BSc Semester: I

GENERALCHEMISTRY

Paper : 1

Course code : US01 CCHE21

PHYSICAL CHEMISTRY

Chemical science is chemistry, likes

In-Organic chemistry

Organic Chemistry

Physical chemistry

and Analytical Chemistry are principal subject of chemistry.

How we are able to understand and gain the knowledge of chemical science ?

Simple, you should read the basic of chemistry and practice to remember as general fact of subject.

Basic chemistry

Periodic table

Electronic configuration

General principle

Characteristic of chemical substance

Properties of substance

Hybridization of orbital

Bond formation.

UNIT: III IONIC EQUILIBRIUM

• Property of substance: Acids and Bases
• Theory of acid and base: Arrhenius theory
• The Lowry-Bronsted theory
• Lewis acid base theory
• Acidity and basicity, PH scale, Hydrolysis, Buffer Solution, Indicator, Sparingly soluble salt, common ions effect, selective precipitation, numerical.

Equilibrium State

When reaction is reversible, the velocity of forward and reverse reaction is approx to be similar, the state is Equilibrium state of reaction.

A(g) + B(g) <=> C(g) + D(g)

N2(g) + 3H2(g) <=> 2NH3(g)

H2O(l) + H2O(l) <=> H3O⁺(aq)+OH^_

The ionic equilibrium is subject of physical chemistry touching with velocity of reaction.

Generally ionic reaction attain equilibrium in aqueous solution.

NaCl(aq) = Na⁺(aq) + Cl^(-)(aq)

​

Theory of acids and bases

Acidic and basic property of substance is explain by different scientist, these are known theory of acids and bases.there are three different theory of acids and bases

(a) Arrhenius theory :

According to this theory

Substance or compound are dissociate H⁺ ions in aqueous solution is called Acids.

(a) Arrhenius theory :

(OR)

Acids is substance which are dissociate H⁺ ions in aqueous medium.

e.g. HCl(aq), H2SO4(aq), HNO3(aq)

CH3COOH(aq) are acids and acidic property of substance.

Base:

Substance or compound are dissociate OH^- ions in aqueous medium is called Bases.

Arrhenius theory

Bases are substance which dissociates OH^- ions in aqueous medium.

NaOH, KOH, CH3NH2, NH4OH

are Basic nature substance.

The proton H+ formed in dissociation of an acid is reason of acidic property of acids.

Similarly OH- ions formed in dissociation of bases is reason of basic character of bases.

This theory does not explain the role of solvent in it.

Limitation of Arrhenius theory

(1)This theory is explain the acidic and basic property of substance in aqueous solution.

(2)This theory does not explain the acidic and basic character of NH3, Na2CO3,FeCl3 etc.

(3) This theory is limited for specific substances.

(4) This concept is based on ionisation reaction.

Lorry–BrØnsted theory

According to this theory acid is substance or species having tendency to lose or donate a proton(H)+) to anther substance.

HCl(aq) + H2O = H3O(aq)+ + Cl(aq)-

acid -1 base-2 acid-2 base-1

According to this theory base is substance or species having tendency to accept proton from other substance.

Na2CO(aq) + H2O = NaHCO3(aq)- + OH-

Base 1 acid 2 acid 1 base 2

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Lorry–BrØnsted theory

According to this theory the Lorry-Bronsted acid base are conjugate

acid-base pair of chemical reaction. This is depend upon on the chemical affinity of substance.

also it is depend upon on the electro negativity and electron density of atom.

Conjugate acid-base Pair

HSO4- + H2O = H3O+ + SO4—

H2O + CO3-- = HCO3- + OH-

H2O + NH3 = NH4+ + OH-

H3PO4 + H2O = H3O+ + H2PO4-

H2S + H2O = H3O+ + HS-

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Lewis Concept

According to this concept, an acid is a substance which is able to accept a pair of electrons (i.e. lone Pair electron acceptor) and a base is substance which is able to donate a pair of electron (i.e. lone pair electron donor)

Following examples of such acid-base reactions given below

:B:F3 + :F: - = [BF4]-

complex

Ag+ + 2 CN- = [Ag(CN)2]--

Acid Base complex

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Lewis Concept

We are see that the Lewis definition extends the acid-base concept to reaction in which protons is not involved. Lewis concept of acid and base is able to explain systems where protonic materials not involved.

It is clear that Lowry-Bronsted bases react by donating electrons to a proton, Lorry-Bronsted base is also a Lewis base.

Many substance like FeCl3, AlCl3, SnCl4, etc are acid.

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Lowry-Bronsted acids and Lewis � acids

According to Lowry-Bronsted concept, an acid must have proton, which it remove in aqueous solution. Acid dissociate proton in solution. HCl,HNO3 and organic acid etc.

BF3, CO2(g) are Lewis acid but there is not a proton in formula. They have tendency to accept lone pair electron from another substance. There fore Lowry Bronsted acids are Lewis acids.

Lewis Concept

However a Bronsted-Lowry acid must have a proton available for transfer to another molecule , but this is not require by the Lewis definition.

Lewis definition of a base covers all substance which would be classed as such according to Lowry- Bronsted concept also.

But his definition of acids includes many substances other than proton donors as well as.

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Lewis Concept

The Lowry–Brϕnsted definition suggests that a strong acid has a tendency to transfer a proton to another molecule and a strong base is one with a affinity for proton.

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Strength of Acids and Bases :

In general , the ionization of an acid HA in water may be represented as :

HA + H2O = H3O+ + A-

The equilibrium constant for this acid-base reaction is given by

Ka = [Keq] = [H3O+][A-] …. …

[HA]

The strength of an acid depends upon its tendency to lose protons and strength of base depends upon its tendency to gain protons .

E.g. HCl is stronger acid than CH3COOH because HCl have a greater tendency to donate proton, CH3COOH have less tendency to donate (loose) proton.

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Strength of acid

Strength of an acid is describe by acid dissociation constant(Ka).The strength of acid is compare with dissociation constant in aqueous solution. For strong acid dissociation constant is large. For strong acid equilibrium is lie on product side. Reaction of acid with water is more useful to decide the acid strength.

CH3COOH + H2O = H3O+ + CH3COO-

Acetic acid is weak acidic because have less tendency to loss proton in aqueous solution.

P^H Scale

The Denis chemist, Sorenson introduced a more convenient scale of hydrogen ion conc. in aq. Solution,

it define as P^H of solution.

It is [ -log10 ^H+ = P^H]

It define small quantity of H⁺ ion in aq. Sol.

e.g 0.1M acid sol. 0.01M …..10^-14M sol.

It also define the acidity and basicity of sol.

For acidic sol. P^H < 7 and basic sol. P^H >7 ,

when P^H = 7, sol. is neutral.

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P^OHscale

Similarly P^OH of sol.

It is [ -log10 ^OH= P^OH]

P^H+ P^OH= 14

Self Ionization of water:

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Self ionization of water

Water ( H₂O) is Amphotaric chemical unit

which behave as an acid and base, this reaction is given by

H₂O_(aq) + H₂O_(aq) = H₃O⁺_(aq) + OH^-_(aq)

Acid -1 base-2 Acid-2 Base -1

According to ionic equilibrium,

[H₃O+][OH^-] = K_w

K_w is Ionic product of water

[H₃O+][OH^-] = 10^-14

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Ionic Product of water

At 25 ^oC temp and 1atm pressure, the ionic product of water is given by following eq

[H₃O+][OH^-] = K_w

K_w is Ionic product of water

At 25 ^oC temp and 1atm pressure, the ionic

product of water is given by following eq

[H₃O+][OH^-] = 10^-14

• At ionic equilibrium

[H₃O+]=[OH^-] = 10^-7 mole/lit

P^H+ P^OH= 14

Now, [ -log10^H+= P^H] and [ -log10^OH-= P^OH]

=> [ -log10^10-7= P^H]

=> P^H = 7 and P^OH = 7

​

=> P^H+ P^OH= 14

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Electrolyte

• The substance are dissociated in (aq) sol. define as electrolyte. They are conducting charge of electric current. they are divide as strong and weak electrolyte.

HCl_(aq), H₂SO₄ are strong electrolyte, CH₃COOH,

CH₃NH₂ are weak electrolyte.

Strong electrolyte and Weak electrolyte

• Strong electrolyte:

The substance are completely dissociated

in (aq) sol. define as strong electrolyte.

KCl, NaCl, HCl, are strong electrolyte

• Weak electrolyte:

The substance are partially dissociated in

(aq) sol. define as weak electrolyte.

CH₃COOH, HCOOH, CH₃OH are weak electrolyte

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Strong electrolyte:

Strong Electrolyte

• Good electric charge conductor and electric current
• They are completely dissociate in (aq) sol.
• The value of ionization constant is high
• KCl, NaCl, HCl, are strong electrolyte

Weak electrolyte:

• Weak electric charge conductor and electric current
• They are Partially dissociated in (aq) sol.
• The value of ionization constant is law
• CH₃COOH, HCOOH, CH₃OH are weak electrolyte.

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Ionization principle

• Le’Chatalier Principle:

‘When a stress or change is applied on system in equilibrium, the system tends to adjust itself so as to reduced the stress’

Common ion effect : The suppression of the dissociation of weak electrolyte on the addition of similar ion is called common ion effect.

[NH4Cl + NH4OH]

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Solubility(S) and Solubility Product(Ksp)

Substance are soluble and insoluble in water, The product of ionic concentration in solution is term as solubility, this is represent as symbol S.

According to solubility of substance, the salts are soluble, insoluble and sparingly soluble.

In a saturated solution of sparingly soluble salt, the product of ionic concentration is constant at given temperature is called solubility product. This constant is given by Ksp.

AgCl(s), PbSO4(s) etc are sparingly soluble salts.

�Solubility and Solubility Product � (Ksp)�

At 25 0C temperature, the solubility of sparingly(partially)soluble salts is define as solubility product of solution. It is represent as (Ksp), It is a product of insoluble salt, conc. of ion dissociates in aq. Solution.

eg. AgCl(s), PbSO₄(s) are insoluble salt, similarly

(OH^-) and S^-- of metal ions are sparingly soluble salt.

Now Ksp[AgCl(s)] = [Ag+][Cl^-] and

Ksp[PbSO₄(s)] = [Pb+2][SO4^-2]

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Solubility (s)

• Solubility (s)=

• The equilibrium between saturated salt

​

• AgCl_(S) = [Ag⁺][Cl^-] = (S x S)
• => (S) =

​

• At equilibrium of reaction, the value of

[Ag⁺][Cl^-] is constant is define as solubility product of AgCl_{(S) .}

The solubility is expressed in mole/lit (M)

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Solubility Product [Ksp]

• The equilibrium constant of the reaction is given by following Equation

K = AgCl_(S

=> Ksp = K[AgCl_(S]

= [Ag⁺][Cl^-]

= S²

S = [Ag⁺]=[Cl^-] =

[Ag⁺] = [Cl^-] = 1.7x10^-5 m

Solubility Product [Ksp]

Salt are AB, AB₂, A₂B, AB₃ types, the solubility of the these sparingly soluble salts is given as

AB, S =

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• AB₂ S= 4 (cube route)

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Cal.

Cal :

Calculate the solubility of AgCl in a solution of .1M AgNO3.

[ K(sp) of AgCl is 2.8x 10^-10 ]

Solution:

AgCl(S) = [Ag⁺(aq)][Cl^-(aq)]

K(sp) = [Ag⁺] [Cl^-]

We know that solubility i.e.

[Ag⁺(aq)]=[Cl^-(aq)] = [K(sp)] / [Ag⁺

Solubility and solubility Product(Ksp)

molarity of AgNO3. (∴ Ag⁺ = 0.1 M )

[Cl^-(aq)] = 2.8x 10^-10 ]/ (0.1)

= 2.8x 10^-9 M

Ans : S = 2.8x 10^-9 M

Solubility and solubility Product(Ksp)

• For example, let us denote the solubility of Ag₂CrO₄ as S mol L^–1. Then for a saturated solution, we have

Ag₂CrO₄(s) ⇌ 2Ag⁺+CrO₂^–4

[Ag⁺] = 2S    [CrO₄^2–] = S.

Ksp = [Ag+]² [CrO₄^2–]

(2S)² (S) = 4S³ = 2.76×10^–12

S = (K_s/4)^1/3 = (6.9 x 10^-13)^1/3 = (0.69 x 10^-12)^1/3 = 3√ (0.69) x 10^-4 = 0.88 x 10^-4

Thus the solubility is 8.8×10–5M.

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� Important effects:�

Important effects:

• For highly soluble ionic compounds the ionic activities must be found instead of the concentrations that are found in slightly soluble solutions.
• Common Ion Effect: The solubility of the reaction is reduced by the common ion. For a given equilibrium, a reaction with a common ion present has a lower Ksp, and the reaction without the ion has a greater Ksp.
•  Salt Effect (diverse ion effect): Having an opposing effect on the Ksp value compared to the common ion effect, uncommon ions increase the Ksp value. Uncommon ions are ions other than those involved in equilibrium.

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Ion Pairs:

• Ion Pairs: With an ionic pair (a cation and an anion), the Ksp value calculated is less than the experimental value due to ions involved in pairing. To reach the calculated Ksp value, more solute must be added.
• Solubility S increases with increase in [H⁺] or decrease in pH

Selective Precipitation

In both the Quantitative and Qualitative

analysis, the advantage of difference

solubility of salt is used to remove one

salt from several salt in solution i.e.

selective precipitation.

Solubility product of [ AgCl_(S)] = 1.7x10^-10 M

And

[Ag₂CrO₄ ] =1.7x10^-12M

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Selective Precipitation

Precipitation of AgCl(s) is before the

Ag₂CrO_4(s)

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salts

Reaction of acid and base is called neutrilization reaction, and gives salts.

This salt is divided according to acid base strength.

(1) weak acid and strong base salt (basic salt)

(2) weak base and strong acid salt(acidic salt)

(3) weak acid and weak base salt (uncerntain)

(4) strong acid and strong base ( neutral)

Dissociation constant (Ka) & Kb

Dissociation of aceticacid

Ka=1.8 x 10^-5 mole/lit

Methanoic acid (HCOOH)

Ka = 1.8 x 10^-4

Similarly the Kb is dissociation of base, this is shows that, the tendency of dissociation is lower. Hence, they are weak acid and base.

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Hydrolysis of weak acid and base

Recction of weak acid and base in aquous soltion is called hydrolysis reaction.

CH3COOH + H2O = H3O⁺ + CH3COO^-

CH3COO^- + H2O = CH3COOH + OH^-

Hydrolyis of weak acid salt (A-) : The reaction of weak acid salt with water is called hydrolysis reaction.

Now,

here weak acid is represent by HA and its salt is A-

Hydrolysis and Hydrolysis constant� (Kh)

Hydrolysis of weak acid salt is given by following generel chemical reaction.

A- + H2O = HA + OH-

according to chemical Equilibrium and mass action law

Keq = [HA][OH-]

[A-][H2O]

Keq[H2O] = [HA][OH-]

[A-]

Kh = [HA][OH-] ..... (1)

[A-]

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Kh

Kh is known as hydrolysis constant.

It is represent the hydrolysis of acid salt.

In above chemical equilibrium other two equilibriums are present with parelal hydrolysis.

(a) Self ionization of (H2O) water.

H2O = H+ + OH-

Similarly, this is given by ionic product of water

Kw = [H+][OH-] ..... (2)

Relation of Kh,Ka and Kw

(b) ionization of Wak acid

HA + H2O =H3O+ A-

from above chemical eqilibrium,

Ka = [H3O+][A-]/[HA]

= [H+][A-]/[HA]

... ... … (3)

by division of (2) with (3), we get following

relation

= [H+][OH-]/ [H+][A-]/[HA]

=[HA][OH-]/[A-]

Kh = Kw/Ka from eq (1) … … … (4)

Relation of Kh,Ka and Kw

from above Eq(4), we are find constant (Kh) and (Ka)

Cal.

Hydrolysis of weak base salt(AB)

Kh = Kw/Kb

B+ + H2O = BOH + H+

The salt weak base and strong acid is acidic in nature. This is due to hydrolysis of salt, in aqueous medium.

e.g. NH4Cl. Is acidic in nature.

Now, AB is general formula of salt of weak base, the hydrolysis of salt is given by following chemical reaction.

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Hydrolysis of Weak base Salt

B+ + H2O = BOH + H+

​

according to chemical equilibrium law,

equilibrium constant is given as,

Keq = [BOH][H+]/ [B+][H2O] … … (1)

According to law of mass action,

Kh = [BOH][H+]/[B+] … … … (2)

Now, we know in this reaction, two other is reaction parallel with hydrolysis,

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Hydrolysis(AB) Salt

• (a) Self ionization of water:

H2O = H+ + OH-

Ionic product of water is given as,

Kw = [H+][OH-] = 1.0 x 10-14 …(3)

(b) Ionization of Weak base:

BOH = B+ + OH-

ionization constant is given as,

Kb = [B+][OH-]/[BOH] … .. (4)

Hydrolysis

From above eq. No (3) and (4)

(3) is divided by (4) we get,

Kh = Kw/Kb ….. (5)

We are calculate the constant (Kh) and (Kb) from data of eq.No (5).

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Buffer Solution

The chemical mixture or substance which are used as protection of change of PH of solution or chemical reaction is called Buffer solution.

Or

The solution which are resist change of PH value, on even addition of small amount of an acid or base in reaction is known as buffer solution.

Buffer solution is mixture of weak acid and salt of strong base.

Similarly, it is mixture of weak base and strong acid salt.

Buffer Solution

e.g.

(1) [CH3COOH+ CH3COONa]

(2) [NH4Cl +NH4OH]

(3) Na2B4O7

etc.

Characteristic of buffer solution:

(1) Buffer solution has definite PH range.

(2) PH of these solutions does not change after long time.

characteristic of buffer solution

(3)The PH of buffer solution does not change with dilution of solution i.e. constant

(4) PH of buffer solution does not change on even addition of small amount of acid or base.

buffer Action: The property of the solution to resist the change in Its PH value on the addition of small amount of strong acid or base is known as buffer action.

Types of buffer solution

There is two types of buffer solutions

(a) Acidic Buffer solution

(b) Basic buffer solution

Acidic buffer solution: an aqueous mixture of weak acid and its strong base salt is called acidic buffer solution.

preparation: When weak acid and its strong base salt is mixing in certain amount or ratio, we are prepare the acidic buffer solution.

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Basic buffer solutions

• The aqueous solution of weak base and its strong acids salt is called basic buffer solution.
• Preparation:
• It is preparing by addition of weak base and its strong acids salt.

NH4Cl + NH4OH ,

NH4OH + NH4NO3

Mechanism of buffer action

As Mixture of CH3COOH and CH3COONa is a acidic buffer solution.

(i) CH3COOH + H2O -🡪 H3O+ + CH3COO-

(ii)CH3COONa + H2O -🡪 CH3COOH + NaOH

according to above reaction weak acid is partially ionized and salt is completely dissociated.

now, on addition of acid or base, the reaction is shift with minimize the addition effect.

On addition of acids solution, the proton is gives

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Buffer solution

• Excess acid is reform CH3COOH with acetate ions,

CH3COO- + H+ = CH3COOH

(CH3COO–) is dissociate by salt and weak acid

Similarly

• and on addition of base (OH-) excess base is reform H2O and CH3COONa,

CH3COOH + OH- = CH3COO-+ H2O

these all reaction minimized the effect of addition of acid-base.

Buffer Action

P (Ka) = PH – log([Salt]/[Acid])

Consider an acidic buffer solution, containing weak(HA) and its salt(AB) with strong base BOH,

Ionization of weak acid HA is given by

HA + H2O = H3O+ + A-

=> [H3O+] = [H+]

= Ka([HA]/[A-])

=> [H+] = Ka [HA]/[A-]

buffer action

• taking [-log] of both side of Eq.(1)

-log[H+] =-log[Ka] – log[acid]/[salt]

=> PH = Pka – log[acid]/[salt]

above eq. is Known as Henderson-Hasselbach equation.

or

PH = Pka + log[salt]/[acid]

PH of buffer solution is calculated from above equation.

Basic Buffer Solution

• Basic buffer solution is containing a B+ (ions), the ionic equilibrium is shown as

B+ + H2O = BOH + H+

Kb = [BOH] [H+]/[B+]

=>[H+] = Kb[B+]/BOH

Apply(-log) on both side of above eq.

-log[H+]= -log[Kb] –log[salt]/[base]

PH = PKb –log[B+]/[BOH]

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Buffer capacity(ß)

• Buffer capacity (ß) =(∆V/∆PH)

∆V = Volume of acids or base

∆PH = 1

Buffer capacity is define as amount of acid or base required to change 1 unit of PH of 1 lit solution.

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Indicator

• Indicators are acids or bases which change the color of solution in terms of PH change of solution. The substance which are the change the color of medium or reaction or titration at end of the reaction is called indicator.
• OR

acidic or basic substance which are able to change the color with change of PH of solution.

Generally they are organic substance, they have characteristics color in two different forms.

Indicator

we consider, as Indicator is acidic substance.

its general structural formula is given by HIn and basic indicator is HOIn.

According to dissociation and ionic equilibrium of state of reaction,

Hin + H2O = H3O+ + In-

color-I color –II

e.g. Phenophthalien is substance shows two different color in acidic and basic solution.

In acidic solution its has no color but in basic

Indiactor

Medium its color is pink.

it is most useful in acid base titration.

There is number of substance which shows these color characteristics, so they are used as an indicator.

different reactions and analysis are operate by using indicators.

Mechanism of indicators

Indicator is represent by following general structural formula. It is a weak acids and bases.

Hin and HOIn

We are see, Hin is dissociates H+ and gives In- ions,

Hin + H2O = H3O+ + In-

[color –I] [color -II]

According to equilibrium of reaction,

Dissociates form

• K(In) = [H3O+][In-]/ [HIn] (1)
• => [H3O+] = K(In) [HIn]/[In]

= K(In) [un-dissociates Form]

[dissociate Form]

= K(In) [color –I]/[color-II]

effective PH limits of color-I and color -II of indicator is given by following equation.

-log[H3O+] = -log[K(In)]–log[color-I]

[color-II] (2)

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Indicator exponent

• PH = P[K(In)] – log[color-I]/[color-II] (2)
• From this eq. we are find PH range of indicators. This is define as useful PH range of Indicator

SQ

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• Define following

(1) solubility and solubility product

(2) P^{H and} P^OH of solution

(3) strong acid and weak acid, give ex.

(4) define and explain sparingly soluble salt.

(5) selective precipitation

(6) Conjugate acid- base pair

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SQ

(7) Explain the different concept of acid-base

Theory (a) Lewis concept

(b) Lowry– Bronsted

(c) Arrhenius concept

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LQ

(1) Discuss brief concept of Lewis acid base

theory with suitable example.

(2) Explain the self ionization of water and

prove P^H+ P^OH= 14

(3) Define sparingly soluble salt and explain

the term Solubility product with suitable

example.

(4) discuss the sort account of selective

precipitation.

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LQ

(5) Discuss the theory of buffer solution.

(6) What is acid–base indicator ,discuss the

brief theory of acid-base indicator.

(7) Discuss the relation between the hydrolysis

constant and dissociation constant of weak

acid and base.

Example.

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