CuS | [Cu2+][S2-] | 8×10-34
AgCl | [Ag+][Cl-] | 1.8×10-10 | Fe2S3 | [Fe3+]2[S2-]3 | 1.4×10-85
Al(OH)3 | [Al3+][OH-]3 | 3×10-34 | MgCO3 | [Mg2+][CO32-] | 3.5×10-8
BaCO3 | [Ba2+][ CO32- ] | 1.1×10-10 | MnS | [Mn2+][S2-] | 3×10-11
CaCO3 | [Ca2+][CO32-] | 3.3×10-9 | PbCrO4 | [Pb2+][CrO42-] | 2.3×10-13
CaF2 | [Ca2+][F-]2 | 3.2×10-11 | PbSO4 | [Pb2+][SO42-] | 1.6×10-8

Applications of solubility product
(a) Determination of Solubility from Ksp
• To calculate solubility from Ksp, the formula of the compound and Ksp value is required.
• The unknown molar solubility S is calculated and the concentration of the ions is determined.

Table 9.6 shows the relationship between the Ksp values and the solubility of some sparingly soluble compounds.

Table 9.6 The relation between solubility and solubility products of some salts
Formula | No. of ions | Ksp | Solubility gdm-3
PbSO4 | 2 | 1.69 × 10-4 | 1.3 × 10-4
Ca(OH)2 | 3 | 6.5 × 10-6 | 1.175 × 10-2
CaF2 | 3 | 3.2 × 10-11 | 2.0 × 10-4
Ag2CrO4 | 3 | 2.6 × 10-12 | 8.7 × 10-5

(b) Common Ion Effect
• The presence of a common ion decreases the solubility of a slightly soluble ionic compound.
• In order to explain it, consider a saturated solution of PbCrO4, which is a sparingly soluble ionic salt.

PbCrO4(aq) ⇌ Pb2+(aq) + CrO42-(aq)

Now add Na2CrO4 which is a soluble salt. CrO42- is the common ion. It combines with Pb2+ to form more insoluble PbCrO4. So equilibrium is shifted to the left to keep Ksp constant.

(c) Predicting Precipitation
The solubility product can also help in predicting whether the precipitation of a salt will occur or not.

Example The solubility product of CaSO4 is 2 x 10-5. If we add 10-2 mol dm-3 solution Ca2+ to 10-2 mol dm-3 solution of SO42- ions at 25°C.
The concentrations of each ionic species can be calculated as

[Ca2+ ] = [SO42- ] = 10-2 / 2 = 5.0×10-3 mol dm -3

[Ca2+ ][SO42- ] = 5.0×10-3 mol dm -3 × 5.0×10-3 mol dm -3
= 2.5× 10-5 mol2 dm -6
= 2.5× 10-5 mol2 dm -6 > Ksp of CaSO4

As the ionic product of concentrations is greater than Ksp, therefore CaSO4 will precipitate out.

Ans. Salt hydrolysis
Definition: The breakdown of a salt (made up of weak acid or base) by reacting with water, resulting in the bond breaking in that salt which changes the pH of water is called hydrolysis.

Explanation
When a salt dissolves in water, it dissociates into its constituent ions. These ions can interact with water, affecting the solution's pH depending on the nature of the acid and base from which the salt is derived. It is called salt hydrolysis.

(i) Salts of strong acids and strong bases
In salts of strong acids and strong bases like sodium chloride (NaCl).
• The conjugate base of a strong acid (Cl- from HCl) is very weak and does not significantly react with water.
• The conjugate acid of a strong base (Na+ from NaOH) is also very weak and does not significantly react with water.
NaCl(s) ⇌[H2O] Na+(aq) + Cl-(aq)
• Na+ is the conjugate acid of NaOH (a strong base) and does not affect the pH.
• Cl- is the conjugate base of HCl (a strong acid) and does not affect the pH.
The solution remains neutral.

(ii) Salts of strong acids and weak bases
In salts of strong acids and weak bases like ammonium chloride (NH4Cl).
• The conjugate base of the strong acid (Cl- from HCl) does not react with water.
• The conjugate acid of the weak base (NH4+ from NH3) reacts with water to produce H3O+ ions, making the solution acidic.
NH4Cl(s) ⇌ NH4+(aq) + Cl-(aq)
• NH4+ is the conjugate acid of NH3 (a weak base) and reacts with water:
NH4+(aq) + H2O(l) ⇌ NH3(g) + H3O+(aq)
• Cl- is the conjugate base of HCl (a strong acid) and does not affect the pH.
The solution is acidic due to the (NH4+ ) ions.

(iii) Salts of weak acids and strong bases
In salts of strong of weak acid and strong base like sodium acetate (CH3COOOH).
CH3COONa(s) + H2O(l) ⇌ CH3COO-(aq) + Na+(aq)
• CH3COO- is the conjugate base of CH3COOH (a weak acid) and reacts with water:
CH3COO-(aq) + H2O(l) ⇌ CH3COOH(aq) + OH-(aq)
• Na+ is the conjugate acid of NaOH (a strong base) and does not affect the pH.
The solution is basic due to the acetate (CH3COO- ) ions.

(iv) Salts of weak acids and weak bases:
In salts of weak acids and weak bases like ammonium acetate ( CH3COO- NH3+).
• The conjugate acid (NH4+) and the conjugate base (CH3COO-) both affect the pH.
CH3COO- NH3+(aq) ⇌ NH3+(aq) + CH3COO-(aq)
• NH4+ hydrolyses to produce H3O+:
NH4+ + H2O ⇌ NH3 + H3O+
• CH3COO- hydrolysis to produce OH-:
CH3COO- + H2O ⇌ CH3COOH(aq) + OH-
The resultant pH of the solution depends on the relative strengths of the conjugate acid and conjugate base. The solution may be acidic, basic, or nearly neutral, depending on which reaction is more dominant.

Ans. Acid base indicators
Definition: An indicator is a substance that changes colour to mark a titration's endpoint.
• Acid-base indicators exhibit one colour in acid and another in base.
• Most indicators used in acid-base titration are weak organic acids or weak organic bases.

Acid solution In solution, a weak-acid indicator (HIn) can be represented by the equation below.
HIn ⇌ H+ + In-

In- is the symbol of the anion part of the indicator. Because the reaction is reversible, both HIn and In- are present. The colours displayed result from the fact that HIn and In- have different colours.

Acidic solutions In acidic solutions, any In- ions that are present act as Bronsted bases and accept protons from the acid. The indicator is then present in largely unionized form, HIn.

Basic Solutions In basic solutions, the OH- ions form the base combine with the H+ ions produced by the indicator.
The indicator molecules further ionize to equalize the loss of H+ ions.
• The indicator is thus present largely in the form of its anion, In-. The solution now show the base-indicating colour, which for litmus is blue.

Colour of indicators
Different indicator change colour at different pH values. The colour depends on the relative amount of HIn and In- at a given pH.

Example Methyl red changes from red to yellow between pH 4.4 and 6.2. At pH 4.4, the indicator exists mostly as HIn molecules, which appear red in the solution.

The indicator ranges are given below for some of the indicators commonly used.

Methyl orange | Orange | Yellow | (color range visualization)
Phenolphthalein | Colourless | (range 0-14 scale) | (color range visualization)

Fig Range and colour changes of some common Acid-Base indicators

Selecting a suitable indicator
The two general criteria for an indicator to be used in a titration are:
• The pH at the end of the titration should be close to the indicator's neutral point.
• The indicator should indicate a sharp colour change near the equivalence point of the titration.

Each pH indicator changes colour over a defined range of pH, known as the indicator range. An indicator changes colour over a range of about 2 pH units.

Ans. Titration curve and Equivalence point
Definition of titration curve: A pH curve is a graph of the pH of the solution verses the volume of titrant added.

Definition of equivalence point The equivalence point is the point at which the amount of titrant added is just enough to neutralize the analyte solution completely.

Explanation of titration curve
• Titration curves show how the pH of an acidic or basic solution changes as a basic or acidic solution is added to it.
• The titration curve to choose an indicator that will show when the titration is complete and we reach the equivalence point.
• The end point of the titration occurs when the indicator changes colour.
• We choose and indicator with an end point close to the equivalence point.

(i) Strong Acid-strong base titration curve
The titration of HCl with NaOH is an example of strong acid-strong base titration.
• The solution contains only the strong acid at the start.
• As the acid is strong, it completely dissociates leading to a high concentration dissociates, leading to a high concentration of H+ ions and the pH is very low.
• By the addition of NaOH, OH- ions from NaOH begin to neutralize the H+ ions from HCl and pH of the solution rises.
• The equivalence point is reached when the amount of OH- added is stoichiometrically equal to the amount of H+ originally present in the acid.
• The pH at the equivalence point in a strong acid-strong base titration is 7.
• The titration curve at this stage shows a steep rise in pH, changing quickly form acidic to neutral.
• The pH value at the endpoint changes about from 4.0 to 10.0.

Conclusion
A pH of 7.0 is at equivalence, so, such phenolphthalein as indicator can be used as they show different colours in this range.

(ii) Strong acid-weak base titration curve
The titration of NH3 with HCl is an example of strong acid-weak base titration.

(a) Initial stage
• The solution contains only the weak base at start, dissociates partially, with a lower concentration of OH- ions compared to a strong base.
• The initial pH will be greater than 7 but lower than the pH of a strong base.
• By the addition of HCl, the concentration of OH- decreases, results in further decrease in pH.

(b) Change in pH
• Before reaching the equivalence point the solution is in a buffer-like region where the pH changes more gradually.
• The presence of the weak base and its conjugate acid (from the salt formed) creates a buffering effect, which helps to moderate the pH change as the base is added.

(c) Equivalence point
The pH at the equivalence point will be less than 7 (pH = 5.27) because the conjugate acid slightly dissociates, releasing H+ ions making the solution acidic (NH4+ → NH3 + H+).

(d) Titration curve
• After the equivalence point, the pH of the solution decreases rapidly. Because the strong acid dissociates completely in water, providing a high concentration of hydrogen ions (H+), which significantly lowers the pH.
• Methyl orange has its colour change in this range, therefore, it can be used as an indicator for strong acid-weak base titrations.

Table 9.7 pH ranges of common indicators
Indicators | Acid color | Base colour | pH Range | Type of Titration
Methyl orange | Orange | Yellow | 3.2-4.5 | Strong acid-strong base / Strong acid-weak base
Phenolphthalein | colorless | red | 8.2-10.0 | Weak acid-strong base