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05/01:

UNIT-1: Buffer Solution

Introduction:

• Some reactions,
  ➢ H2O + CO2 ---> H2CO3 (Carbonic Acid) [soln becomes Acidic]
  ➢ H2O + SiO2 (Silica) ---> H4O4Si (Silicic acid/Orthosilicic acid)
  ➢ H20 + CaO (Lime)---> Ca(OH)2 (Calcium hydroxide
  ➢ H20 + SO2 (sulphur dioxide/sulfurous anhydride) ---> H2SO3 (Sulfurous acid)
  ➢ H20 + SO₃ (Sulfur trioxide) ---> H₂SO₄ (Sulfuric acid)

• Buffers are solutions that contain a weak acid and its a conjugate base or vice versa; as such, they can absorb excess H+ions or OH– ions, thereby maintaining an overall steady pH in the solution.

• 2 Types:
  Acidic Buffer (weak acid + strong salt i.e. salt of the weak acid with strong base)([pH < 7]
  Basic Buffer (weak base + strong salt i.e. salt of the weak base with strong acid) [pH > 7]

• Generally inorganic/mineral acids are strong and organic are weak.

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05/02:

Law of mass action:

• Can only be used for weak electrolyte. Strong electrolyte's equation is irreversible hence no equilibrium

  ➢ A + 2B ⇌ C
  ➢ Rf ∝ k1[A][B]2 (k1= Rate constant for forward reaction )
  ➢ Rb = k2[C] (k2= Rate constant for backward reaction)
  ➢ k=[𝐶]/[𝐴][𝐵]2 (k= k1/k2, equilibrium constant)

Buffer Capacity:

• the number of moles of an acid or base necessary to change the pH of 1 liter of buffer solution by 1.

Common Ion Effect

• the addition of a common ion prevents the weak acid or weak base from ionizing as much as it would without the added common ion.

Calculation of pH of acidic buffer / Henderson equation (Hessel Balch equation) for acidic buffer: [Imp Question!!!!]

• dissociation (equilibrium) constant ‘ka’ of weak acid and the concentration of weak acid and its salt used are required

  ➢ HA ⇌ H+ + A- (Consider a weak acid ‘HA’ which dissociates as..)
  ➢ ka = [𝐻+][𝐴−]/𝐻𝐴 [Law of mass action]
  ➢ [H+] = 𝑘𝑎[𝐻𝐴]/[𝐴-] -- (i)
  ➢ BA --> B+ + A- (highly ionizable salt of acid, ‘BA’ is added to it)
  ➢ [A-] = [BA] --- (ii) [all the A- ions have come from the salt (BA) due to common ion effect]
  ➢ pH = Pka + log [𝑆𝑎𝑙𝑡]/[𝐴𝑐𝑖𝑑] [From (i) and (ii), taking -ve log on both sides]

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05/08:

Le Chatelier's principle :

• When a settled system is disturbed, it will adjust to diminish the change that has been made to it
• This method is also used to make weak electrolyte strong
  

Calculation of pH of basic buffer / Henderson equation (Hessel Balch equation) for basic buffer: [Imp Question!!!!]

• Quick derivation

  ➢ BOH ⇌ B+ + OH- (Consider a weak base ‘BOH’ which dissociates as..)
  ➢ kb = [B+][OH−]/BOH [Law of mass action, kb= dissociation constant of base]
  ➢ [OH-] = 𝑘b[BOH]/[B+] -- (i)
  ➢ BA --> B+ + A- (highly ionizable salt of base BA is added)
  ➢ [B+] = [BA] --- (ii) [concentration of unionized base is nearly equal to original concentration due to common ion effect]
  ➢ pOH = Pkb + log [𝑆𝑎𝑙𝑡]/[Base] [From (i) and (ii), taking -ve log on both sides]

Significance of Henderson's equation

• PH of buffer solution can be calculated
• The dissociation constant of weak acid /base can be determined
• A buffer solution of desired pH can be prepared

Buffer Mechanism of acidic buffer solution [Sure Question!!!!]

• Link
  

• TLDR

  ➢ CH3COONa --> CH3COO- + Na+ [strong electrolyte which ionized almost completely]
  ➢ CH3COOH ⇌ CH3COO- + H+ [weak electrolyte ionizes feebly]
  ➢ Ka = [CH3COO-][𝐻+]/[CH3COOH] [law of mass action ]
  ➢ When (strong) acid added, H+ ion concentration increases which changes pH of solution. So, with Le Chatelier's principle, almost half of H+ from added acid combines with the ion given by strong salt and forms undissociated weak acid.
  ➢ When (strong) base added, OH- ion concentration increases which reacts with the H+ ions of weak acid and disturbs the equilibrium and changes pH of solution. So, with Le's principle, the weakly ionizable weak acid now quickly ionizes so as to fulfill the loss of H+ ions.
  ➢ This method is also used to make weak electrolyte strong

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05/09:

Buffer Mechanism of basic buffer solution [Sure Question!!!!]

• NH4Cl --> NH4+ + Cl- [strong electrolyte which ionizes almost completely in aq. solution]
• NH4OH ⇌ NH4+ + OH- [weak electrolyte ionizes feebly]
• kb = [NH4+][𝑂𝐻−]/[NH4OH] [law of mass action]
• pOH = Pkb + log [𝑆𝑎𝑙𝑡]/[𝐵𝑎𝑠𝑒] [Henderson’s eqn for basic buffer solution]
• When (strong) acid added, the H+ ion conc. increases which reacts with the OH- ion of the weak base. To maintain equilibrium, the week base now quickly ionizes to fulfill the loss of ON- ions.
• When (strong) base added, OH- ion concentration increases which changes pH of solution. So, large number of OH- contributed by the added base combines with the common ion given by strong salt and forms undissociated weak base.

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05/15:

Numerical problems related to buffer [Sure Question!!!!]

pH and pH scale [NTI - Not that Important]

• -ve common logarithm of Hydronium ion (H30+) concentration expressed in the unit of molarity
• the scale which expresses the degree of acidiy of a solution in terms of pH

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05/16:

UNIT-2: Electrochemistry

Types of cells

• Electrolytic cells

  ➢ electrical energy to chemical energy
  ➢ electroplating of copper

• Electrochemical Cells

  ➢ chemical energy to electrical energy
  ➢ Daniell cell

Cell Terminologies

• Direct/Indirect Redox Reaction

  ➢ both oxidation and reduction occur simultaneously with oxidizing and reducing agents being in direct or indirect contact

• Anode/Cathode

  ➢ electrode where oxidation/reduction half takes place
  ➢ acquires -ve/+ve charge
  ➢ more/less electro positive
  ➢ In electrolytic cell, cathode (-ve electrode) acquires cations (+ve ion) and anode(+ve electrode) acquires anions (-ve ion)

• Electrolyte

  ➢ salt solution used in cell

• Cell Potential (emf)

  ➢ maximum potential difference between two electrodes of a galvanic or voltaic cell

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05/22:

Galvanic cell diagram or Symbolic Representation

• Zn/Zn2+(a=0.2M) // H+/H2(Pt) E=+1.1V

Calculation of emf of cell

• in terms of reduction

  ➢ E(cell) = E(cathode) - E(anode)

• in terms of oxidation

  ➢ E(cell) = E(cathode) + E(anode)

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05/23:

Electrolytic V/s Galvanic cell

• Link
  

Galvanic cell

• Every Metal loose electrons in their surface before the circuit is completed. The crowdness of free electron in metal electrode surface causes electron pressure which develops potential.
• After a circuit is completed with a conductor, the electrons in surface flows from high pressure to low and metal ions falls to electrolyte surface. Which makes the neutral metal acquire -ve or +ve charge and becomes anode or cathode performing oxidation or reduction.

Flow of current

• When Eapplied < Ecell,

  ➢ cell -> external circuit
  ➢ forward reaction
  ➢ more electro positive to less (Zn to Cu)
  ➢ natural

• When Eapplied > Ecell,

  ➢ external circuit -> cell
  ➢ backward reaction
  ➢ less electro positive to more
  ➢ nonspontaneous

2 types: [Imp Question!!!!]

• Reversible cell and its illustration

  ➢ Zn + Cu cell
  ➢ When Ecell infinitesimally > Eapplied, Zn + Cu2+ ---> Zn2+ + Cu
  ➢ When Ecell infinitesimally < Eapplied, Cu + Zn2+ ---> Cu2+ + Zn
  ➢ Hence the diff is negligibly small and when Eapplied infinitesimally > Ecell, the cell reaction occurs exactly in reverse reaction

• Irreversible cell and its illustration

  ➢ Zn + Ag cell
  ➢ When Ecell infinitesimally > Eapplied, Zn + H+ --> Zn2+ + H2
  ➢ When Eapplied infinitesimally > Ecell, 2Ag + 2H+ --> Ag+ + H2
  ➢ the cell reaction doesn't occur exactly in reverse direction

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05/29:

Single Electrode (Half Cell) Potential and its origin [Sure Question!!!!]

• 3 cases when electrode is dipped into its ion solution:

  ➢ metal ion of the solution may collide at the surface of the electrode and no changes takes place (no charge developed)
  ➢ after colliding,metal atom of electrode looses its electron in its own surface and goes into the solution (-ve charge developed in electrode)
  ➢ after colliding, metal ion of electrolyte gains electron from the surface of electrode and gets discharged into elemental metal (+ve charge developed in electrode)

• if metal ion of solution has greater tendency to undergo reduction, case 3 happens

• if metal atom of electrode has greater tendency to undergo oxidation, case 2 happens

• thus individual electrode develops potential with respect to the solution which is single electrode potential

Hydrogen Electrode

• Standard H-electrode (S.H.E) is prepared by dipping a pt-electrode into 1M H+ ion solution and H2 gas under 1atm pressure is bubbled through a gas hood at 25 degree Celcius

Experimental Determination of single electrode potential

• If standard H- electrode is at cathode and given electrode is at anode:

  ➢ Eo cathode = 0 volt
  ➢ Eo cell = -Eo anode [Eo cell = Eo cathode - Eo anode]

• if standard H- electrode is at anode and given electrode is at cathode:

  ➢ Eo anode = 0 volt
  ➢ Eo cathode= Eo cell [Eo cell = Eo cathode - Eo anode]

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05/30:

No recording for the class :(

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06/05:

Application of Nernst equation

• Calculation of electrode potential of a half cell

  ➢ Eo – 2.303 𝑅𝑇/𝑛𝑓 𝑙𝑜𝑔 [1]/[𝑀^𝑛+]

• to calculate the cell potential

  ➢ Ecell = Eocell – 2.303 𝑅𝑇/𝑛𝑓 log⁡𝑘
  ➢ less reduction electrode potential [more electro positive, easily e- loss, occurs oxidation] - anode
  ➢ more reduction electrode potential [less electro positive, hardly e- loss, occurs reduction]-cathode

• to study the effect of electrolyte on the electrode potential

• calculate the equilibrium constant for the cell reaction

  ➢ Log k = 𝑛×Eocell /0.0591

• to find the concentration of one of the ionic species in the cell if that of other is known

• to find the pH of a solution containing H+ ions

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06/06:

Electro chemical series and its Application

• the sequential arrangement of various electrodes in the order of their increasing standard reduction potential values

• Calculation of standard cell potential

  ➢ Eocell = Eocathode – Eoanode

• predicting feasibility of cell reaction

  ➢ del G = -nfE (the value of del G should be -ve in order for rxn to be feasible)
  ➢ i.e. if Eocell is negative, the cell is not possible and is wrong

• predicting oxidizing and reducing tendency of electrode material

  ➢ anode - less E0 value - more electropositive - great tendency to loose e-s & undergo oxidation
  ➢ more oxidation tendency are strong reducing agent. thay oxidize themselves but reduces others
  ➢ from top to buttom, the tendency to reduce others (self oxidizing power) goes on decreasing while the tendency to oxidize others (self reducing power) goes on increasing

• predicting the direction of metal displacement reaction

  ➢ when metal having less Eo value is dipped into the salt solution of another metal having more Eo value then metal with less Eo value displaces the another metal

• Prediction of liberation of H2 gas from acid

  ➢ Metals having less Eo value than H+ ions when dipped into the solution of H+ ions, easily reduce the H+ ions into molecular hydrogen

• To compare the Speed of a redox reaction

  ➢ the electrodes which are closer in the series react with slower rate than those which are farther in the series

• Used for the selection of electrode material

  ➢ The electrode having less electrode potential are used to make the material of anode and electrodes having more Eo values are used as the material of cathode.

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06/12:

No recording for the day :(

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06/13:

Prevention of corrosion

• surface coating

  ➢ liquor like enamel, paint, varnish

• coating a layer of corrosion resistance metal

  ➢ Metals like Ni, Cr

• Galvanization

  ➢ coating a layer of Zn (Tin - tinning) over the surface

• forming alloys of metals

  ➢ Alloys are relatively less sensitive than the component metals

• depositing a layer of phosphate or chromate

  ➢ dip into the solution of phosphoric acid and the solution of chromate salt then a layer of metallic sulphate or chromate gets deposited over the surface

• cathodic protection or sacrificial protection

  ➢ connect metal to to the more electropositive metal like connecting iron to aluminium

Mechanism of corrosion

• Fe --> Fe2+ + 2e- (Anode)
• O2 + 4H+ + 4e- --> 4OH- (Cathode)
• 2Fe + O2 + 4H+ --> 2Fe2+ + OH- (overall)
• Fe2+ + OH- --> Fe(OH)2
• 4Fe(OH)2 + O2 +2H2O --> 4Fe(OH)3 --> 2Fe2O3 (3H2O) --> Fe2O3 .3H2O

Numerical problem related to Nernst equation [Sure Question!!!!]

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06/21 - :

UNIT-3: Organic Chemistry

Characteristics of S and P-Block Elements

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